Basic Ideas in Chemistry

Matter in Chemistry

When you think of studying chemistry, an image of someone mixing chemicals in a lab likely comes to mind. However, you'll learn that studying

chemistry involves many aspects. Chemistry is the study of the composition, structure, properties, and transformations of matter.

As you learned in previous science courses, matter is anything that occupies space and has mass. Mass is a measure of the amount of matter that

an object contains. Most of Earth is made up of matter. You, the chair you're sitting on, the computer you're using, and even the air around you are

all made of matter. You may wonder what is not made up of matter. Radio waves, heat from a �re, and the sound from a car horn are examples of

forms of energy and aren't made of matter. Your thoughts and ideas also aren't matter because they don't occupy space or have mass. �ink about

the matter around you, what these objects are made from, and how they were made.

From Alchemy to Chemistry

�roughout human history, people have tried to convert matter into more useful forms. Stone Age humans chipped pieces of �int into useful tools

and carved wood into statues and toys. �ese endeavors involved changing the shape of a substance without changing the substance itself. As

knowledge increased, humans began changing the composition of substances. Humans learned to control �re and use it to cook, make pottery,

and smelt metals. Hides were cured to make garments, copper ores were transformed into copper tools and weapons, and grain and other

ingredients were combined and cooked into bread. Subsequently, humans learned to separate and use speci�c components of matter. Various

drugs, including aloe, myrrh, and opium, can be isolated from plants. Dyes such as indigo and Tyrian purple were extracted from plant and animal

matter. Metals were combined to form alloys. For example, copper and tin can be mixed to make bronze. More complex smelting techniques

produced iron. Alkalis were extracted from ashes, and soaps were made by combining alkalis with fats. Alcohol was produced by fermentation and

puri�ed by distillation.

Attempts to understand the behavior of matter began more than 2500 years ago. Until modern times, these attempts were carried out by

philosophers, not scientists. As early as the sixth century BC, Greek philosophers discussed a system in which water was the basis of all things. You

may have heard of the Greek theory that matter consists of four components: earth, air, �re, and water.

In the Middle Ages, alchemy was common. In this specialized branch of philosophy, alchemists believed that it was possible to transform some

materials into others. Alchemists sought to transform "base metals" such as lead into "noble metals" such as gold. �ey also worked to create

elixirs to cure disease and extend life. Alchemy was the forerunner of chemistry. Like chemistry, alchemy sought to understand and manipulate

matter. Unlike chemistry and other branches of modern science, alchemy was based on a supernatural, mystical view of the world. Modern

science is based on observations and reason.

In the seventeenth century, science began to resemble its modern form. Until this time, it was believed that earth, air, �re, and water were the four

basic components of matter. However, this notion was dispelled when Robert Boyle, an Irish chemist, published a book in which he proposed a

theory describing an "element" as something that couldn't be broken down into smaller pieces. Boyle de�ned how experimental science should

be conducted, helping to distinguish chemistry from alchemy.

Science development accelerated rapidly during the Age of Reason in the eighteenth century. In addition to Boyle, many early scientists made key

contributions to chemistry: Joseph Priestley, Jacques Charles, Joseph Proust, John Dalton, Amedeo Avogadro, and Antoine Lavoisier. Lavoisier is

widely considered to be the father of modern chemistry, and the following are just some of his contributions:

Developing a system of naming chemical compounds

Discovering the role of oxygen in combustion

Helping to formulate the metric system

Writing the �rst modern chemistry textbook

Chemistry has continued to advance, improving scientists' understanding of matter and the ability to harness and control its behavior.

Describing Matter

�e study of chemistry requires that you understand the nature of matter, which includes methods for describing and categorizing matter.

Matter can be described physically or chemically. Physical properties—density, weight, temperature, hardness, and color—are characteristics

that can describe a substance. A physical change doesn't change the composition of a substance.

[1]

Example 1: If wax melts, the substance is still wax even though it's no longer solid.

A chemical change involves a substance changing composition.

Example 2: When iron combines with oxygen, this forms a new substance—rust or iron oxide.

While physical changes are easily reversible, chemical changes are often di�cult or impossible to reverse. Chemical properties describe a

substance's ability to undergo chemical changes, such as pH and �ammability. You'll study physical and chemical properties in more depth later

in this lesson.

Physical versus Chemical Changes

Extensive properties depend on the substance's amount. Some examples of extensive properties are weight and volume. Intensive properties

don't depend on the amount of matter. Temperature is an example of an intensive property. If you measure the temperature of water in a bucket,

the water's temperature doesn't depend on the amount of water in the bucket, so this is an intensive property. However, the volume of water in the

bucket does depend on how much water is in the bucket and is an extensive property.

Quantifying Matter

Physical properties can be categorized as qualitative or quantitative. Consider how an object can be described physically. For example, you get a

cup of co�ee on your way to class. You may note that the cup is �lled to the eight-ounce mark. You may also notice the color of the co�ee, that it's

hot, or that it smells good. �ese are all ways to describe the physical properties of the co�ee. Its volume, eight ounces, is a quantitative description

of the co�ee. If you measure the temperature of the co�ee as 180 °F this would also be a quantitative description. �e color, aroma, and labeling of

the liquid "hot" are examples of qualitative descriptions.

Quantitative descriptions use numbers to express properties and may use these instruments to make quantitative measurements: rulers, scales,

thermometers, timers, or graduated cylinders. Qualitative descriptions don't use numbers but are impressions produced by your senses.

Qualitative and quantitative properties are used to describe observations. For example, if you make pancakes, the physical changes to the batter

after it's cooked can be described in several ways.

Qualitative observations:

Comparing the odor of the batter to that of the cooked pancakes

Noting the color change from yellow to brown

Describing the texture change from a gooey liquid to a soft, �u�y solid

Quantitative observations:

Measuring the temperature of the batter versus that of the warm pancakes

Comparing the density of the batter with that of the cooked cakes by measuring the weight and volume of each

Note that while these are physical properties, cooking the pancakes also causes a chemical change.

Pancake batter and a stack of cooked pancakes

Weight is a common quantitative measurement. Yet it's important to distinguish between mass and weight. �e mass of an object is a measure of

the amount of matter in it. Although weight is related to mass, it isn't the same thing. Weight refers to the force that gravity exerts on an object,

which is directly proportional to the object's mass. �e weight of an object changes as the force of gravity changes, but the object's mass doesn't

change. An astronaut's mass doesn't change on the moon, but her weight on the moon is only one-sixth of her Earth weight. �is is because the

moon's gravity is only one-sixth of Earth's. She may feel "weightless" when she experiences the moon's lower gravity, but she can never be

"massless."

�e mass of an object is typically determined by using a balance to compare its mass with a standard mass, commonly called "weighing" an object.

An object's mass can also be determined by measuring the force required to accelerate the object, but this is a more di�cult method for

determining an object's mass. It takes much more force to accelerate a car than a bicycle because of the car's greater mass.

�e di�erence between mass and weight may not seem important now, but its signi�cance will become apparent as your studies progress. Keep in

mind that di�erent units are used to measure mass and weight because weight is a force that gravity exerts on a mass.

Measurements in science are generally made using the International System of Units, commonly referred to as the SI system of measurement.

You've worked with SI units in other classes.

Typical SI units of measure in science and math include the following:

Grams or kilograms for weight

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Meters or centimeters for length

Liters for volume

Degrees Celsius for temperature

Weight (a force) is measured in a unit called a Newton in the SI system. Density, the ratio of a substance's mass to volume, is most often measured

in grams per cubic centimeter (g/cm ).

In the United States, most measures are given in the American Customary Units system, sometimes called the English or American system. Typical

American measures for length are inches, feet, yards, and miles. For volume, ounces, cups, pints, quarts, gallons, and barrels are used. You're used

to expressing weight in pounds, but a pound is a measurement unit for mass. In the American system, the pound-force is the correct unit for

weight. You may feel reluctant to embrace the SI system, but you'll better understand it after experiencing the time-consuming and mundane work

of performing conversions from the American system.

Footnotes

1. Attribution: Flowers, Paul, Klaus �eopold, Richard Langley, and William R Robinson, Ph.D. Chemistry 2e. Houston, TX:

OpenStax2019. https://openstax.org/books/chemistry-2e/pages/1-introduction

3

Using the Periodic Table

�e periodic table is an incredibly useful tool. Once you know how it's arranged, you can predict many physical and chemical properties of

elements.

Refer to the interactive periodic table and select "trends" to see examples of this.

Previously in this section, you were introduced to the arrangement and parts of the table. As mentioned, elements in the periodic table are

arranged by increasing atomic number, from left to right and top to bottom. Elements have common properties and predictable trends by column,

row, and block. �ese trends are due to the atomic number and the arrangement of electrons.

The Periodic Table and Atomic Structure

Elements in the same period have the same principal quantum number. Recall from the last lesson that the principal quantum number, or shell

number, indicates the energy level of an atom. Elements in the �rst period, for example, have electrons in the �rst shell. Elements in the second

period have electrons in the �rst and second shells. Elements in the third period have electrons in the �rst, second, and third shells, and so on.

Periodic table showing electron shells.

Periodic trends are patterns in the properties of elements that are revealed by the periodic table. Some periodic trends occur as you move across a

row. As you move from left to right across a period, the number of protons in the nucleus of elements increases because the atomic number

increases, and the number of electrons increases accordingly. For example, in period two, lithium has three protons in its nucleus, beryllium has

four, boron has �ve, and so on. �e number of neutrons is roughly the same as the number of protons but varies in di�erent isotopes. Recall that

protons and neutrons make up the nucleus. �erefore, the size of an element's nucleus increases as you move across a period from left to right.

Atomic mass generally increases as you move from left to right across a row also. Since the number of neutrons can vary for a given element

because of di�erent isotopes of the element, there are a few instances where atomic mass decreases as you move from left to right on the table.

You'll learn about other periodic trends as the course progresses.

Groups on the periodic table also have common properties because elements in a group have the same number of electrons in their outermost

shell. While atomic number increases by only one as you move from left to right within a period, you should note that the atomic number

increases by certain intervals as you move down a column. �is is a result of how electrons are arranged in each element, which you'll learn more

about later in the course.

�e grouping of elements as main-group elements, transition elements, and inner transition elements is based on how their electrons populate

subshells. �is will also be discussed in depth later. �e main group elements have outer shell electrons in s- and p-orbitals, transition elements

have outer shell electrons in d-orbitals, and inner transition elements have outer shell electrons in f-orbitals.

Periodic table showing the orbitals that the outer electrons populate in main group, transition, and inner transition elements.

Metals, Nonmetals, and Metalloids

Most of the elements in the periodic table are classi�ed as metals, including iron, chromium, sodium, potassium, and tin. Except for mercury,

metals are solids at room temperature. Metals tend to be shiny, malleable (can be bent or pounded into sheets), and ductile (can be stretched into

thin wires). �ey have relatively high melting and boiling points and high densities, combine readily with oxygen or water to form corrosion

products and conduct electricity well. Di�erent metals display these properties to di�ering degrees. Metals tend to lose electrons in chemical

reactions. Metallic character decreases as you move across a period from left to right and increases as you move down a group.

Periodic table showing the orbitals that the outer electrons populate in main group, transition, and inner transition elements.

Nonmetals are found on the right side of the periodic table. Oxygen, bromine, carbon, and iodine are examples of nonmetals. Nonmetals can be

solids, liquids, or gases at room temperature. �ey tend to be less dense and more brittle than metals, with lower melting and boiling points.

Nonmetals are not good conductors of electricity. �ey tend to gain electrons in chemical reactions. �ere are far fewer nonmetals than metals.

Hydrogen, the most abundant element in the universe, is a nonmetal. Carbon, also a nonmetal, is an important element for life. One form of the

nonmetal phosphorus is highly reactive and spontaneously ignites in air, which is why it's shown in water in the �gure.

Chromium is an example of a metal.

Phosphorus is an example of a nonmetal.

Boron is an example of a metalloid.

Metalloids have properties between those of metals and nonmetals. �ese elements run in a diagonal line between nonmetals and metals in the

periodic table. Metalloids have fewer members than metals or nonmetals: boron, silicon, germanium, arsenic, selenium, antimony, tellurium,

polonium, and astatine. Some sources even disagree on whether antimony, astatine, and polonium are metalloids. Metalloids tend to be fair

conductors of electricity, brittle, and less dense than metals. Because they're brittle, metalloids are not malleable or ductile. Metalloids may gain

or may lose electrons in chemical reactions. �ey're commonly used to make semiconductors in the electronics industry.

You were introduced to quantum numbers in the last section. In the next lesson, you'll see how electrons �ll orbitals, how to express electron

con�gurations using quantum numbers, and how this is re�ected in element properties and trends in the periodic table.

The Periodic Table History and Structure

Scientists long ago recognized that groups of elements had similar properties. Some elements existed naturally as gases, some were metals, some

conducted electricity, and so on. Scientists knew the value of organizing phenomena systemically according to their physical and chemical

properties, and several attempted to create such a system for the elements. As early as 1789, Antione Lavoisier grouped elements into broad

categories of gases, metals, earths, and nonmetals. German scientist Johann Dobereiner proposed a "law of triads" in 1829. Dobereiner had

noticed groups of three elements with similar physical properties. For example, the triad of lithium (Li), sodium (Na), and potassium (K) each

react with chlorine (Cl) in a 1:1 ratio to form lithium chloride (LiCl), sodium chloride (NaCl), and potassium chloride (KCl). Li, Na, and K all react

readily with water at room temperature. Dobereiner found that the atomic weight and density of the middle element of each triad could be closely

predicted by the arithmetic mean of the other two elements. In the case of Li, Na, and K, the atomic weight of Na (22.99) is very close to the mean

(23.02) of that of Li (6.94) and K (39.10).

Another of Dobereiner's triads is beryllium (Be), magnesium (Mg), and calcium (Ca). Each member of this triad reacts with Cl in a 1:2 ratio to form

BeCl , MgCl , and CaCl . Be has an atomic weight of 9.01, Mg has an atomic weight of 24.31, and Ca has an atomic weight of 40.08. �e mean of the

atomic weights of Be and Ca is 24.55, very close to the actual atomic weight of Mg. While not all the known elements could be arranged in triads,

Dobereiner's approach was useful because it identi�ed groups of elements with similar properties and revealed an orderly pattern in some of their

physical and chemical properties. �e concept of triads related the properties of an element to its atomic weight.

