Atoms: the building blocks of matter

Chemistry chapter 3 1

Atoms: the building blocks of matter

Chapter 3


The atom The atom – smallest piece of matter that has

the properties of an element. Made of

Protons Neutrons Electrons

Each specimen of a specific subatomic particle is the same If we split an atom, we no longer have a specific

element


Early atomic theory - Democritus Greek philosopher about 400 B.C. Gave us the word atom

Atomos - indivisible. Thought

The world was made of empty space and particles called atoms.

There were different types of atoms for different types of materials.

Theory was not supported by experimental evidence.


Early atomic theory – Aristotle Aristotle did not believe in atoms

thought matter was continuous He was very influential, so Democritus’s

theory was not accepted for many centuries.


17th century

People began to express doubts in Aristotle’s theory.

Isaac Newton and Robert Boyle published articles stating their belief in the atomic nature of elements, but they had no proof.

Their theory also had no ability to predict the unknown.


Antoine Lavoisier – late 1700s Law of conservation of mass

during a chemical change in a closed system, no mass is lost


Joseph Proust – late 1700s to early 1800s Law of definite proportions

specific substances always contain elements in the same ratio by mass


Law of multiple proportions

Some elements form more than one compound with each other.

If two or more different compounds are composed of the same two elements, then the ratio of their masses always contains small whole numbers


John Dalton – early 1800s

Studied experimental observations of chemical reactions

Proposed explanation of these three laws


Dalton’s Hypothesis1. All matter is composed of very small particles

called atoms.

2. All atoms of an element are exactly alike; atoms of different elements are very different.

3. Atoms cannot be subdivided, created, or destroyed.

4. Atoms unite with other atoms in simple ratios to form compounds

5. In chemical reactions, atoms are combined, separated, or rearranged.


Did Dalton’s theory work?

Conservation of mass the atoms are simply rearranged because they

cannot be created or destroyed Laws of definite and multiple proportions

Only whole atoms can combine, giving small whole numbers in ratios


Gas research J.L. Gay-Lussac

Under constant temperature and pressure Volumes of reacting gases and gaseous products are in

a ratio of small whole numbers.

Amadeo Avogadro explained Gay-Lussac’s work with Dalton’s theory. Equal volumes of gases, under the same

temperature and pressure, have the same number of molecules.

Helped Dalton’s theory get accepted


Dalton’s theories

Atomic theory and law of multiple proportions have been tested and accepted as correct.

However, there some major exceptions to the rules. Splitting atoms Different atoms of the same element


Discussion

Section review on page 69


Cathode tubes

Anode – positive electrode Cathode – negative electrode When the tube is on, cathode rays appear

that begin at the cathode and travel to the anode.


Cathode rays and electrons

1897 – J.J Thomson tested cathode rays and discovered that they were electrons. Rays turned a paddlewheel – they had mass Rays deflected by a magnet just like current-

carrying wire – they were negatively charged He determined the ratio of the electron’s

charge to its mass.


Charge on an electron

Robert Millikan’s famous oil drop experiment. Tiny oil drops fell through a chamber

gravitational force offset by applying an opposing electrical force.

Charge on oil drops determined This charge was always a whole number

multiple of one small charge


Charge on an electron

This small charge was the charge on one electron.

This is now the standard unit of negative charge (1-). It can be written e-.

e- can also represent an electron


Mass of an electron

Using Thomson’s ratio and Millikan’s charge, determined to be 9.1 x 10-31 kg

It was found that it’s mass is only 1/1837 the mass of the lightest atom known – the hydrogen atom. Most of the mass must be somewhere else Since atoms are neutral, there must be some

positive charge


Thomson’s plum pudding model

In this model, the raisins were the electrons and the pudding was the positive charge.

Sort of like chocolate chip cookie dough. The chips are the electrons and the dough is the

positive charge. Explained the experiments that had been

done so far.


Testing the plum pudding model See page 72 fired alpha particles at a very thin (a few

atoms thick) sheet of gold foil. They expected the particles to go right

through because the spread out positive charge in the “pudding” wouldn’t be strong enough to deflect them.


