Chapter 3: Atoms

The Building Blocks of Matter

An atom is the smallest particle of an element that retains the chemical properties of that

element.

Section 1

The Atom: From Philosophical Idea to Scientific Theory

Page 67

The Early Atom

As early as 400 B.C., Democritus called nature’s basic particle the “atomon” based on the Greek word meaning “indivisible”.

Aristotle succeeded Democritus and did not believe in atoms. Instead, he thought that all matter was continuous. It was his theory that was accepted for the next 2000 years. (Read page 43 of your textbook.)

Three Basic Laws of Matter:

Law of Conservation of MassLaw of Definite ProportionsLaw of Multiple Proportions

Basic Laws of Matter

Law of Conservation of Mass- mass is neither created nor destroyed during ordinary chemical reactions or physical changes.

CH4 + 2O2 → 2H2O + CO2

16g + 64g → 36g + 44g

Antoine Lavoisier

stated this about 1785

Basic Laws of Matter Law of Definite Proportions – no matter how much

salt you have, it is always 39.34% Na and 60.66% Cl by mass.

Example: Sodium chloride always contains 39.34% Na and 60.66% Cl by mass.

2NaCl → 2Na + Cl2100g → 39.34g + 60.66g116.88g → ? + ?

Joseph Louis Proust

stated this in 1794.

Basic Laws of Matter Law of Multiple Proportions- Two or more

elements can combine to form different compounds in whole-number ratios.

Example

John Dalton proposed this

in 1803.

Dalton’s Atomic Theory In 1808, Dalton proposed a theory to

summarize and explain the laws of conservation of mass, definite proportions, & multiple proportions.

I was a school teacher at the

age of 12!

Dalton’s Atomic Theory

1. All matter is composed of extremely small particles called atoms.

2. Atoms of a given element are identical in size, mass, and other properties.**

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

4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds.

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

**Today, we know these parts to have flaws.

John Dalton - 1808

Flaws of Dalton’s Theory…

2. Atoms of a given element are identical in size, mass, and other properties.

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

Isotopes – atoms with the same number of protons but a different number of neutrons

Subatomic particles – electrons, protons, neutrons, and more

Section 2

The Structure of the Atom Page 72

The Atom

Atom - the smallest particle of an element that retains the chemical properties of that element.

CARBON

The Structure of the Atom

The atom is composed of two main regions, the nucleus & the electron cloud.

Nucleus of an Atom Nucleus- very small region located at the center of the

atom. The nucleus accounts for most of an atoms mass but very little volume, making it a very dense region.

The nucleus contains protons, neutrons, and more.

proton = p+ neutron = no others – neutral, too

MD=

V

Electron Cloud of an Atom The electron cloud is the negatively

charged region of the atom that accounts for most of the atom’s volume but very little of the atom’s mass.

electron = e-

The electron cloud is composed of a number of electrons, of which depends the element.

Checking for Understanding

What are the two main regions of the atom?

Does an electron from gold, act like gold?

What is the charge on the nucleus?

NO, an electron is like any other electron, no matter the source.

The nucleus and the electron cloud are the two main regions.

The nucleus is positive since it holds protons (+), neutrons (0) and other neutral particles.

Subatomic Particles

Protons- positively charged particles found in the nucleus of an atom.

Neutrons- neutral particles found in the nucleus of an atom.

Electrons- negatively charged particles found in the electron cloud.

Others – photon, boson, gluon, lepton, muon, quark, tau, neutrino, meson, …

Properties of Subatomic Particles

Particle Symbol Charge Mass # Relative Mass (amu)

Actual Mass (g)

Electron e- -1 0 0.0005486 9.109 X 10-28

Proton p+ +1 1 1.007276 1.673 X 10-24

Neutron no 0 1 1.008665 1.675 X 10-24

1 amu (atomic mass unit) = 1.660540 x 10-27 kg or exactly 1/12 the mass of a carbon-12 atom

Discovery of the Subatomic Particles

The discovery of the subatomic particles came about from the study of electricity & matter.

Benjamin Franklin’s kite experiment in 1752 demonstrated that lightning was electrical.