In 1865, English chemist John Newlands ordered the elements by their atomic weights and recognized that properties of elements seemed to

repeat at regular intervals. Newlands identi�ed seven "families" of elements and observed that there were seven elements between elements with

similar properties. Based on this, he created a table based on musical octaves. �is table placed some dissimilar elements together and didn't

account for new elements.

Mendeleev's Periodic Table

In 1869, Dimitri Mendeleev of Russia created a table that listed elements according to their atomic weights and their chemical properties. Around

the same time, Lothar Meyer of Germany suggested a similar arrangement but focused on physical properties rather than chemical properties.

Mendeleev's table proved to be superior, and he is now recognized as the creator of the periodic table of elements.

Mendeleev had the foresight to use his table to predict the existence of elements that had not yet been discovered and left blank spaces in his table

where he believed unknown elements existed. He even predicted the properties of these elements. Over time, his predictions were validated.

Mendeleev believed some of the atomic masses accepted at the time were incorrect. He placed these elements according to their properties rather

than their atomic masses, and some of their atomic masses were incorrect. Other elements' atomic masses were proved to be correct. In 1913, the

English scientist Henry Moseley developed the concept of atomic number. Arranging elements by atomic number rather than by atomic mass

further validated Mendeleev's arrangement of the elements. While not all elements were known at Mendeleev's time, the basic arrangement of his

table endured.

In the modern version of the periodic table, shown in the �gure, elements are listed by the atomic number. �e atomic number is now recognized

as the primary identi�er of an element. Mendeleev's law, also known as the periodic law, states that properties of the elements are periodic when

the elements are arranged according to their atomic number. Periodicity means that something occurs at regular intervals and describes how

these properties follow patterns. �e patterns in properties are due to the way that electrons are arranged in the outer shells of elements. For

example, elements in the �rst column of the periodic table have only one electron in their outermost shell. �is will be discussed further in the

next lesson—after you learn more about electron con�guration.

2 2 2

e periodic table of the elements.

The Modern Periodic Table

While they were not always called "elements," carbon, iron, silver, gold, sulfur, tin, lead, and mercury have long been known and used by humans.

Over time, more elements have been isolated and named. By the time Mendeleev organized the known elements into his periodic table, 63

elements had been identi�ed, including some that were radioactive. More elements were discovered over time, and today there are 118 elements

in the periodic table. Of these, 94 occur naturally and 24 are synthetic, made in particle accelerators or nuclear reactors. You'll learn more about

synthetic elements in future lessons on nuclear chemistry.

Each square in the periodic table represents an element. �e element's symbol, atomic number, and average atomic mass are listed. Some

versions of the table also detail the electron con�guration of each element, shown in the upper right corner of each element square in the �gure.

Elements are listed in order of increasing atomic number from left to right and top to bottom. Recall that the atomic number is the number of

protons in an element's nucleus. �e number of electrons in a neutral atom is equal to the number of protons, so as the atomic number increases,

the number of electrons also increases.

�e table has seven horizontal rows called periods and 18 vertical columns called groups. Less often, a column is called a family and a row is

called a series. Columns are labeled with numbers one through 18. Sections of two rows (elements 57-71 and 89-103) are typically removed and

placed beneath the main body of the table so that it more easily �ts on one page and is less confusing. �ese are the lanthanides and the

actinides. Some versions of the periodic table use Roman numerals, shown in the �gure below the column numbers. �ere are common

properties and trends in the periodic table by column, row, and block.

One way to divide the periodic table by blocks is to classify elements as metals, nonmetals, or metalloids. Metals tend to be shiny, good

conductors of electricity, malleable (bendable) and ductile (can be drawn into wires), and conduct heat and electricity well. Nonmetals, in

contrast, often appear dull and are poor conductors of heat and electricity. Metalloids have some properties of metals and some of nonmetals, and

they conduct heat and electricity moderately well. In the table below, metals are shown in gray, nonmetals in green, and metalloids in yellow to

help you visualize where they're located on the periodic table. Most of the known elements are metals. Note that in this version of the periodic

table, helium has been moved next to hydrogen and the lanthanides and actinides have been placed in the main body of the table rather than

below it.[1]

A simple periodic table of the elements indicating which elements are metals, nonmetals, and metalloids.

Another way elements can be classi�ed is as main-group elements, transition elements, or inner transition elements. Main-group elements are

sometimes also called representative elements and consist of the elements in columns 1, 2, and 13-18. Transition elements are sometimes called

transition metals and are the elements in columns 3-12. Inner transition elements are sometimes referred to as inner transition metals and are the

lanthanides and actinides mentioned earlier. �ese two rows of elements are typically removed from the main table and placed underneath the

rest of the table. �e lanthanides are in the top row and the actinides are in the bottom row. In the next lesson, you'll see that these blocks are

de�ned by the suborbital that an element's outermost electrons occupy. Again, in this version of the periodic table, helium is located next to

hydrogen and the inner transition elements have been placed in the main body of the table rather than below it.

A simple periodic table of the elements indicating which elements are main group elements, transition elements, and inner transition elements.

�e elements in each column or group of the periodic table have similar properties. Some groups have special names. For example, when forming

molecules, the elements in group 1 form compounds that consist of one atom of the element and one atom of hydrogen. Except for hydrogen, the

elements in group 1 are known as alkali metals. �e elements in group 2 form compounds consisting of one atom of the element and two atoms

of hydrogen. �ere are known as alkaline earth metals. Other groups with speci�c names are the halogens in group 17 and the noble gases in

group 18, also known as inert gases. �e noble gases are typically found as elements rather than compounds because they have very low reactivity

and generally don't bond with other elements. Noble gases are generally colorless and odorless. Elements in the same group are congeners.

Hydrogen is a unique nonmetal with properties similar to both group 1 and group 17 elements. For that reason, hydrogen may be shown at the top

of either of these groups or by itself.

Looking at the periodic table, you might have noticed that the atomic masses of some elements are given in square brackets. �is is true of element

43 (technetium), element 61 (promethium), and most of the elements with atomic number 84 (polonium) and higher. �is indicates that the

elements are unstable, radioactive isotopes, which you'll learn more about in the nuclear chemistry chapter. An average atomic weight can't be

determined for these elements because their radioisotopes vary signi�cantly in relative abundance, depending on the source, or may not even

exist in nature. �e number in square brackets is the atomic mass number (and approximate atomic mass) of the most stable isotope of that

element.

Visit this page for an interactive periodic table, as well as detailed information, videos, and podcasts about each element.

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The Electron Cloud Model

With the acceptance of wave-particle duality, Bohr's model of the atom had to be modi�ed. Austrian scientist Erwin Schrödinger did this by

extending de Broglie's work. Schrödinger derived what's today known as the Schrödinger equation. By applying this equation to hydrogen-like

atoms, he was able to reproduce Bohr's expression for atomic energy. Schrödinger described electrons as waves in his functions, represented by

the Greek letter psi, ᴪ. A few years later, the German scientist Max Born proposed an interpretation of an electron's wave function ᴪ that's still

accepted today: Wave functions can be used to describe the probability of an electron existing in a certain location. In other words, you can use

wave functions to determine the distribution of electrons with respect to the nucleus in an atom. Schrödinger, Heisenberg, Einstein, Planck, and

other scientists de�ned a new branch of physics to describe the behavior of matter at the subatomic scale, generally referred to today as quantum

mechanics.

In Bohr's model, an electron was visualized as a particle orbiting a nucleus in a path, much like the planets orbiting around the sun. Schrödinger

and Born re�ned this idea by instead describing the probability of �nding an electron at di�erent locations. Instead of a �xed point in space, an

electron can be pictured as existing in a charged cloud. �e cloud is denser where there's a higher probability of �nding the electron. �e cloud is

less dense in places where there's a low probability that the electron will be found. �e �gure shows an orbital for hydrogen, which has only one

electron. An orbital can be thought of as a three-dimensional picture that shows the probability of where you'll �nd an electron.

Comparison of the Bohr model and the electron cloud model

Orbitals are described with three quantum numbers. Recall that the energy of an electron in an atom is quantized—it has speci�c values and can

jump from one energy level to another but not transition smoothly or stay between these levels. �ese energy levels are labeled with n values, n =

1, 2, 3, and so on, called the principal quantum number or the shell number. In general, the energy of an electron is greater with greater values of

n. You can imagine shells as concentric circles radiating out from the nucleus. Electrons in a speci�c shell are most likely to be found within the

corresponding spherical volume. �e farther away from the nucleus that an electron is located, the higher the shell number will be and the higher

the energy level will be.

In the diagram, the circle closest to the nucleus is the n = 1 shell. �e next closest circle is the n = 2 shell, and so on. �e number of electrons in

each shell and how electrons form chemical bonds will be discussed later.

[1]

Electron shell diagram

Electron distribution in shells for a few elements is shown in the next �gure.

�e principal quantum number, n, is one of three quantum numbers used to characterize an orbital. �e other quantum numbers are 𝑙 the

angular momentum quantum number, and m , the magnetic quantum number. Just as n de�nes the general energy of the orbital, the angular

momentum quantum number speci�es the orbital shape, and the magnetic quantum number speci�es the spatial orientation of the orbital.

Angular momentum quantum numbers are integers from 0 to n – 1. An orbital with n = 1 only has 𝑙 = 0, but n = 2 permits 𝑙 = 0 and 𝑙 = 1, and so on.

Orbitals with the same value of 𝑙 de�ne a subshell.

Orbitals with 𝑙 = 0 are called s orbitals and make up the s subshells. �e value 𝑙 = 1 corresponds to p orbitals. �e orbitals with 𝑙 = 2 are called the

d orbitals, followed by the f-, g-, and h-orbitals for 𝑙 = 3, 4, and 5.

�e s subshell is spherical, and the p subshell has a dumbbell shape. �e d and f orbitals are more complex. Remember that these shapes

represent the three-dimensional regions within which electrons are likely to be found.

l

Shapes of the s, p, and d orbitals

�e magnetic quantum number (m ) speci�es the relative spatial orientation of a particular orbital. �e m can be an integer from –𝑙 to 𝑙, so the

total number of possible orbitals with the same value of 𝑙 (that is, in the same subshell) is 2𝑙 + 1.

�ere's one s orbital in an s subshell (𝑙 = 0). . .

. . .three p orbitals in a p subshell (𝑙 = 1). . .

. . .�ve d orbitals in a d subshell (𝑙 = 2). . .

. . .seven f orbitals in an f subshell (𝑙 = 3), and so forth.

�e shape of an s orbital and the shapes and orientations of the three p orbitals are shown in the �gure. Orientations of d, f, and higher subshells

are more complicated and will be presented later.

l l

e spherical shape of an s orbital and the shape and orientations of the three p orbitals.

�e next �gure shows what an s orbital and three p orbitals together look like.

An s orbital and three p orbitals in an atom.

IDENTIFYING ORBITAL COMPONENTS

�e principal quantum number de�nes the electron energy.

�e angular momentum quantum number determines the orbital shape.

�e magnetic quantum number speci�es the spatial orientation of the orbital.

An orbital is usually described with the principle quantum number as a number and the secondary quantum number as a letter, such as 1s, 3p, or

4d. �e third quantum number is given as a subscript; however, it isn't always necessary, so it isn't always listed.

One last piece of information used to describe electrons in orbitals is spin. Electron spin is a di�erent kind of property that quantum number. Each

electron acts like a tiny magnet. You can imagine electrons spinning around like tops. When two electrons are in the same orbital, each spins in a

direction opposite to the other—one clockwise and the other counterclockwise. �e magnitude of electron spin only has one value, and an

electron can only "spin" in one of two quantized states. Spin is described by convention as either ±1/2.

Quantum numbers help us to visualize the distribution of electrons in an atom. �ey're also a way to understand the relative energies of electrons

within an atom due to their positions relative to the nucleus. Now that you've been introduced to quantum numbers, you'll later learn about how

the electrons �ll these shells and subshells. Electron number and distribution determine many of an element's properties and de�ne how

elements form chemical bonds. You'll �rst need to learn more about how elements are arranged in the periodic table.

[1]

Footnotes

1. Attribution: Flowers, Paul, Klaus �eopold, Richard Langley, and William R Robinson, Ph.D. Chemistry 2e. Houston, TX:

OpenStax2019. https://openstax.org/books/chemistry-2e/pages/1-introduction

Electrons: Particles or Waves?

Initially, the patterns of atomic spectral lines were puzzling. In 1885, Johann Balmer, a Swiss scientist, came up with a formula that �t the data for

these lines well. But the formula didn't explain why the lines occurred. Scientists also studied what happened when solid materials were heated

until they glowed and emitted spectra. Data from these experiments didn't �t with previous views of energy. Max Planck, a German scientist

studying radiation from solids, resolved the problems by proposing that energy emitted from substances could only be released in discrete

"packets." Planck described the smallest quantity of energy that could be released as a quantum and said that energy could only be released as

multiples of this quantum, called quanta. �is was a revolutionary idea. Planck's equation for a quantum of energy is where λ is the

wavelength of the radiant energy, c is the speed of light, and h is a constant (6.63 × 10 Js). �e h was named Planck's constant in his honor.

�e next breakthrough came when German scientist Albert Einstein applied Planck's theory to the photoelectric e�ect, a problem puzzling

scientists at the time. �e photoelectric e�ect was observed when electromagnetic radiation hit a metal surface and the surface emitted electrons.

Electron emission depends on the radiation frequency. No electrons are emitted at low frequencies. Frequency is increased until the radiation

reaches a certain "threshold" frequency, where emission occurs. As the frequency is increased above the threshold frequency, the number of

electrons doesn't increase but the energy of the emitted electrons does. Einstein proposed that energy exists as quanta called photons. He used

Planck's constant to describe the energy of a quantum, E = hv, where E is the energy of a quantum of light or a photon, h is Planck's constant, and νis frequency. �is implies that light is made of bundles of energy—that it behaves as a particle. A photon is this discrete particle of light.

To visualize this in the photoelectric e�ect, imagine a photon as a particle of light hitting an electron in an atom. An electron is held by an atom

with a certain amount of energy, the electron binding energy. For the electron to be released from the atom, the photon must have enough energy

to overcome the electron binding energy. If the photon doesn't have enough energy, it just bounces o� the electron. If it has just enough energy to

overcome the electron binding energy, the electron is knocked free, and this de�nes the threshold frequency. If a photon has more than enough

energy, the electron is released, and the extra energy is transferred to the electron as kinetic energy.