What happened

Most of the particles did go right through without being deflected at all.

Some were deflected at large angles. Ernest Rutherford explained it:

the positive charge on the atom was concentrated at a small core – now called the nucleus.


The atom as we now “know” it The nucleus contains all of the positive

charge and most of the mass. The negatively charged electrons have very

small mass and are located around the nucleus in the electron cloud.

Most of an atom is empty space.


Protons

same charge as an electron; opposite sign. standard unit of positive charge (1+) Much larger mass than the electron:

1.67 x 10-27 kg The number of protons determines the atom’s

identity.


Neutrons

Weren’t discovered until the 1930s. Neutral – no charge – harder to detect

Slightly more mass than a proton: 1.68 x 10-27 kg


Nuclear or Strong Force

The force that holds protons and neutrons together.

It is effective only for very short distances – about 10-15 m.


Dalton’s theory

Dalton thought that atoms were indivisible discovery of electrons, protons, and neutrons did

not fit with his theory. Led to major revisions in atomic theory


Isotopes

Thomson discovered what seemed to be two kinds of neon atoms. Same chemical properties; different masses.

Atoms of the same element that differ in mass are called isotopes. Have the same number of electrons and protons

but different number of neutrons.


Atomic number

Number of protons in an atom Atoms are electrically neutral, the number of

electrons must equal the number of protons. The number of protons determines the

identity of the atom and the number of neutrons determines the isotope.


Modification of Dalton’s theory All atoms of an element contain the same

number of protons but can contain different numbers of neutrons.

So we have to use average mass of an atom.


Nucleons

Particles in the nucleus – protons and neutrons


Mass number

Total number of nucleons : protons plus neutrons

Number of neutrons = mass number minus atomic number


Designating Isotopes

Hyphen notation Uranium-235 Carbon-14 Carbon-12

The number refers to the mass number


Nuclide

General term for any isotope of any element


Atomic mass units

There must be a standard for all units of measurement.

A Carbon-12 atom with 6 protons and 6 neutrons was chosen as the standard

C126


Atomic mass unit

Defined as 1/12 the mass of that carbon atom.


Average atomic masses

Many elements have an average atomic mass close to the number of nucleons in their nuclei – near whole numbers.

Some don’t – look at Chlorine The periodic table shows average atomic

masses.


Weighted averages

We then use a weighted average to find the average mass of an atom of a given element.

This is called the average atomic mass or just atomic mass.


Finding a weighted average

A class of 25 students took a test. 10 of them got 80%. 12 got 90%. 3 got 100%. What was the average score?

Not 90% - probably less than that.

%2.8725

%1003%901280%10average weighted


You try

Neon has two isotopes. Neon-20 has a mass of 19.992 amu and neon-22 has a mass of 21.991 amu. In any sample of 100 neon atoms, 90 will be neon-20 and 10 will be neon-22. Calculate the average atomic mass of neon.

20.192 amu


You try

Compute the average atomic mass of silver, if 51.83% of the silver atoms occurring in nature have mass 106.905 amu and 48.17% of the atoms have mass 108.905 amu.

107.9 amu


The Mole

SI unit for amount of substance Abbreviated mol A counting unit 6.022 x 1023 particles

Avogadro’s number Based on carbon-12, 12 g of C-12 contains

1 mol of atoms


Molar mass

The mass of 1 mol of a pure substance g/mol Numerically equal to the atomic mass in amu

On the periodic table the number with a decimal is the atomic mass in amu AND the molar mass in g/mol


conversions

Grams to moles or moles to grams Use the molar mass


Example

What is the mass in grams of 5.60 mol of sulfur?


Example

How many moles of carbon are in a sample with a mass of 567 g?


Example

How many atoms of lithium are in a sample with a mass of 76.2 g?


You try

How many moles of rubidium are in 3.01 x 1023 atoms of rubidium?


You try

How many moles are in 0.255 g of zinc?


You try

What is the mass of 1.20 x 1025 atoms of helium?

Atoms: the building blocks of matter

Documents

Transcript of Atoms: the building blocks of matter