Charged Particles

In 1832, Michael Faraday proposed that objects are made of positive and negative charges.

Discovery of the Electron In the late 1870’s many experiments were

performed in which electric current was passed through gases at low pressures due to the fact that gases at atmospheric pressure don’t conduct electricity well.

These experiments were carried out in glass tubes called cathode-ray tubes or Crookes tubes.

Sir William Crookes developed these tubes.

Crookes TubeCRT

Discovery of the Electron

When current was passed through the cathode ray tube, the surface of the tube, directly opposite the cathode, glowed.

It was thought that this glow was caused by a stream of particles called cathode rays.

The rays traveled from cathode (negative) to anode (positive).

Discovery of the Electron

Negatively charged objects deflected the rays away.

Therefore, it was determined that the particles making up the cathode rays were negatively charged.

Cathode Ray Tube Experiment Accomplishments

Proved that the atom was divisible and that all atoms contain electrons.

This contradicted Dalton’s Atomic Theory. This allowed a new model of the atom.

Plum-Pudding Model of the Atom

Checking for UnderstandingCathode Ray Tube Why were the cathode rays deflected?

Why did they assume there was a positive portion to the atom?

How did this contradict Dalton’s model of the atom?

They were negatively charged, so they were repelled from the negative plate and attracted to the positive plate.

They knew the atom was neutral, so by default, there must be a positive portion if there are negative particles.

Dalton stated that atoms cannot be subdivided. Electrons are subatomic particles.

Oil Drop Experiment Millikan dropped negatively charged microscopic

oil particles into a chamber containing metallic plates and viewed them with a microscope.

By applying voltage to the metallic plates, Millikan created an electric field.

He was able to suspend the oil droplets by adjusting the electric field to the appropriate strength and direction to overcome gravity.

Oil Drop Experiment

Knowing the mass of the droplets and the strength of the electric field necessary to suspend them, he was able to calculate the charge of the electron.

He noticed that the charge was always a whole-number multiple of 1.602 X10-19 Coulombs.

He determined that the charge of the electron to be 1.602 X 10-19 C.

Checking for UnderstandingOil Drop Experiment What year did Millikan perform this

experiment?

How did he view the oil droplets?

He did NOT measure the charge on the electron; he calculated it. What did he measure?

1909

He viewed them with a microscope.

He knew the mass of the droplets and the strength of the electric field.

Discovery of X-Rays

In 1895 William Conrad Roentgen discovered X-rays, a form of radiation.

Radioactivity In 1896, the French scientist

Henri Becquerel was studying a Uranium mineral. He discovered it was spontaneously emitting high-energy radiation.

In 1898, Marie and Pierre Curie attempted to isolate radioactive components of the mineral.

Radioactivity

In 1899, Ernest Rutherford, a British scientist, began to classify radiation: alpha (), beta (), and gamma ().

Radiation Look closely at the paths of radiation. Do

you notice something about the amount of deflection of each type of particles?

Radiation

Discovery of the Nucleus

In 1911, Ernest Rutherford performed a Gold Foil Experiment.

He and his colleagues bombarded a thin piece of gold foil with fast moving, positively charged alpha particles.

Alpha Particles

Alpha () particles are Helium-4 nuclei. This means they are two protons and two

neutrons (with no electrons). Thus, they are positive.

4 +22 He

Discovery of the Neutron In 1932, James Chadwick discovered the

neutron. Rutherford predicted that there were

massive, neutrally charged particles in the nucleus, but it was Chadwick who proved their existence.

Bohr’s Model of The Atom

Forces in the Nucleus

The nucleus is positive (p+ and no) Like charges repel each other…so

shouldn’t the p+ in the nucleus repel each other?

But…when 2 p+ are close together in the nucleus there is a strong attraction between them. The same holds true for neutrons.

REMEMBER:

Forces in the Nucleus

no act like the “glue” that holds the nucleus together. They help to stabilize the nucleus.

Nuclear forces are the short-range p+-no, p+- p+ , and no-no forces that hold the nuclear particles together.

Atomic Number

atomic number (Z) - the number of protons in the nucleus of each atom of a given element.