For explaining the photoelectric e�ect, Einstein's equation �t scientists' observations, but it caused quite a problem. �e question of the nature of

light was again debated. Was light best described as a wave or a particle?

e photoelectric e�ect

E =hc

λ

-34

Previously, you learned about the evolution of atomic theory. Recall that Rutherford's model improved on previous models but wasn't compatible

with classical mechanics. �e simplest atom is hydrogen, consisting of a single proton as the nucleus, around which a single electron exists in an

orbital. �e electrostatic force, which attracts the negatively charged electron to the positively charged proton, depends on the distance between

the two particles. An electron in an elliptical orbit would be accelerating by changing direction. According to classical electromagnetism, the

electron should also continuously emit electromagnetic radiation. �is loss in orbital energy should result in the electron's orbit getting

continually smaller until it spirals into the nucleus. Rutherford's model implies that atoms are inherently unstable.

Niels Bohr attempted to resolve the atomic paradox in his model by ignoring classical electromagnetism's prediction that the orbiting electron in

hydrogen would continuously emit light. Instead, he combined Planck's ideas of quantization with Einstein's �nding that light consists of photons

whose energy is proportional to their frequency. Bohr assumed that the electron orbiting the nucleus would not normally emit any radiation, but it

would emit or absorb a photon if it moved to a di�erent orbit. Instead of allowing for continuous values of energy, Bohr assumed the energies of

these electron orbitals were quantized, or had discrete values. Each of these energy levels that an electron could exist in was called a stationary

state. �e stationary state with the lowest energy was called the ground state and given the designation n = 1. Higher states are called excited

states and are designated n = 2, n = 3, and so on.

A fundamental law of physics is that matter is most stable with the lowest possible energy. �e electron in a hydrogen atom usually moves in the n

= 1 orbit, the orbit in which it has the lowest energy. When the electron is in this lowest-energy orbit, the atom is in its ground state. If the atom

receives energy from an outside source, it's possible for the electron to move to an orbit with a higher n value, making the atom exist in an excited

state with higher energy. Bohr said that the energy di�erence when an electron moved from one state to another was ΔE = hv. In this formula, ΔE is

the change in energy, v is the frequency of a photon emitted or absorbed, and h is Planck's constant.

When an electron transitions from an excited state (higher energy) to a less excited state, or the ground state, the di�erence in energy is emitted as

a photon. Similarly, if a photon is absorbed by an atom, the energy of the photon moves an electron from a lower energy state to a more excited

state. Atoms don't absorb or emit energy when they're in a stationary state. .[1] [1]

e Bohr model, showing the release of a photon as an electron moves to a lower state.

Electrons exist only at particular energy levels, and they emit or absorb photons as they move between energy levels. �is idea explains the speci�c

wavelengths of light observed in atomic emission spectra. �ere's a speci�c energy change associated with an electron moving from one energy

level down to the next. �e photon that it emits will correspond to this energy change, . Similarly, a �xed amount of energy is

required for an electron to move from one level up to another. Higher energy levels aren't as stable as lower energy levels, so the electron will drop

back to the lower state and release a photon. Two of the quantities in this equation, h and c, are constants, so ΔE is inversely proportional to

wavelength. For an electron to move between two speci�c energy levels, a �xed wavelength of light will be emitted. �is energy and wavelength

will be di�erent for di�erent values of n and di�erent elements.

Watch the video, Bohr Model of the Atom, for an explanation of the energy of electron movement between levels.

�e idea that electrons "jump" between states isn't compatible with a classical, Newtonian view of physics. However, Bohr's model explained the

observed behavior of hydrogen atoms. Additionally, it was possible to measure the energy changes between stationary states using the atomic

spectrum of hydrogen.

Δ = hv =hc

λ

Bohr's model of the hydrogen atom provided insight into the behavior of matter at the atomic level, but it didn't account for interactions between

electrons in atoms with more than one electron. It did introduce several important ideas to describe the distribution of electrons in an atom.

In an atom, the energies of electrons, or energy levels, are quantized: they're described by quantum numbers. �ese are integers with

speci�c values and can characterize the arrangement of electrons in an atom (n = 1, n = 2, n = 3, and so on).

An electron's energy increases with increasing distance from the nucleus.

�e discrete lines in emission spectra of elements correspond to quantized electron energies.

�e most important of these ideas involves the quantized energy levels for an electron in an atom. �is model laid the foundation for the quantum

mechanical model of the atom. Bohr won the Nobel Prize in Physics in 1922 for his contributions to scientists' understanding of atomic structure

and line spectra emissions.

Wave-Particle Duality

Planck, Einstein, Bohr, and others proposed theories in which electromagnetic radiation acted as both waves and particles. �e idea that

something can have characteristics of both a wave and a particle is called wave-particle duality.

In 1924, the French scientist Louis de Broglie proposed that if light can have particle characteristics, particles also have wave characteristics. He

derived an equation, , to calculate the wavelength of an electron. In this equation, λ is wavelength, h is Planck's constant, m is mass, and v

is velocity. Up to this point, an electron had been regarded as a particle. Experiments showed that electrons do indeed have an associated

wavelength.

A surprising implication of de Broglie's equation is that matter has both wave-like and particle-like characteristics, and even macroscopic objects

have wavelengths. For large objects, this wavelength is so small it isn't noticeable. For example, the wavelength of a 60 kg person moving at 1 m/s

can be calculated as follows:

�is wavelength, 1.1 × 10 m, is so short that it's unobservable. Since Planck's constant is so small, this wavelength is relevant only for extremely

small masses, such as subatomic particles.

Notice that in the equation above, joule is written out in di�erent units. A joule is equal to a newton of force applied over a meter. A newton is the

force needed to accelerate one kilogram by one meter per second squared. �erefore, a joule is equal to a kg m /s . By expanding the joule to this

form, you can see that many of the units cancel each other out and leave only the meter, which is the appropriate unit for wavelength.

With the acceptance of wave-particle duality, electrons needed to be described as both particles and waves. �e German scientist Werner

Heisenberg studied the implications of the wave-particle duality and formulated the uncertainty principle. �is principle states that it's

impossible to know both the exact momentum and position of an object at the same time. Any method used to �nd the momentum of a particle

changes its position and vice versa. Heisenberg's uncertainty principle helps scientists understand the limits of what can be known about

subatomic particles, and it was critical in shaping the modern view of electrons in atoms.

Footnotes

1. Attribution: Flowers, Paul, Klaus �eopold, Richard Langley, and William R Robinson, Ph.D. Chemistry 2e. Houston, TX:

OpenStax2019. https://openstax.org/books/chemistry-2e/pages/1-introduction

[1]

λ = hmv

λ = = =

= 1.1 × 10−35 m

h

mv

6.6 × 10−34Js

(60 kg) (1 )ms

6.6 × 10−34 kg m2 s

s2

60kg m

s

–35

2 2

Electromagnetic Waves

In the last lesson, you learned that atoms are made of protons, neutrons, and electrons, and you were introduced to some characteristics of these

particles. Recall that protons are positively charged, neutrons have no charge, and electrons have a negative charge. Protons and neutrons make

up the small, dense nucleus of an atom, which takes up a very small portion of the volume of an atom. Electrons exist in orbitals surrounding the

nucleus, and these orbitals where electrons can be found take up a much larger portion of the volume of the atom compared to the nucleus.

Protons and neutrons are similar in weight and are largely responsible for the mass of an atom. Electrons weigh very little compared to protons

and neutrons and contribute little to the mass of an atom.

You also learned that an ion was an atom that has gained or lost electrons. �e distribution of electrons in atoms determines how atoms combine

and chemical bonds are governed by the nature of electrons in atoms. In this section, you'll learn more about electrons and their role in atoms.

First, you'll learn a little about electromagnetic radiation, particles, and waves.

The Electromagnetic Spectrum

You may not have thought much before about the nature of light beyond light enabling you to see things. You "see" objects through this process:

1. Light re�ects o� of an object.

2. �e light is sensed by the rods (for black and white) and cones (for color) on the retinas of your eyes.

3. It's changed into an electrical signal.

4. �e electrical signal is sent to your brain, where it's interpreted as an image.

How we see things

Light is a type of radiant energy, but there are other types of radiant energy as well. A microwave oven uses radiant energy to cook things. X-rays

and radio waves are radiant energy. You may use sunscreen when outside to protect yourself from ultraviolet rays, yet another type of radiant

energy.

Scientists have wondered about the nature of light for centuries. Around the seventeenth century, there were two competing views. Isaac Newton's

investigations of light led him to propose that light was composed of small particles. Another scientist, Christian Huygens, theorized that it was a

wave. In the nineteenth century, Scottish scientist James Maxwell discredited the particle view of light. Maxwell's theory of electromagnetic

radiation showed that light was the visible part of a vast spectrum of electromagnetic waves. Di�erent types of electromagnetic radiation,

including light, x-rays, and ultraviolet radiation, make up the electromagnetic spectrum. You'll explore the entire spectrum later in this section.

By the end of the nineteenth century, scientists viewed the physical universe as comprising two separate domains:

1. Matter composed of particles moving according to Newton's laws of motion

2. Electromagnetic radiation consisting of waves governed by Maxwell's equations

Today, these domains are referred to as classical mechanics and classical electrodynamics, or classical electromagnetism. �is framework didn't

explain some physical phenomena, but scientists were con�dent that these phenomena would ultimately be resolved within this framework.

However, as you'll learn, these phenomena led to a contemporary framework that has superseded the classical view by connecting particles and

waves at a fundamental level. Visible light and other forms of electromagnetic radiation play important roles in chemistry because they can be

used to infer the energies of electrons within atoms and molecules.

A wave is an oscillation or periodic movement that can transport energy from one point in space to another. Common examples of waves are all

around you:

Shaking the end of a rope transfers energy from your hand to the other end of the rope.

Dropping a pebble into a pond causes waves to ripple outward along the water's surface.

�e expansion of air that accompanies a lightning strike generates sound waves, called thunder, that can travel for several miles.

In each of these cases, kinetic energy is transferred through matter (the rope, water, or air) while the matter essentially remains in place.

Waves don't require matter for travel. Maxwell showed that electromagnetic waves consist of an electric �eld oscillating in step with a

perpendicular magnetic �eld. Both the electric �eld and the magnetic �eld are perpendicular to the direction of travel. �ese electromagnetic

waves can travel through a vacuum at a constant speed of 2.998 × 10 m/s. �is is the speed of light and is denoted c.8

Electromagnetic waves are a combination of electric and magnetic �eld waves in phase with each other. ey travel perpendicular to the direction of these �elds.

Previously, you learned the di�erence between kinetic and potential energy. Recall that energy is the ability to do work, kinetic energy relates to

the motion of objects and the work they're capable of due to their motion, and potential energy relates to an object's position in a �eld, such as a

magnetic or gravitational �eld. �e previous �gure showed a magnetic and an electrical �eld, and waves have motion. You may be wondering if

electromagnetic waves are potential or kinetic energy. An electromagnetic wave is moving, so it conveys energy and is considered kinetic. You can

also consider the fact that when electromagnetic energy hits matter, at least some of it is converted to heat, seen when you microwave food. Heat is

a measure of kinetic energy. �is may make more sense to you later in the section after you learn about photons.

All waves, including forms of electromagnetic radiation, are characterized by three things:

1. A wavelength (λ, the lowercase Greek letter lambda)

2. A frequency (ν, the lowercase Greek letter nu)

3. An amplitude

As can be seen in the �gure, the wavelength is the distance between two consecutive peaks or troughs in a wave. Wavelength is measured in

meters in the SI system. Electromagnetic waves have an enormous range of wavelengths and are grouped according to their wavelengths.

Wavelengths of kilometers (10 m) to picometers (10 m) have been observed. Frequency is the number of wave cycles that pass a speci�ed point

in space in a speci�ed amount of time (measured in seconds). A cycle corresponds to one complete wavelength. �e unit for frequency, expressed

as cycles per second [s ], is the hertz (Hz). Common multiples of this unit are megahertz (1 MHz = 10 Hz) and gigahertz (1 GHz = 10 Hz). �e

amplitude of a wave corresponds to the magnitude of the wave's displacement. �e amplitude is half the height between the peaks and troughs.

Amplitude is related to the intensity of the wave, which is the brightness of light and the volume of sound.

Parts of a wave

In general, the product of a wave's wavelength (λ) and its frequency (ν), λν, is the speed of the wave. For electromagnetic radiation in a vacuum,

speed is equal to the fundamental constant c, the speed of light.

c = λν

Wavelength and frequency are inversely proportional. As wavelength increases, frequency decreases. �e �gure shows the electromagnetic

spectrum, the range of all types of electromagnetic radiation. Visible light makes up only a small portion of the electromagnetic spectrum. Notice

that di�erent parts of the electromagnetic spectrum are typically referred to by a variety of units. �is is because the technologies operating in

various parts of the electromagnetic spectrum developed at di�erent times. For example, radio waves are usually speci�ed as frequencies

(typically in MHz) while the visible region is usually speci�ed in wavelengths (typically in nm or angstroms).

3 −12

−1 6 9

e electromagnetic spectrum

Many valuable technologies operate in the radio frequency region (3 kHz–300 GHz) of the electromagnetic spectrum. At the low-frequency, long-

wavelength end of this region are AM (amplitude modulation) radio signals (540–2830 kHz) that can travel long distances. FM (frequency

modulation) radio signals are used at shorter wavelengths, corresponding to higher frequencies (87.5–108.0 MHz). In AM radio, the information is

transmitted by varying the amplitude of the wave. In FM radio, by contrast, the amplitude is constant and the instantaneous frequency varies. As

you can see, the same types of waves (such as radio waves) can di�er in amplitude and wavelength.

Other technologies also operate in the radio portion of the electromagnetic spectrum. Examples include the following:

4G cellular telephone signals are approximately 880 MHz.

Global positioning system (GPS) signals operate at 1.228 and 1.575 GHz.

Local area wireless technology (Wi-Fi) networks operate at 2.4 to 5 GHz.

Highway toll sensors operate at 5.8 GHz.

�e frequencies associated with these applications are convenient because these waves tend not to be absorbed much by common building

materials.

Atomic Spectrum of an Element

In the seventeenth century, Isaac Newton refracted sunlight with a prism and showed its component colors.

[1]

Light split into component colors with a prism

Later, scientists experimenting with light found that pure elements in the gas phase emitted light when heated. In contrast to the continuous

spectrum of colors that Newton saw, the spectra produced by heated gases had a relatively small number of narrow bands separated by dark areas,

termed spectral lines. �e following �gure shows the spectrum of lines emitted by hydrogen gas.

Spectral lines of hydrogen

Scientists found that each element investigated produced a unique pattern or "�ngerprint" of discrete spectral lines. �is pattern, now called an

atomic spectrum, can be used to identify an element. �e lines emitted by elements aren't just in the visible portion of the spectrum but are in the

infrared and ultraviolet regions as well.

Footnotes

1. Attribution: Flowers, Paul, Klaus �eopold, Richard Langley, and William R Robinson, Ph.D. Chemistry 2e. Houston, TX:

OpenStax2019. https://openstax.org/books/chemistry-2e/pages/1-introduction

Atomic Number and Mass

You learned about elements in earlier sections. Each element has an atomic number, which is the number of protons in its nucleus. �e atomic

number de�nes the identity of an atom. For example, an atom with one proton is hydrogen, and an atom with two protons is helium. If the atom is

neutral, it must have the same number of electrons and protons. If not, the atom will have a charge and is called an ion. Atoms can gain or lose

electrons in a variety of ways, which will be discussed later. Anions are atoms that gain electrons and have a negative charge. Cations are atoms

that lose electrons and have a positive charge.

Atomic mass number is the number of protons plus the number of neutrons in an atom. �e number of neutrons in an atom is the mass number

minus the atomic number.

As mentioned, isotopes are variations in the number of neutrons in an atom's nucleus. Common isotopes of hydrogen and helium are shown in

the �gure, along with the notation used to signify their mass number. Hydrogen most commonly has no neutron and only one proton in its

nucleus. However, isotopes of hydrogen exist, such as a hydrogen atom with one neutron that has a mass number of two. �e mass number of an

isotope is written as a superscript (2 in this example) before the element symbol.

Sometimes, you'll see the atomic number as a subscript before the element symbol, but this is often left o�. In text, this would be written as

"hydrogen-2." �e isotope of hydrogen with two neutrons in its nucleus has a mass number of three (two neutrons plus one proton), written in text

as "hydrogen-3." A few isotopes have common names, such as deuterium for hydrogen-2 and tritium for hydrogen-3.

Isotopes of hydrogen and helium. e most abundant isotope of each element is denoted by a green box.

�e occurrence and natural abundance of isotopes can be experimentally determined using an instrument called a mass spectrometer. Mass

spectrometry (MS) is used in chemistry, forensics, medicine, environmental science, and many other �elds to analyze and help identify the

substances in a sample of material. In a typical mass spectrometer, the sample is vaporized and exposed to a high-energy electron beam. �is

beam causes the sample's atoms or molecules to become electrically charged, typically by losing one or more electrons. �ese cations then pass

through a variable electric or magnetic �eld that de�ects each cation's path to an extent that depends on both its mass and charge. �is is similar

to how the path of a large steel ball bearing rolling past a magnet is de�ected to a lesser extent than that of a small steel BB. �e ions are detected,

and this makes a mass spectrum.

How a mass spectrometer works

A mass spectrum is a plot of the relative number of ions generated versus their mass-to-charge ratios. �e height of each vertical feature or peak

in a mass spectrum is proportional to the fraction of cations with the speci�ed mass-to-charge ratio. Since its initial use during the development of

modern atomic theory, MS has evolved to become a powerful tool for chemical analysis in a wide range of applications. [1]

A mass spectrum shows the relative abundance of di�erent Zr isotopes in a sample.

Watch this Simple Explanation of Mass Spectrometry to see how mass spectrometry works.

Atomic Mass

Atomic mass is the mass of an atom of a particular element. Atoms and the particles that comprise them are very small. To describe such small

objects, instead of using conventional units, special units are de�ned. An atomic mass unit (amu) is approximately 1.66 x 10 kg. An amu is

de�ned as one-twelfth the mass of an atom of carbon-12, the most common isotope of carbon. A proton has a mass of 1.0073 amu, a neutron has a

mass of 1.0087 amu, and an electron has a mass of about 0.00055 amu. Atomic weight, also called relative atomic mass, is the average mass of an

atom of an element, accounting for di�erent isotopes of the element and the relative abundance of these isotopes.

Since they're very small compared to protons and neutrons, electrons contribute little to the mass of an atom. Protons and neutrons have a mass of

around one amu each, so the mass number of an element gives you an approximate atomic mass for an atom. For example, carbon has an atomic

number of six (six protons) and an atomic weight of 12.011. �e most common isotope of carbon (carbon-12) has six neutrons. However, like

hydrogen and helium, carbon has several other common isotopes, such as carbon-13 and carbon-14. To get atomic weight, scientists use the

weight of these isotopes along with a percentage of how frequently they occur to determine the average weight of each element.

atomic mass = Σ(mass of isotope)(relative abundance)

For example, oxygen is a gas at room temperature, and it has three common isotopes. �ese are shown in the table here along with their relative

abundance. On average, a sample of oxygen contains 99.759% oxygen-16, 0.037% oxygen-17, and 0.204% oxygen-18. �ere are other isotopes of

oxygen, but they occur in much smaller percentages and therefore aren't signi�cant. �e mass of oxygen-16, determined with a mass

spectrometer, is 15.99491. As you would expect from the additional neutron, the mass of oxygen-17 is 16.99914 and the mass of oxygen-18 is

17.99916. To �nd the atomic weight of oxygen, multiply each isotope mass by its abundance, and then add these together.

Isotope Relative Abundance Mass of Isotope (amu) (% abundance)•(mass of isotope)

Oxygen-16 99.759% 15.99491 (15.99491)(0.99759) = 15.95636

Oxygen-17 0.037% 16.99914 (16.99914)(0.00037) = 0.00629

Oxygen-18 0.204% 17.99916 (17.99916)(0.00204) = 0.03672

Sum 15.9994 amu

-27

Atomic number and atomic weight are both given for each element in most representations of the periodic table. An example for oxygen is shown

in the �gure.

Footnotes

Attribution: Flowers, Paul, Klaus �eopold, Richard Langley, and William R Robinson, Ph.D. Chemistry 2e. Houston, TX:

OpenStax2019. https://openstax.org/books/chemistry-2e/pages/1-introduction

Discovering Atomic Structure

Over time, the idea that atoms are the smallest quantity of an element was accepted. However, the nature of atoms wasn't well understood. Many

important contributions and key discoveries led to the present understanding of the atom. J.J. �ompson is credited with discovering the �rst

known subatomic particle, the electron. �ompson made this discovery using a device called a cathode ray tube. �is tube contains almost no air,

and an electric current is passed through the tube, creating a "cathode ray," which is now understood to be a beam of electrons. Particles were able

to travel through the tube without bumping into air molecules and were detected at the end of the tube.

ompsons’s cathode ray tube apparatus

�ompson applied an electric �eld, shown in the �gure as blue "+" and red "-" plates, by connecting these plates to a battery. �is caused the beam

of particles to impact the disk higher when the plates were connected as shown in the �gure. �e impact was lower when the plates were reversed.

�ompson concluded that the particles hitting the disk were attracted to the positive plate and repelled by the negative plate, indicating that the

particles were negatively charged. He also demonstrated that the particles could be de�ected by a magnetic �eld.

Next, he put a scale on the disk and measured the amount of de�ection, trying di�erent materials to generate his cathode ray. �e de�ection was

always the same. From these experiments, �ompson showed that the particles were negatively charged, indistinguishable from each other, and

smaller than atoms. �ompson also calculated the charge-to-mass ratio of these particles. �ese particles were later named electrons, negatively

charged particles within an atom. As with many groundbreaking ideas, �ompson's ideas weren't immediately accepted.

Another breakthrough in atomic theory was made by an American scientist, Robert Millikan, who determined the charge of an electron using the

apparatus shown in the �gure. Very small oil drops were sprayed into the chamber and allowed to fall through a hole in the top plate. �e drops

were electrically charged, both by the x-ray sources and by static electricity. Millikan could slow the fall of the drops by applying an electric �eld

made by the positively and negatively charged plates. Varying the electric �eld strength varied the speed of the drops' fall. By measuring the �eld

strength and the speed of the drops, he calculated that an electron has a charge of 1.6 × 10 C. C stands for coulomb, which is the SI unit of

electrical charge. From �ompson's and Millikan's experiments, the charge-to-mass ratio of an electron and the charge of an electron were

known. �erefore, the mass of an electron could be calculated and was found to be 9.1 × 10 kg.

-19

-31

Apparatus used in Millikan’s “oil drop” experiment

At this point, scientists knew that atoms contained electrons, that atoms had a neutral charge, and that electrons were negatively charged. �ey

inferred that positively charged particles must also exist in atoms. An early model, proposed by �ompson in 1897, was known as the "plum

pudding model." It envisioned the atom as a positively charged mass with negatively charged electrons embedded in it, much like plums in a

pudding. �is idea persisted until experiments by Ernest Rutherford established that atoms had a nucleus.

Rutherford, a physicist from New Zealand, spent much of his scienti�c career in Canada and England. He performed a series of experiments using

a beam of high-speed, positively charged alpha particles produced by the radioactive decay of radium. Alpha particles consist of two protons and

two neutrons, and you'll learn more about radioactive decay later. Rutherford and his colleagues Hans Geiger (later famous for the Geiger counter)

and Ernest Marsden aimed a beam of alpha particles at a very thin piece of gold foil. �e radiation source was embedded in lead to absorb most of

the radiation. �e alpha particles passed through the foil and hit a luminescent screen, after which the scattering of the particles was examined.

Most particles passed through the foil without being de�ected. However, some particles were diverted slightly, and a very small number were

de�ected almost straight back toward the radiation source.

"It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you �red a

15-inch shell at a piece of tissue paper and it came back and hit you."

Ernest Rutherford describing the experiment's results

Rutherford’s gold foil experiment apparatus

Rutherford deduced that because most of the fast-moving alpha particles passed unde�ected through the gold, they must have traveled through

essentially empty space inside the gold atoms. Alpha particles are positively charged, so de�ections arose when they encountered another positive

charge because like charges repel each other. �e few positively charged alpha particles that abruptly changed paths must have hit, or closely

approached, another body that also had a highly concentrated positive charge. Since very few of the particles were de�ected, this positive charge

only occupied a small amount of space in the gold foil.

Analyzing a series of such experiments in detail, Rutherford drew two conclusions:

1. �e volume occupied by an atom must consist of a large amount of space.

2. A small, relatively heavy, positively charged body, the nucleus, must be at the center of each atom.

In 1911, Rutherford proposed a model in which an atom consists of a very small, positively charged nucleus where most of the atom's mass is

concentrated. �e nucleus is surrounded by negatively charged electrons so that the atom is electrically neutral. After many more experiments,

Rutherford discovered that the nuclei of other elements contain the hydrogen nucleus as a "building block." He named this the proton, the

positively charged subatomic particle found in the nucleus. �is nuclear model of the atom, proposed over a century ago, is largely still used today.

In 1913, Niels Bohr, a Danish physicist, proposed that instead of a mass of electrons around the nucleus, electrons orbited the nucleus in �xed,

stable orbits, much like the planets orbit the sun. �is model was contrary to classical Newtonian physics. Bohr's model proposed that electrons

gain or lose energy by moving between these orbits and that these jumps release or require discrete amounts of energy. Bohr made many other

contributions to science and is an important �gure in the history of quantum mechanics.

Next, the Austrian physicist Edwin Schrodinger proposed treating electrons as "matter waves." He came up with an equation to describe the

probability of �nding an electron in a particular area. �ese probability areas de�ned the "orbitals" you'll learn about in the coming lessons.

�e nucleus was known to contain almost all the mass of an atom, with the number of protons only providing half, or less, of that mass. �e

existence of neutral particles in the nucleus was one of several proposals to explain the remaining mass. In 1932, the English physicist James

Chadwick found evidence of neutrons, which are uncharged subatomic particles in the nucleus with a mass approximately the same as that of

protons.

During the early 1900s, scientists isolated several substances that appeared to be new elements from radioactive ores. Analyses showed that some

of these substances were chemically identical to known elements but had di�erent atomic masses. �is result led the English chemist Frederick

Soddy to realize that some elements existed as several types, with di�erent masses that were chemically indistinguishable. �ese di�erent types of

the same elements are called isotopes. Isotopes are forms of the same element that di�er only in the number of neutrons in the nucleus—and

therefore have a di�erent mass. Isotopes are chemically identical because they have the same number of protons.

Timeline of major atomic models

�anks to the ideas and experiments of many scientists, it's widely accepted that the atom is mostly space with a small, dense nucleus of protons

and neutrons while electrons exist in orbitals around the nucleus. In this course, you'll learn more about the shapes of electron orbitals, what

happens when electrons move from one orbital to another, and how subatomic particles can explain macroscopic properties and changes.

�e electric charge of protons and electrons means that they in�uence how electricity moves through a material. Protons and electrons have equal

and opposite charges, so they're attracted to each other. A proton has a positive charge, written +1, and an electron has a negative charge, written

-1. Neutrons have no electrical charge. When an atom has an equal number of electrons and protons, the positive and negative charges balance

each other, and the atom is neutral.

In old textbooks or �lms, atoms are often depicted as acting like a solar system. Electrons are shown orbiting the nucleus like the planets orbit the

sun in the solar system. Although this planetary model for the atom was described by Niels Bohr, the concept isn't accurate. Electrons don't orbit

the nucleus in the same way that planets orbit the sun. Instead, there's a high probability of �nding electrons in regions called orbitals around the

nucleus. �e image below shows two orbitals.

Example of orbitals in an atom

Orbitals can be grouped into electron shells, and the �rst electron shell can hold only two electrons. �e next electron shell holds up to eight

electrons. Subsequent shells can hold more. In the periodic table, the row in which an element is listed tells you how many electron shells its atom

contains. For example, sodium (Na) appears in the third row, so it has three electron shells. Calcium (Ca) is in the second column on the fourth

row. �is means that calcium has four electron shells. �e �rst shell in calcium only contains two electrons, the second and third each contain

eight electrons, and the fourth outermost electron shell contains two electrons.

�e outermost electrons in atoms play an important role in bonding between atoms. Elements that have a full outer shell are inert. �is means

that they don't react with other elements to form compounds. Examples include helium, neon, and argon. For elements that don't have a full outer

shell, the outermost electrons, called valence electrons, can interact with the outermost electrons of nearby atoms to create chemical bonds.

Chemical bonds are formed when atoms share electrons between adjacent outer shells, which doesn't occur if the shells are already full. Just as

orbitals can be considered as clouds of probability for where an electron is located, shared electrons exist in both outer shells with some

probability.

�e diameter of an atom is on the order of 10 m, whereas the diameter of the nucleus is roughly 10 m, about 100,000 times smaller. To get a

picture of the relative sizes of an atom and its nucleus, consider this: If the nucleus were the size of a blueberry, the atom would be about the size

of a football stadium.

Footnotes

−10 −15

[1]

Attribution: Flowers, Paul, Klaus �eopold, Richard Langley, and William R Robinson, Ph.D. Chemistry 2e. Houston, TX:

OpenStax2019. https://openstax.org/books/chemistry-2e/pages/1-introduction

Atomic Theory

Early History of Atomic Theory

�e earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers, the scientists of their day. In the �fth

century BC, Leucippus and Democritus argued that all matter was composed of small, �nite particles that they called atomos, a term derived from

the Greek word for "indivisible." �ey thought of atoms as moving particles that di�ered in shape and size and that could join together.

Later, Aristotle and others theorized that matter consisted of various combinations of the four "elements"—�re, earth, air, and water—and could

be in�nitely divided. �ese philosophers thought about atoms and "elements" as philosophical concepts. �e idea of performing experiments to

test ideas wouldn't be adopted until the seventeenth century, when the scienti�c method was developed.

English chemist John Dalton hypothesized that the behavior of matter could be explained using an atomic theory. First published in 1807,

Dalton's atomic theory was the �rst time that a scientist tried to describe matter in terms of atoms. Many of Dalton's hypotheses about the

composition of matter are still valid in modern atomic theory, and the hypotheses have been modi�ed as well.

Postulates of Dalton's Atomic Theory

1. Matter is composed of exceedingly small particles called atoms. An atom is the smallest unit of an element that can participate in a

chemical change.

2. An element consists of only one type of atom, which has a mass that's characteristic of the element and is the same for all atoms of

that element. A macroscopic sample of an element contains an incredibly large number of atoms, all of which have identical chemical

properties. Atoms of one element di�er in properties from atoms of all other elements.

3. A compound consists of atoms of two or more elements combined in a whole-number ratio. In a given compound, the numbers of

atoms of each of its elements are always present in the same ratio.

4. Atoms are neither created nor destroyed during a chemical change but are instead rearranged to yield substances that are di�erent

from those present before the change.

Dalton's atomic theory provides a submicroscopic explanation of the macroscopic properties of matter that you've learned about. For example, if

an element such as copper consists of only one kind of atom, then it can't be broken down into substances composed of fewer types of atoms. If

atoms are neither created nor destroyed during a chemical change, as stated by the law of conservation of matter, then the total mass of matter is

constant when matter changes from one type to another.

The Law of Constant Composition

Dalton was in�uenced by the law of conservation of matter, but he also knew of the experiments of the French chemist Joseph Proust. Proust

demonstrated that all samples of a pure compound contain the same elements in the same proportion by mass. �is statement is known today as

the law of de�nite proportions, or the law of constant composition. According to Dalton's postulate, in a compound, the ratio of atoms of

di�erent elements is constant. �is is consistent with Proust's observations.

�e following �gure illustrates the law of constant composition with two samples of water. Sample A is a glass of water from the tap, and Sample B

is a bottle of water. �e chemical formula of water is H O, so each molecule of water is made of two atoms of hydrogen (H) and one atom of oxygen

(O). �e data tables on these samples a�rm the law of constant composition. Samples of a pure compound always contain the same elements in

the same mass proportion.

[1]

2

Illustration of the Law of Constant Composition

�e mass of the water in the glass is 10.000 g. �e water in the glass is 11.19% hydrogen and 88.81% oxygen by mass. �e water in the bottle has a

mass of 27.000 g but has the same composition as the water in the glass (11.19% hydrogen and 88.81% oxygen).

It's worth noting that although all samples of a particular compound have the same mass ratio, the opposite isn't always true. �at is, samples that

have the same mass ratio aren't necessarily the same substance.

The Law of Multiple Proportions

One of the postulates of Dalton's atomic theory is now known as the law of multiple proportions. �e law of multiple proportions is an extension

of the law of constant composition. �is law states that when two elements react to form more than one compound, a �xed mass of one element

will react with masses of the other element in a small, whole-number ratio.

To illustrate the law of multiple proportions, this table shows some compounds that can be made from oxygen and nitrogen. Imagine combining

14 g of nitrogen with enough oxygen to form nitrous oxide (N O). You would need 8 g of oxygen. If you were to make nitric oxide (NO) instead—

still using 14 g of nitrogen—you would need 16 g of oxygen. To make dinitrogen trioxide (N O ) with 14 g of nitrogen, you would need 24 g of

oxygen.

Compound Mass of N Mass of O Relative Mass of O

Nitrous oxide, N O 14 g 8 g

Nitric oxide, NO 14 g 16 g

Dinitrogen trioxide, N O 14 g 24 g

[1]

2

2 3

mass O

mass of N

2 ≈ 0.57148

14= 1

8

14

8

14

≈ 1.142916

14= 2

16

14

8

14

2 3 ≈ 1.714324

14= 3

24

14

8

14

�ese ratios themselves may not seem interesting or informative. However, the ratio of these ratios, shown in the last column, is a small whole

number. �is fact is critical in understanding stoichiometry, which is the atomic ratio of elements in a compound.

Both the law of constant composition and the law of multiple proportions began as theories to explain atoms based on observations of how matter

behaved. Dalton's explanations of atoms weren't immediately accepted. No one had seen an atom, and many didn't think such small particles

were possible. Some scientists thought that samples of pure substances could have di�erent compositions. Both theories began to be considered

"laws" after they gained wide acceptance over time. �is is an illustration of the scienti�c method, a process that has characterized science since

the seventeenth century.

The Scienti�c Method

�e ancient Greeks are generally credited with emphasizing reason as the basis for knowledge, and early Greek philosophers didn't believe that

measurement was important. Later, philosophers came to recognize that quantifying information was essential. Over hundreds of years, European

and Persian philosophers recognized that observation, reading others' work and experimentation, documenting and publishing methods and

results, and having others review one's work were important aspects of advancing knowledge. During the Renaissance, the development of the

modern scienti�c method accelerated. �is required contributions from many famous historical individuals: Francis Bacon, Galileo, Isaac

Newton, Johannes Kepler, and others. �e goal of science is to improve and expand knowledge of the universe, and the scienti�c method provides

a guideline for how to accomplish this. Analyze the following �gure to understand the steps of the scienti�c method. You can also view this �gure

as a pdf here.

Steps in the Scienti�c Method

�e �rst part of this process is making observations and using reasoning to propose possible answers to questions. Dalton knew from Proust's

earlier work that elements combined in �xed proportions to form pure compounds—the law of constant composition. He was also familiar with

the law of conservation of mass, developed by the French chemist Antoine Lavoisier. Dalton performed his own experiments and observed that

sometimes when elements combined, more than one type of compound formed. At the time, he was working on formulating his atomic theory.

Dalton carefully measured reactants and products; extensive experimentation allowed him to propose that if elements combine to form more than

[2]

one compound, they do so in �xed ratios of small whole numbers. Over time, further experiments by many scientists con�rmed this theory, and it

eventually became known as the law of multiple proportions. Even this "law" has limitations, though, because larger molecules and polymers

have atoms in much more complicated ratios.

�is is how the scienti�c method works:

1. Identify and de�ne a problem or question, such as the question of how the same elements can form di�erent compounds.

2. Read what others have done, and make initial observations.

3. Formulate a hypothesis, which is an educated guess based on observation and information.

4. Design and conduct experiments to test your hypothesis.

5. Document details of how you conduct your experiments and your results.

6. Analyze your data and share what you found. If your data supports your hypothesis and stands up to the scrutiny of others, you've

contributed to the body of knowledge about this topic. If not, revise your hypothesis and try again. Your results may contradict an

existing, accepted theory.

�eories tested over long periods, subject to thorough scrutiny, become laws. �at being said, there are very few laws in science, and even those

can change or be revised over time as new evidence emerges. �e scienti�c process is ongoing, constantly modifying and even discarding theories

as new information is obtained.

Footnotes

Attribution: Flowers, Paul, Klaus �eopold, Richard Langley, and William R Robinson, Ph.D. Chemistry 2e. Houston, TX:

OpenStax2019. https://openstax.org/books/chemistry-2e/pages/2-introduction

1. From https://openstax.org/books/chemistry-2e/pages/2-1-early-ideas-in-atomic-theory

Physical Properties and Changes

Familiar examples of physical properties include density, color, hardness, melting and boiling points, and electrical conductivity. Some physical

properties, such as density and color, may be observed without changing the physical state of the matter. Other physical properties, such as the

freezing temperature of water, can only be observed as matter undergoes a physical change.

�e physical property volatility refers to how readily a substance vaporizes or changes to the gas phase at a given temperature and pressure. �e

more easily a substance evaporates, the higher its volatility. Substances with low volatility tend to stay in the liquid or solid phase. Volatility can

apply to a liquid going to the gas phase or to a solid that becomes a gas in a process called sublimation. Substances that are gases at room

temperature have very high volatility, and substances with high volatility have low boiling points. You can compare volatilities of substances by

comparing their boiling points at a chosen pressure.

As with boiling points, di�erent substances also have di�erent freezing points. Recall that molecules or lone atoms in liquids and solids are held

together with attractive intermolecular forces that are weaker than chemical bonds. One factor a�ecting the boiling and freezing points of a

substance is the relative strength of these attractive forces, or how tightly molecules hold onto each other.

Phase Changes

It was mentioned previously that some physical properties require a physical change to observe. One of the most common physical changes is a

phase change. To illustrate, you know that freezing is the process of a liquid becoming a solid, and melting is the process of a solid becoming a

liquid. You also know that vaporization is the process of a liquid becoming a gas and that condensation is the process of a gas becoming a liquid.

�ese terms apply to all substances, not just water. Oxygen, for example, can be condensed to a liquid and frozen to a solid given the right

temperature and pressure. Rocks can melt to form liquids, as with lava in a volcano.

Yet two terms you may not have heard before are deposition and sublimation. Sublimation is the process of a solid becoming a gas without

passing through a liquid phase. You may have seen dry ice (solid CO ) sublimate.

Watch this video, Sublimation and Deposition, to see what sublimation of CO looks like. Mothballs (naphthalene) and solid air fresheners are two

examples of common household products that sublimate.

Deposition is the process of a gas becoming a solid without an intermediate liquid phase. �e frost that forms on a window on a cold morning is

an example of deposition. Water vapor in the air undergoes deposition on the window without �rst becoming a liquid.

Deposition is used frequently in the making of consumer products—especially in the electronics industry—to deposit very thin layers of material.

Computer chips have thin and precisely placed layers of various metals. Several deposition techniques are used to form these layers. Many

products, from potato chip bags to dental implants to jewelry, have thin �lms applied using this process.

Phase changes of solids, liquids, and gases.

[1]

2

2

Besides phase changes, common physical changes you'll encounter include mixing, size reduction (grinding), and separation. Distillation is an

example of a physical-separation method that uses the di�erences in substances' boiling points to separate them.

Calculations Involving Density

Density is a commonly used physical property that you'll need to calculate and work with. Recall that density is a measure of a substance's mass-

to-volume ratio. In chemistry, density is often reported in the units grams per cubic centimeter (g/cm ). You may see the Greek letter rho, ρ, used

for density. Some sources use a lower case d.

Densities of some common substances are provided in the table shown here. Note that while mass remains constant, the volume of a gas can

change with temperature and pressure, so density is dependent on these variables. �e volume, and therefore density, of solids and liquids change

very little with temperature and pressure.

Substance Density, g/cm

Water (liquid) 1.000

Ethanol 0.791

Helium 0.00018

Oxygen 0.0014

Cement 2.7–3.0

Mercury 13.6

To calculate the density of a substance, weigh the substance and determine its volume, then divide mass by volume. Weighing a substance by

placing it on a scale to determine its mass is straightforward. For a liquid, the volume can be measured in a piece of glassware such as a graduated

cylinder.

Volume can be determined for some uniform solid objects by measuring dimensions. For example, a cube 2 cm in length on each side has a

volume of:

(2 cm)(2 cm)(2 cm) = 8 cm

However, you'll rarely work with uniform objects like this. One common method for determining the volume of solids is to measure the liquid that

the object displaces by using a graduated cylinder. To do this, you �rst place some liquid in a graduated cylinder and read the volume. �en, place

the solid object into the liquid and read the new volume of the liquid. �e di�erence between these two volumes is the volume of the object.

3

d =mass

volume

3

3

Finding volume using liquid displacement in a graduated cylinder.

You should be able to calculate an object's density, mass, or volume given the other two values. You may need to use conversion factors if the

values aren't given in the required units. Don't forget to use the correct number of signi�cant �gures and to include units.

Example 1:

Find the density of a substance that has a volume of 50.0 cm and a mass of 52.0 g.

Use the density formula with the given values.

Example 2:

A solution has a density of 1.2 g/cm . What is the mass of 4.0 L of this solution?

To �nd mass, use the formula for density:

Substitute the known values for volume (4.0 L) and density (1.2 g/cm ), and then use the conversion factor 1 L = 1000 cm :

4.0 L of a solution with a density of 1.2 g/cm has a mass of 4800 g.

Footnotes

Modi�ed from https://openstax.org/books/chemistry-2e/pages/1-2-phases-and-classi�cation-of-matter

3

d = = = 1.04massvolume

52.0 g

50.0 cm3

g

cm3

3

d = → mass = (volume)(d)massvolume

3 3

mass = (4.0 L )( )(1.2 ) = 4800 g1000 cm 3

1 L

g

cm 3

3

Energy and Mass in Physical and Chemical Changes

So far, the physical and chemical changes that matter can undergo have been discussed in terms of mass. Another important component of these

changes is energy. Mass was de�ned earlier as a measure of the amount of matter a substance contains. Energy is the ability to do work. Some

forms of energy are shown in the �gure. �e many forms of energy can be broadly categorized as kinetic or potential. Types of kinetic energy have

to do with moving objects and the work they're capable of due to their motion. Types of potential energy have to do with position in a �eld, such

as a magnetic or gravitational �eld. �e drawn bow and arrow in the �gure show a type of potential energy. Once the arrow is in motion, the

energy is kinetic. Electrical current is kinetic energy since it's made up of �owing (moving) electrons. Magnetic energy is a type of potential energy.

Molecules have both kinetic and potential energy. While physics is the branch of science more focused on energy, you must grasp some basic

concepts about energy to understand chemical and physical changes of matter.

Before discussing this further, you need to know two basic principles, the conservation of mass and the conservation of energy.

Some forms of energy.

Conservation of Mass

�e law of conservation of mass states that there's no detectable change in the total quantity of matter present when matter converts from one

type to another (a chemical change) or changes among solid, liquid, or gaseous states (a physical change).

Although this conservation law is true for all conversions of matter, convincing examples are few; outside of controlled laboratory conditions,

scientists rarely collect all of the material produced during a conversion. For example, when you eat, digest, and assimilate food, all of the matter

in the original food is preserved. However, some of the matter is incorporated into your body, and much is excreted as various types of waste,

making it challenging to verify by measurement.[1]

Burning wood in a camp�re can provide an example of the conservation of matter. During burning, components of the wood are combined with

oxygen in the air to form carbon dioxide, water, and other gases. Any moisture in the wood is vaporized when it gets hot enough. For the

components of the wood not burned at high enough temperatures to convert to gas, these remain as ash after the �re goes out.

�e total mass of the substances involved doesn't change in this process. All the atoms that were present at the start of the process still exist at the

end, but many are in di�erent molecules. If you were able to collect and weigh all of the air and wood that combine to burn, and collect and weigh

all of the gases and ash products, there would be no change between the combined weight of the starting materials and the combined weight of

the ending materials. �e atoms in the starting molecules would be accounted for in the atoms of the ending molecules.

Wood combines with oxygen in the air to form water vapor, carbon dioxide, and ash.

Conservation of Energy

�e law of conservation of mass says that mass is conserved during physical or chemical changes. Similarly, the law of conservation of energy

says that energy is conserved during such changes. Energy may change form from one type to another, but it doesn't disappear.

Atoms and molecules have both kinetic and potential energy. Kinetic energy is due to the motion of atoms and molecules. Atoms and molecules

are always in motion even if this isn't apparent. Temperature is proportional to the average kinetic energy of the atoms or molecules in a

substance. A higher temperature indicates that a substance has more kinetic energy. Frozen water (ice) has relatively little motion of atoms or

molecules and low kinetic energy. Molecules in a solid don't change position, but they do vibrate. Increasing the temperature until the ice melts

produces a liquid, which has higher kinetic energy—or more movement of molecules. Liquid water molecules move around relative to one

another, although they remain very close to each other. Increasing the temperature enough to vaporize the water results in a gas, whose molecules

have even more motion and more kinetic energy than liquid water molecules. Gas molecules move around a great deal and don't touch unless

they happen to bump into each other.

Energy and temperature increase as a substance moves from solid to liquid to gas.

Atoms and molecules also have potential energy stored in the bonds between atoms in a molecule and the attractive forces between molecules or

between atoms in a substance. Since bonds hold atoms together in molecules, it takes energy to break bonds, and energy is released when bonds

form. When atoms or molecules react, some bonds are broken and others are formed as new substances are made. Sometimes, the new bonds

have more energy than the old ones, and sometimes the old has more energy than the new.

Energy can also change from kinetic to potential or from potential to kinetic. Dynamite, for example, has a lot of potential energy in the form of

chemical bonds. When it explodes, a great deal of this potential energy is changed to kinetic energy.

The Sum of the Kinetic and Potential Energy

�e energy associated with a substance is the sum of the kinetic and potential energy of all the atoms and molecules in the substance. Energy and

mass changes in chemistry are related and often are considered together. Substances that undergo chemical or physical changes also undergo

energy changes.

As an example of how energy changes during a chemical reaction, consider the burning of wood in a camp�re illustrated previously. Components

of the wood combine with oxygen in the air to form carbon dioxide and water vapor. You know that mass is conserved in this reaction, and the

energy of the system is also conserved. �e sum of the kinetic and potential energy of the starting substances is equal to the ending substances,

although forms of energy change. �e heat, light, and sound given o� by the �re are all results of energy change during the combustion reaction.

�e potential energy or bond energy of the wood and oxygen is greater than the potential energy or bond energy of the combustion products. �is

energy is conserved, changes to other forms of energy, and is felt as heat (the product gases are hot), seen as light, and heard as sound. �e

products of the reaction have more kinetic energy and a higher temperature than the starting substances, wood and oxygen, which had more

potential energy in the form of chemical bonds.

Wood combines with oxygen in the air to form water vapor, carbon dioxide, and ash.

Exothermic vs. Endothermic

When a chemical or physical change releases energy to its surroundings, it's said to be exothermic. �e combustion of wood is an example of an

exothermic chemical reaction. Energy is released as heat to the surroundings since the products of the reaction are at a higher temperature. It's

also released in the form of light and sound given o� by a �re. Condensation of water vapor to a liquid is an exothermic physical change because

the vapor loses energy to become a liquid. Exo- means "outside," or "external."

When a chemical or physical change requires the addition of energy to occur, it's said to be endothermic. Evaporation is an example of an

endothermic physical change. Heat must be added to liquid water for it to evaporate. Cold packs used for sports injuries are an example of an

endothermic chemical reaction. When the chemicals in the sports pack mix, they react and take heat from their surroundings. Endo- means

"internal," or "inside." Sometimes, you may �nd it hard to decide if a change is endothermic or exothermic. �ink about how the surrounding

temperature changes as a result of the change.

�e symbol ΔH, discussed in relation to the combustion of wood, indicates a transfer of heat to the surroundings in exothermic processes (if it's

listed with the products) and a transfer of heat from the surroundings in endothermic processes.

Endothermic vs Exothermic

In science, the primary unit of measurement for energy is the joule, abbreviated as J and measured in kgm /s . A unit of energy you may be more

familiar with is the calorie. �e calorie content of food is typically written on food labels. �e word "calorie" that commonly describes the energy

content of food is actually a kilocalorie to scientists. One calorie is equal to about 4.19 joules, but the calories reported on food labels are equal to

about 4187 joules. Calories are used in some areas of science instead of or in addition to joules as a measurement of energy. You may see calories

in chemistry calculations, but joules will be used most often.

Footnotes

Modi�ed from https://openstax.org/books/chemistry-2e/pages/1-2-phases-and-classi�cation-of-matter

2 2

Properties and Changes of Matter

Recall that chemical changes are those in which the identities of substances change. For example, one carbon atom (C) reacts with two oxygen

atoms (O) to form one molecule of a new substance, carbon dioxide (CO ). �e identity and chemical properties of carbon dioxide are di�erent

from those of carbon and oxygen.

Also, recall that physical changes are those in which the chemical identities of substances don't change. Carbon dioxide can be frozen to form

what's commonly called "dry ice." Both dry ice and gaseous carbon dioxide are CO .

Common examples of physical changes include the following:

Wax melting

Sugar dissolving in co�ee

Steam condensing into liquid water

Physically separating materials, such as by distillation or �ltering

Magnetizing and demagnetizing metals (such as antitheft security tags)

Grinding solids into powders

In each of these examples, the physical state, form, or properties of the substance are changed, but its chemical composition isn't.

Chemical Properties and Changes

You're familiar with chemical properties such as �ammability, toxicity, solubility, and acidity. Another important chemical property is reactivity, a

substance's tendency to react with other substances. Some substances react readily with other substances, while others are relatively unreactive.

�e element potassium, for example, is quite reactive. Neon and helium, however, have very low reactivity.

Chemical properties require a chemical change to be observed and quanti�ed. Acidity is a chemical property that you'll have the chance to

observe and quantify. Flammability, toxicity, and reactivity are other examples of chemical properties. Measuring �ammability requires igniting

(combusting) a substance. As you become more familiar with the periodic table and learn more about the structure of molecules, ways to describe

these properties will be discussed further.

�e term chemical change indicates that a chemical reaction takes place. Such reactions occur when atoms combine with other atoms or

molecules, or when molecules combine with each other, to form substances with di�erent chemical identities. �e starting substances are called

reactants, and the resulting substances are called products. Atoms in molecules are held together by strong attractive forces called bonds. �e

breaking and forming of bonds between atoms in substances results in changes in the amount of energy in the substances. To start understanding

the nature of chemical reactions, you must know the language and symbols for describing chemical substances and reactions.

Chemical Formulas and Symbols

Recall that all elements have a symbol consisting of one or two letters. See the following table for some common examples.

Element Symbol

Hydrogen H

Carbon C

Copper Cu

Chromium Cr

2

2

Sodium Na

Iron Fe

Nitrogen N

Chlorine Cl

Nickel Ni

Sulfur S

Oxygen O

You know that most elements don't exist as single atoms. Atoms join together chemically, forming bonds to make molecules. When molecules are

formed, the element symbols are written together, with subscripts to indicate the number of atoms of each element. When there's only one atom of

an element, no subscript is used. �is is called a molecular formula, or a chemical formula. For example, CH , methane, is a molecule that

contains one carbon atom and four hydrogen atoms. HCN, hydrogen cyanide, is a molecule formed from one hydrogen atom, one carbon atom,

and one nitrogen atom. CO, carbon monoxide, is a molecule formed from one oxygen atom and one carbon atom.

In addition to using chemical formulas in equations, you'll see chemical formulas written in sentences in place of a molecule's name. Although

these examples are fairly simple, some molecules, especially in biology, can be quite large and complex.

When two atoms combine to form a molecule, this is a diatomic molecule. If the two atoms are both of the same element, the molecule is a

homonuclear diatomic molecule. Oxygen is most often found as two oxygen atoms bonded together, written as O . Other molecules of oxygen

are possible but they aren't diatomic. For example, three oxygen atoms bonded together form an ozone molecule, O .

When discussing chemical reactions, you may see a number in front of a molecule. �is coe�cient indicates the number of molecules. For

example, 3CO means there are three carbon dioxide molecules. 4N O refers to four molecules of N O, nitrous oxide. Don't confuse a coe�cient

with a subscript.

�e �gure here represents the reaction of hydrogen and oxygen to form water. Hydrogen and oxygen both typically exist as diatomic molecules, H

and O . H , O , and H O are the molecular formulas for hydrogen, oxygen, and water molecules. �e coe�cients in front of hydrogen and water

(2H and 2H O) show that two hydrogen molecules are used in the reaction and two water molecules are produced. You know that this is a

chemical change since water is a di�erent substance from the starting substances, hydrogen and oxygen.

Reaction of oxygen with hydrogen to form water

4

2

3

2 2 2

2

2 2 2 2

2 2

Structural Formulas

Another type of formula used to describe molecules is a structural formula. Structural formulas are similar to molecular formulas, but they show

how atoms are bonded together. �e structural formula for chloroform, CHCl , is shown in the �gure. �e molecular formula tells you that the

molecule has one carbon atom, one hydrogen atom, and three chlorine atoms. �is doesn't give you information about how the atoms are

connected. �e structural formula, however, shows that carbon is bonded to each of the other four atoms and that the chlorine atoms and the

hydrogen atom are only bonded to the carbon atom.

Structural formula for chloroform

As another example, the structural formula for ammonia, NH , is shown. �is structural formula shows that the nitrogen atom is bonded to each of

the hydrogen atoms.

Structural formula for ammonia.

Empirical Formulas

A molecular formula tells how many of each atom are in a molecule. An empirical formula is another type of formula that only gives the ratio of

atoms in a molecule. For example, hydrogen peroxide has the molecular formula H O , and its empirical formula is HO. �ere are two H atoms

and two O atoms, so their ratio in the molecule is 1:1. To determine an empirical formula, you �nd the ratio of the atoms and then divide by the

lowest common denominator. For example, acetic acid has a molecular formula of C H O . Carbon, hydrogen, and oxygen are in a ratio of 1:2:1, so

the empirical formula is CH O. �e molecular formula will always be a whole number multiple of the empirical formula.

Sometimes, the empirical formula will be the same as the molecular formula. �e molecular formula for carbon monoxide is CO. Carbon and

oxygen are in a 1:1 ratio, and the molecular formula is already in this form.

Introduction to Nomenclature

You know many chemicals by their common names, such as water, table salt, ammonia, bleach, baking soda, and aspirin. �ese names don't give

any information about the composition of these substances, and they don't convey the nature of these substances to someone who is unfamiliar

with the common names. �ere's a system of naming substances in chemistry that gives this type of information, referred to as chemical

nomenclature.

�e chemical name of a substance allows you to distinguish it from other substances. �e name needs to convey the composition of substance—

that is, how many of each type of atom are in the substance. Multiple substances can have the same chemical formula but their atoms are arranged

di�erently, giving these substances di�erent properties and behaviors. A chemical name must distinguish between such substances as well. A

chemical name indicates both the composition of the substance and the arrangement of the atoms in the substance.

3

3

2 2

2 4 2

2

Isomers are compounds with the same chemical formula but di�erent molecular structures, and they're an example of the importance of a clear

system of naming chemicals. For example, acetic acid and methyl formate both have the molecular formula C H O . Methyl formate is used in

manufacturing as an insecticide and for quick-drying �nishes. It has an oxygen atom between two carbon atoms, di�ering from the atomic

arrangement of acetic acid molecules. Acetic acid and methyl formate are structural isomers, compounds in which the same atoms are

connected di�erently. �is small di�erence in the atoms' arrangement has a major e�ect on the molecules' chemical properties. You certainly

wouldn't want to use a solution of methyl formate as a substitute for a solution of acetic acid (vinegar) when making salad dressing.

Molecules of (a) acetic acid and (b) methyl formate are structural isomers; they have the same formula (C H O ) but di�erent structures and therefore di�erent

chemical properties.

Many types of isomers exist. Besides structural isomers, there are various types of spatial isomers, in which the relative orientations of the atoms

in space can be di�erent. For example, the compound carvone is found in caraway seeds, spearmint, and mandarin orange peels. It consists of two

isomers that are mirror images of each other. S-(+)-carvone smells like caraway, and R-(−)-carvone smells like spearmint.

S-(+)-carvone and R-(−)-carvone are spatial isomers because they only di�er in the relative orientations of their atoms in space.

Arrangements of Atoms and Molecules

So far, the formulas that you've learned don't indicate how molecules are arranged relative to one another in solids, liquids, and gases. Structural

formulas show the arrangement of atoms in a single molecule, indicating which atoms are bonded together. �ough less common, some

substances are composed of atoms rather than molecules. For example, at typical pressures and temperatures, gold (Au) is a solid, and neon (Ne)

is a gas. Both of these substances are composed of individual atoms, not molecules. �e element mercury (Hg) can be found as a liquid.

Recall that molecules or lone atoms in a gas move around freely. In a liquid, the molecules or atoms are close together but can still move in

relation to one another. In a liquid, the attraction between individual molecules keeps them close to one another, similar to bonds that hold atoms

together in molecules, but weaker. �e nature of these intermolecular attractions can help explain the properties of substances. In solids, where

individual molecules don't move relative to one another, this attractive force is stronger than the attractive forces found between liquid molecules

but still not as strong as the bonds between atoms within a molecule.

In solids, molecules or lone atoms are arranged in regular repeating patterns, or as repeating subunits composed of groups of atoms or molecules

in a pattern. How atoms or molecules are arranged in a solid helps you understand the properties of a substance. A few examples of how lone

atoms or molecules are arranged in solids are shown in these �gures. �e �rst �gure shows how water molecules in ice are arranged as sheets of

connected hexagons.

2 4 2

2 4 2

[1]

Structure of ice. Oxygen atoms are shown in red and the much smaller hydrogen atoms are in pink. Molecules arrange in hexagonal sheets.

Gold atoms arranged in a cube.

�is �gure shows how solid-gold atoms arrange to form a cube.

Footnotes

Modi�ed from https://opentextbc.ca/chemistryatom�rst2eopenstax/chapter/chemical-formulas/

Mixtures

Most materials aren't pure substances but are instead mixtures. Mixtures are composed of two or more substances—compounds, molecules, or

elements—and can be separated by a physical change. An exception to this de�nition is an alloy, discussed later in this section. In a mixture, the

identities of the substances remain the same. Mixtures can be either heterogeneous or homogeneous.

Homogeneous mixtures have the same properties throughout. �e pre�x homo- means "same," and geneous means "kind," or "type." Examples of

homogeneous mixtures are saltwater, air, and brewed co�ee. In a cup of saltwater, any sample of saltwater from the cup would have the same

composition and properties of any other sample taken from that cup of saltwater.

Heterogeneous mixtures have nonuniform properties. Hetero- means "di�erent." Some examples of heterogeneous mixtures are oil and vinegar

salad dressing and fruit salad. Properties of a salad dressing sample vary depending on where you take a sample. Randomly taking a sample, you

may get more oil or you may get more vinegar.

�e �gure here summarizes the ways matter can be classi�ed when broken down as pure substances or mixtures. Click here to view this chart as a

pdf.

Categories of Matter

Solutions, Colloids, and Suspensions

A homogeneous mixture with only one phase is referred to as a solution. Solutions are single-phase homogeneous mixtures of two or more

substances, where one substance is completely dissolved in the other. �e substance present in a smaller amount is called the solute, and the

substance present in the larger amount is called the solvent. For example, carbonated water is carbon dioxide dissolved in water. Since water is

present in the larger amount, it's the solvent, and carbon dioxide is the solute. Solubility is a measure of how much of a solute can be dissolved in

a given solvent.

A common type of solution is an aqueous solution. An aqueous solution has water as the solvent. Aqua is from the Latin word for "water." To

indicate that a compound is dissolved in water, (aq) is written after a compound. For example, sodium chloride dissolved in water is an aqueous

solution of NaCl, and this would be written "NaCl (aq)."

�e word solution suggests a liquid, but solutions can be liquid, solid, or gas. For example, air is a solution composed mainly of oxygen and

nitrogen. As with all mixtures, air can be separated into components by physical methods. Alloys are usually homogeneous mixtures that are solid

solutions but di�cult to separate via physical methods. Chemical changes sometimes can be used to more easily separate alloys. You'll learn more

about alloys shortly.

In chemistry, two other types of homogeneous mixtures are important to understand. A suspension is a homogeneous mixture with particles that

settle out over time. Examples include muddy water and dust particles in the air. A colloid is a homogeneous mixture with solid or liquid particles

that don't settle out on their own. Examples of colloids include jelly and fog.

Alloys

Alloys are mixtures containing at least one pure metal mixed with one or more other metal(s) or nonmetal(s). Some common alloys are brass (a

mixture of copper and zinc), steel (a mixture of iron and carbon), and bronze (a mixture of copper and tin). Elements of an alloy are melted

together at high temperature and then cooled to a solid. During this process, atoms of the elements making up the alloy become mixed in a way

that's di�cult to separate by physical means.

For example, the �gure here shows a brass alloy. �e copper atoms arrange themselves in a cube shape, common for solid-metal atoms. A zinc

atom occupies a space inside the copper cube in this brass alloy. Alloys are made to enhance some property of a singular metal element. Although

alloys are typically homogeneous, they may be heterogeneous, depending on the process used to make them.

Lattice structure of copper-zinc brass.

Tarnish on a silver cup is a thin layer of silver sul�de formed when silver reacts with oxygen and hydrogen sul�de in the air.

Sterling silver, another alloy, is a solid solution of silver and another metal, usually copper. Silver tarnishes easily, so other metals are added to

lessen this tendency and increase the strength of silver. Tarnish refers to a thin layer of corrosion that forms on some metal products. Silver reacts

with oxygen and hydrogen sul�de in the air to form silver sul�de.

As mentioned, it's di�cult but not impossible to separate alloys. Some alloys may be melted and allowed to settle over time into their constituent

elements. Alternatively, the alloy may be vaporized and separated by distillation, which is a separation by di�erences in boiling points of

constituent elements. Melting, settling, and distillation are all physical changes. Often, an easier method is to chemically extract the original

elements of the alloy using an acid. Alloys provide an example of when a mixture can be separated by a chemical change.

Di�erences between Alloys and Compounds

Alloy Compound

Not easily separable by physical changes but sometimes

separated by a chemical changeOnly separated by chemical changes

Components can be mixed in a range of ratiosFound in �xed ratios of atoms, listed by the

chemical formula

May not have bonds Held together by chemical bonds

Has some properties of its base metal, the metal with the

highest proportion in the alloy

May have vastly di�erent properties than its

elemental components

Has at least one metal element as a component May or may not contain a metal element

Classifying Matter

Now that you've learned about measuring matter, you'll focus on how matter is classi�ed.

Pure Substances

Materials can be classi�ed as either pure substances or mixtures. Pure substances have a uniform chemical composition and uniform properties

throughout. Examples include table salt (sodium chloride), copper, and water. When a sample is taken from a pure substance at any point, it will

have the same chemical composition and properties, such as melting point, density, and color, as any other sample taken from that substance.

Pure substances have chemical and physical properties that can identify them. Mixtures are de�ned and discussed later in this section.

Elements, Atoms, Molecules, and Compounds

Elements are pure substances that can't be broken down into smaller components by chemical changes. Oxygen and copper are examples of

elements. If you were to take some quantity of an element and divide it repeatedly, you would eventually have a particle of that same element that

couldn't be divided any further, called an atom. Recall that an atom is the smallest unit of matter that has the characteristics of a chemical

element. Recall also that a single atom has the same chemical properties as a million atoms together. It's possible to break an atom into smaller

pieces called subatomic particles, but these smaller pieces wouldn't have the properties of the atom that was broken apart.

Most substances aren't made up of single atoms but of molecules. Atoms can join together to form molecules. Atoms in molecules are held

together by bonds, the "glue" that holds molecules together. You'll learn about several di�erent types of bonds and how they a�ect molecular

behavior. Some types of bonds hold atoms together in molecules. Another type of attractive force between molecules, or between atoms that aren't

in a molecule, holds particles together to varying degrees in solids and liquids.

Atoms of the same element can join together to form molecules, or atoms of di�erent elements can combine to form molecules. Oxygen, for

example, exists in air as an oxygen molecule—two oxygen atoms bonded together. Ozone is a molecule formed by three oxygen atoms bonded

together.

Oxygen and ozone molecules

Molecules can be composed of just two atoms (as in O ) or many atoms. Some molecules contain thousands of atoms! Molecules may also contain

atoms from many di�erent elements. If a molecule is made of more than one type of element, it's called a compound. Carbon dioxide and carbon

monoxide are examples of compounds since they're formed by two di�erent types of atoms, carbon and oxygen. Table salt (NaCl) is another

example of a compound. �ere are countless compounds but only 118 elements. Of the 118 elements, 98 occur naturally on Earth. �e oxygen and

ozone molecules previously shown aren't compounds since they're composed of atoms that are all the same element.

2

Carbon monoxide and carbon dioxide are compounds.

Common Elements

Every chemical element has been assigned a symbol consisting of one or two letters of the alphabet. If only one letter is used, it's capitalized. If

two letters are used, the �rst is capitalized and the second is lowercased. Hydrogen, for example, has the symbol H, and lithium has the symbol Li.

All of the known elements are organized in the Periodic Table.

Hydrogen (H) is the most common element in the universe, followed by helium (He). In comparison, the other elements are found in very small

quantities. On Earth, however, oxygen (O) is the most abundant element. �e following �gure shows the composition of the Earth's crust. Again,

oxygen is the most abundant element, mostly found in combination with other elements as compounds in rock. Silicon (Si) is the second most

common, followed by aluminum (Al), iron (Fe), calcium (Ca), and sodium (Na).

Chemical Composition of the Earth’s Crust

A chemical formula is used to show the number and types of atoms that make up a molecule. Constituent elements are listed, and the number of

atoms of each are given as subscripts. When there's only one atom of that element, no subscript is used. For example, the chemical formula for

water is H O. �is formula tells you that one molecule of water is made up of two atoms of hydrogen and one atom of oxygen. A few examples of

elements and molecules are shown in the table that follows, along with their chemical symbols or formulas.

Elements Molecules and their Chemical Formulas

Gold (Au) Oxygen (O )—composed of two oxygen atoms

Mercury

(Hg)Methane (CH )—composed of one carbon atom and four hydrogen atoms

Neon (Ne) Ammonia (NH )—composed of one nitrogen atom and three hydrogen atoms

2

2

4

3

Radon (Rn)Sulfuric acid (H SO )—composed of two hydrogen atoms, one sulfur atom, and four oxygen

atoms.

Natural versus Synthetic Compounds

Some compounds occur naturally, while others are synthetic. Synthetic compounds are made in a lab or industrial manufacturing facility,

sometimes in imitation of natural compounds. Often, it's less expensive to manufacture a synthetic substance than to obtain it from natural

sources. For example, rubber is a naturally occurring compound, but synthetic rubber has mostly replaced natural rubber. Manufactured rubber is

both less expensive to make and has better properties than natural rubber when used to make tires. All plastics are synthetic compounds.

Many of the products you use are synthetic compounds or contain synthetic compounds as additives. Unless the clothes you're wearing are

completely made of a natural material such as cotton or wool, you're wearing synthetic materials or natural-synthetic blends. Even if you're

wearing purely natural fabrics, they're likely colored with synthetic dyes. Other examples of synthetic compounds include ammonium nitrate (a

fertilizer), aspirin, many vitamins, and most pharmaceutical drugs.

2 4

Converting Units of Measurement

Units of measure will often need to be converted from one type to another, both within the SI system and between SI and other measurement

systems. In the United States, a mix of SI units and English units is common. Both systems use seconds for time, so no conversions are needed.

Converting units within the SI system is fairly easy. Converting between the American system and SI takes a bit more work, but you likely have

done similar conversions in math classes. A few examples are shown here as a refresher, starting with a few conversions within the SI system. To

convert from one unit to another, multiply by the appropriate conversion factor.

Example 1: Convert 6540 km to m.

Example 2: Convert 42 g/cm to kg/m .

Example 3: Convert 5 g/mL to kg/L.

Example 4: Convert 1 L to cm .

Pay special attention to the conversion for one liter:

1 L = 1000 cm = 1 × 10 cm

�is is a conversion factor that you'll use frequently in chemistry, especially when working with solutions.

To convert from the English to the SI system, you need to know some conversion factors. Several common conversion factors are given in the

following table.

American to SI SI to American

inches to centimeters 1 in = 2.54 cm centimeters to inches 1 cm = 0.3937 in

ounces to grams 1 oz = 28.35 g grams to ounces 1 g = 0.03527 oz

gallons to liters 1 gal = 3.785 L liters to gallons 1 L = 0.2642 gal

°F to °C °C to °F

Here are a couple of examples to help you practice converting between American and SI units.

Example 5: Convert 25 °C to degrees Fahrenheit.

(6540 km )( ) = 6, 540, 000 m = 6.54 × 106 m1000 m

1 km

3 3

(42 )( )( )3

= 42, 000

= 4.2 × 104

g

cm3

1 kg

1000 g

100 cm

1 m

kg

m3

kg

m3

(5 )( )( ) = 5g

ml

1 kg

1000 g

1000 ml

1 L

kg

L

3

(1 L )( )( )3

= 1000 cm3 = 1 × 103 cm31 dm3

1 L

10 cm

1 dm

3 3 3

(F − 32) = C5

9C + 32 = F

9

5

C + 32 = F95

(25 ∘C) + 32 = 77 ∘F95

Example 6: Convert 105 °F to degrees Celsius.

Example 7: Convert 25 pounds to grams.

Example 8: Convert 15 in to cm .

Example 9: Convert 57 yards to meters.

Using only the conversion factors listed in the tables provided, along with knowing there are 12 inches in a foot and 3 feet in a yard, you would

work this conversion in the following way:

Accuracy and Precision

When counting everyday objects, such as the number of students in a class, it's possible to get an exact number. Although you can count discrete

objects since they have de�ned boundaries, you'll have other situations in which there's a degree of uncertainty in the measurements. It's

important to understand and be able to communicate this degree of uncertainty.

Scientists typically take several measurements of the same quantity to ensure that the measurement is reasonable. You may be familiar with the

terms accuracy and precision from past science classes. Accuracy refers to how close a value is to its true or accepted value. Precision is how close

multiple measurements are to each other. �e following table of targets illustrates the di�erence between accuracy and precision. In the upper-left

target, shots aren't close to the center of the target or each other, so they're neither precise nor accurate. In the upper-right target, shots are all

close to the center of the target but not to each other, representing accuracy without precision. In the lower-left target, shots are grouped close

together but not near the center of the target, showing precision without accuracy. Finally, in the lower-right target, shots are both close together

and near the center; they're both accurate and precise.

Not Accurate or Precise Accurate but not Precise

Not Accurate but Precise Both Accurate and Precise

(F − 32) = C59

(105 ∘F − 32) = 41 ∘C59

(25 lb )( )( ) = 11, 340 g16 oz

1 lb

28.35 g

1 oz

3 3

(15 in3)( )

3

= 246 cm32.54 cm

1 in

(57 yd )( )( )( )( ) = 52 m3 ft

1 yd

12 in

1 ft

2.54 cm

1 in.1 m

100 cm

Signi�cant Figures

You can use a variety of tools, such as scales, rulers, and timers, to take measurements. However, you need to observe some rules for reporting

measurements. �is ensures that a measurement doesn't seem more accurate than the instrument being used is capable of. You can only report

certain numbers called signi�cant �gures. �ese show to what degree of certainty a measurement is known.

Apples on a scale have a weight of 456 g.

For example, the scale in the �gure shows that the weight of the apples is 456 g. We would write this as 456 g, not 456.0 g or 456.00 g, because this

scale doesn't report weight to a tenth or a hundredth of a gram. In the measurement 456 g, the 4 in the hundreds place and the 5 in the tens place

are certain. �e last number, 6 in the ones place, is an estimate. �is is true for measurements in general. Unless otherwise indicated, the last

number on an instrument is an estimate, accurate to the smallest scale division. In this example, the weight is known to ± 1 g. �e actual weight

could be 455 g or 457 g. You would say this measurement has three signi�cant �gures: two certain numbers (the 4 and the 5) plus the last,

estimated number (the 6).

Digital thermometer showing a temperature of 101.4 °F.

Here's another example. �e thermometer in the �gure shows a temperature of 101.4 °F. �is measurement has four signi�cant �gures—three that

are certain (1, 0, and 1) and one that's an estimate (4). �e true value is between 101.3 and 101.5 °F.

When reporting measurements, all nonzero digits are signi�cant. For example, 45.97 has four signi�cant �gures. It can be confusing to determine

whether a zero counts as a signi�cant �gure. A few rules for zeros are listed in the following table.

Rule for Zeros Example

Zeros between nonzero digits are signi�cant.305.22 has �ve signi�cant �gures.120.003 has six

signi�cant �gures.

Lone zeros after a decimal point are signi�cant.45.0 has three signi�cant �gures.1267.0 has �ve signi�cant

�gures.

Zeros that only set the decimal point are not

signi�cant.

50,000 has one signi�cant �gure. 0.0034 has two

signi�cant �gures.

When rounding, round up if the last number is �ve or greater. Round down if the last number is four or less. When performing calculations with

signi�cant �gures, you want to understand how accurate a result is. In general, the result of a calculation can't be more accurate than the least

accurate number or measurement used in the calculation.

When adding or subtracting, the result can't have more decimal places than the least accurate measurement.

Example 10:

145.0 g of water ← 1 place after the decimal

+ 0.627 g of NaCl ← 3 places after the decimal

145.627 g 145.6 g ← The answer can have only 1 place after the decimal.

Example 11:

45.003 m ← 3 places after the decimal

-10.01 m ← 2 places after the decimal

34.993 m 34.99 m ← The answer can have only 2 places after the decimal.

When multiplying or dividing, the result can't have any more signi�cant �gures than the least accurate measurement.

Example 12:

(75.4 m/s)(22 s) = 1658.8 m 1700 m

(3 signi�cant �gures)(2 signi�cant �gures) ← 2 signi�cant �gures

Example 13:

(10 g) / (5.02 g/m ) = 1.992 m 2 m

(1 signi�cant �gure) / (3 signi�cant �gures) ← 1 signi�cant �gure

You may come across a case that isn't straightforward. For example, suppose a sample that needs to be weighed is placed on a scale that's accurate

to the nearest gram. �e scale shows a weight of 300 g. According to the previous rules, writing this as 300 g would indicate only one signi�cant

�gure. You can't write 300.0 g since this would be a more accurate measure than the scale is capable of. Instead, you would write this measurement

as 3.00 × 10 g.

3

3

3 3

3 3 3

2

Measurement in Chemistry

Previously, you learned that chemistry evolved from alchemy. To move away from alchemy, there was an emphasis on quantifying or measuring

observations. To make measurements in chemistry, you'll need to study the SI system of units, convert between and within systems of measure,

di�erentiate between accuracy and precision, and identify how the degree of certainty in measurements is expressed.

A Common Language of Measurement

�roughout history, people have used many ways of measuring quantities. Measuring often used comparisons with something familiar or

common. One "mile" was originally designated by Romans as the distance covered in 1000 double paces. A "foot" was the length of a typical man's

foot. A "cubit" was the length from the end of the middle �nger to the elbow. �e obvious problem with these older measurement systems is the

lack of consistency. For example, not everyone's stride is the same length, so di�erent people pacing o� 1000 steps will come up with di�erent

distances. A consistent, reproducible set of measuring standards was necessary.

�e International System of Units (SI) has emerged as the common measurement system of science accepted internationally. �ere are seven

base units in the SI system, shown in this table. You'll mainly use the �rst �ve of these base units in chemistry.

Measurement Base Unit Unit Abbreviation

length meter m

mass kilogram kg

time second s

temperature Kelvin K

amount of a substance mole mol

electric current ampere A

luminous intensity candela cd

Although you may not be familiar with the mole, this unit of measure is used extensively in chemistry and will be explained later in greater detail.

A mole is a quantity, much like "dozen" is used to mean 12. Yet a mole is a much larger number: 6.022 × 10 .

Another unit that may be new to you is Kelvin for temperature. You've probably used degrees Celsius (°C) as a unit of measure for temperature,

and, from the American system, degrees Fahrenheit (°F). Water freezes at 32 °F and boils at 212 °F. �e Celsius scale is based on water, with 0 °C

being the point at which water freezes and 100 °C, the point at which water boils. On the Celsius scale, the theoretical coldest point attainable is

called absolute zero, and its value is –273.15 °C. At absolute zero, molecular motion stops. �e Kelvin scale has the same increments as the Celsius

scale but de�nes absolute zero as its starting point, 0 K. Note that temperatures are expressed in Kelvin without a degree sign. �e freezing point of

water is 273.15 K, and the boiling point of water is 373.15 K. Since 0 K is the same as –273.15 °C, simply add or subtract 273.15 to convert between

degrees Celsius and Kelvin.

C = K – 273.15

K = C + 273.15

Example 1: Convert 240 K to degrees Celsius.

23

To convert to degrees Celsius, subtract 273.15 from the Kelvin temperature.

240 – 273.15 = –33.15 °C

Example 2: Convert 35 °C to Kelvin.

To convert from degrees Celsius to Kelvin, add 273.15 to the Celsius temperature.

35 °C + 273.15 = 308.15 K

Kelvin, Celsius, and Fahrenheit temperature scales

�e SI system uses powers of 10 to derive other units, with pre�xes to indicate the scale of the units. A kilometer (km), for example, means 1000

meters because "kilo" means 10 . A centimeter (cm) is one hundredth (or 10 ) of a meter. Common pre�xes in the SI system are shown in the

following table.

Pre�x (symbol) Factor

3 –2

tera (T) 10

giga (G) 10

mega (M) 10

kilo (k) 10

deka (da) 10

deci (d) 10

centi (c) 10

milli (m) 10

micro (µ) 10

Nano (n) 10

pico (p) 10

femto (f) 10

Units of measure may also be derived from combinations of base units or other derived units. For example, the standard unit of measure for

volume is a cubic meter (m ). However, the unit most commonly used to measure volume is a liter (L), which is a cubic decimeter (dm ). Density

is typically expressed in grams per cubic centimeter (g/cm ) for solids and grams per liter (g/L) for gases.

12

9

6

3

–1

–2

–3

–6

–9

–12

–15

3 3

3

Phases of Matter

An important physical property of matter is its phase. On Earth, matter has four natural phases, or states:

1. Solid

2. Liquid

3. Gas

4. Plasma

Solids and liquids are matter: It's clear that they take up space and have mass. Gases are also matter. If gases didn't take up space, a balloon

wouldn't in�ate, or increase in volume, when �lled with gas.

A solid is rigid and possesses a de�nite shape. A liquid �ows and takes the shape of its container, except that it forms a �at or slightly curved upper

surface when acted upon by gravity. In zero gravity, liquids assume a spherical shape. Both liquid and solid samples have volumes that are almost

independent of pressure. A gas takes both the shape and volume of its container.

e three most common states of matter on Earth.

You've probably heard the term "plasma" in describing certain TVs. Plasma is a state of matter. Occurring naturally inside stars, it's a gaseous state

of matter that contains signi�cant numbers of electrically charged particles. �e presence of these charged particles gives unique properties to

plasma, which distinguish its matter classi�cation from gases. In addition to stars, plasmas are found in both natural and man-made, high-

temperature environments: lightning strikes, specialized instruments for detecting trace amounts of metals, cutting tools, and certain electronic

screens.

Plasma precision cutting tool.

Some matter samples appear to have properties of solids, liquids, and/or gases at the same time. �is can occur when the sample is composed of

many small pieces. For example, you can pour sand like a liquid because it's composed of many small grains of solid sand. Matter can also have

properties of more than one state when it's a mixture. Clouds are a good example of this: �ey appear to behave like gases but are mixtures of air

(gas) and tiny particles of water (liquid or solid).

A lesser-known state of matter is a Bose-Einstein Condensate (BEC). BEC is a state of matter that exists when a material is cooled to near absolute

zero. Its existence was theorized by Albert Einstein and the Indian physicist Satyendra Bose in the 1920s. BECs were �rst created in a lab in 1995. At

extremely low temperatures, matter becomes very dense, and groups of atoms behave as a single atom. BECs are useful in chemistry and physics

research. Several other states of matter have been identi�ed—some like BECs at extremely low temperatures, and others at extremely high-energy

or high-pressure environments, such as the centers of stars.

To learn about how BECs were created in a lab, watch Creation of a Bose-Einstein Condensate.

Macroscopic, Microscopic, and Submicroscopic

Matter can be studied at the macroscopic, microscopic, or submicroscopic level. Macro is a Greek word that means "large." �e macroscopic

world is familiar to you since it contains objects that can be directly sensed by human sight or touch. �is includes everyday observations and

laboratory chemistry, where you can observe and measure physical and chemical properties such as mass, density, solubility, and temperature.

Micro also comes from Greek and means "small." �e microscopic world of chemistry is visible through standard optical instruments such as

microscopes. Schools typically have microscopes that can magnify up to between 200x and 400x. (Remember that "200x" means the image is 200

times larger than it appears to the naked eye.) For example, the butter�y wing shown here at 30x is 30 times larger than it appears to the naked eye.

Although this isn't an incredibly high magni�cation level, it still reveals unexpected details of the butter�y's wing.

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A butter�y wing magni�ed 30x.

You're familiar with microscopes that use light to look at images and are called light microscopes. Submicroscopic particles such as atoms are too

small to be seen with even the best light microscopes. Chemical elements are substances that can't be broken down any further chemically. An

atom is the smallest unit of matter that has the characteristics of a chemical element. If you took some quantity of a substance and divided it

repeatedly, you would eventually have a particle of that element that can't be divided any further; this particle is an atom. A single atom has the

same chemical properties as millions of atoms put together. While it's possible to break an atom into smaller pieces (called subatomic particles),

these smaller pieces would no longer have the properties of the atom.

Sophisticated instruments in university and professional labs are capable of imaging entities as small as some atoms. Scanning electron

microscopy and transmission electron microscopy are two tools used to image samples at a very small scale. A scanning tunneling microscope

(STM) uses a quantum phenomenon called "tunneling" to produce images of, and even manipulate, individual atoms.

A look into the Ultra High Vacuum (UHV) chamber of a UHV Scanning Tunneling Microscope (STM). Several grippers are mounted to move samples back and forth

between the holder for multiple samples and the STM microscope, which is the tubular gold-capped structure held by a spring suspension.

Watch IBM Atomic Short: History of the scanning tunneling microscope to �nd out more about STMs.

Most submicroscopic particles are too small to be seen even with the most advanced imaging instruments and may only be pictured in the mind.

Techniques like STM allow scientists to observe electron density and chemical bonds. Still, subatomic particles such as electrons, protons, and

neutrons are too small to be directly seen with the currently available instruments.

Applying Symbols to Behaviors and Properties

Symbolic representations help you understand chemistry from the macroscopic to the submicroscopic level. Common examples of chemical

symbols are found in the periodic table, chemical formulas, chemical equations, graphs, drawings, and calculations.

�e following example about water illustrates the terms macroscopic and submicroscopic, showing how symbols can help scientists explain

properties and behaviors.

Macroscopic Observations

Water is a liquid at moderate temperatures, freezes to form a solid at lower temperatures, and boils to form a gas at higher temperatures.

Solid, liquid, and gas are macroscopic descriptions of water.

Submicroscopic Level

As water changes from solid to liquid to gas, the interactions between individual molecules are a�ected. As a solid, water molecules are held

tightly together and are packed in a pattern. In a liquid, water molecules are close together but can move freely. In a gas, water molecules have

space between them.

From left to right: Water molecules in the solid phase, liquid phase, and gas phase.

Symbolic Representation

Water molecules are composed of two hydrogen atoms and one oxygen atom. �e formula H O and the following �gure show two hydrogen atoms

connected to one oxygen atom. �ese are examples of symbols that help explain the structure and behavior of water molecules. 2

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Water can be represented by its formula, H O, and by a physical model, where red and white balls represent atoms, with sticks representing the bonds between atoms.

Here's another illustration of how macroscopic and submicroscopic views are used to understand matter. Diamond and graphite are both made of

carbon.

Cut diamond

2

A chunk of raw graphite and a pencil, showing the “lead” made from mixing graphite with clay.

Although both materials contain only carbon, their macroscopic properties are very di�erent. �e two materials have distinct appearances, and

diamond is very hard, while graphite is not. Diamond doesn't conduct electricity, but graphite does. Macroscopic di�erences between these

materials can be understood by looking at how the carbon atoms are arranged on a submicroscopic level.

Arrangement of Atoms

�e �gure provides a symbolic representation of the submicroscopic arrangement of atoms in diamond and graphite to help visualize how the

atoms are connected to one another. In graphite, carbon atoms are arranged in sheets that can slide over one another, causing graphite to be

relatively soft. In diamond, atoms are arranged in a rigid structure and are unable to move, making diamond very hard. In later sections, you'll

learn how the �ow of electricity is a�ected by how atoms are bonded or connected. Both materials can be represented symbolically by the letter C,

for carbon.

Representations of how carbon atoms are arranged in a crystalline form in diamond (left) and sheets in graphite (right).

Footnotes

1. Attribution: Flowers, Paul, Klaus �eopold, Richard Langley, and William R Robinson, Ph.D. Chemistry 2e. Houston, TX:

OpenStax2019. https://openstax.org/books/chemistry-2e/pages/1-introduction