The number of p+ identifies the element.

Atomic Number increases from left to right on the periodic table.

Electrons

The number of electrons in a neutral atom is equal to the number of protons in that atom.

e- = p+

•Electrons can be lost or gained.

• When electrons are lost or gained, ions are formed.

Ionsion- an atom with a positive or negative

charge.cation- an atom with a positive charge Cations are formed when an atom loses

negatively charged electrons. Ca+2 is formed when calcium loses 2

electrons. Ca+2 has 2 less electrons than protons.

the lithium atom

the lithium ion

Ions

anion- an atom with a negative charge Anions are formed when an atom gains

negatively charged electrons. N-3 is formed when nitrogen gains 3

electrons. N-3 has 3 more electrons than protons.

Noble gases are very stable

and don’t react.

Every element on the periodic

table will try to react to be

stable, like the noble gases.

Metals vs. Nonmetals

Metals form cations.

Na Na+ + 1e-

Nonmetals form anions.

Cl + 1 e- Cl-

Charge determination

Group 1 – forms +1 Group 2 – forms +2 Group 13(B and Al) – forms +3 Group 15 – forms -3 Group 16 – forms -2 Group 17 – forms -1 Group 18 – doesn’t forms ions easily!

(WITH SOME EXCEPTIONS!)

Mass Number

mass number (A)- the number of p+ & no in the nucleus of an atom.

# of neutrons = mass number – atomic number

Why aren’t electrons included when determining the mass number of the atom?

Isotopes

isotope- two or more atoms having the same atomic number (same #p+) , but different mass numbers (due to different #no).

nuclide- general term for a specific isotope of an element.

Isotopes

Isotope Notation

Nuclear Notation

Hyphen NotationUses the elements symbol followed by a hyphen & the mass

number.

C-12

How many protons, neutrons & electrons are there in the following?

Cl-38 35Cl-1

Br-80 32S-2

N-14 56Fe+3

Changes in the Nucleus

Nuclear Reaction- changes that occur in the atom’s nucleus.

Nuclear reactions can change the composition of an atom’s nucleus permanently.

Types of Radiation Produced in Nuclear Reactions

Alpha ()

Beta ()

Gamma ()

Nuclear Stability

Atoms with unstable nuclei are radioactive. Most atoms have stable nuclei and are,

therefore, not radioactive.

Nuclear Stability

Neutrons help to stabilize the nucleus.

Elements 1-20 have p+ = no

Above element 20, increasingly more no are needed than p+ to maintain nuclear stability.

Element 84 and up, all atoms are radioactive so the nucleus cannot be stabilized regardless of the number of no.

Types of Radioactive Decay

Alpha Radiation ()- stream of high energy alpha particles. Consists of 2 protons & 2 neutrons making it identical to

a He-4 nucleus. Alpha particles can be represented by:

Most alpha particles are able to travel only a few centimeters through air and are easily stopped by clothing etc.

4 4 +2 42 2 2 He He

Alpha Decay

239 235 494 92 2Pu U + He

234 230 492 90 2U Th + He

parent daughter

Types of Radioactive Decay

Beta Radiation () – consists of a stream of high speed electrons. These electrons are not electrons that are in motion around the atom’s nucleus.

Beta particles can be represented by:

Can penetrate through clothing and damage skin.

0 -1 0 01 1 1 e e

Beta Decay

6 6 02 3 -1He Li + 24 24 011 12 -1Na Mg +

parent daughter

Types of Radioactive Decay

Gamma Rays ()- energetic form of light that cannot be seen.

Does not contain particles. Gamma particles can be represented by:

Can penetrate heavy material including skin. Can only be stopped by lead or concrete.

00

Other Types of Nuclear Reactions positron –

proton -

neutron -

01e

10 n

11H

Half Lifehalf life- the time required for half of the atoms in any given quantity of a radioactive isotope to decayEach particular isotope has its own half-life.

Half Life

p.689 Sample Problem B

Phosphorus-32 has a half-life of 14.3 days. How many milligrams of phosphorus-32 remain after 57.2 days if you start with 4.0 mg of the isotope?

Ans: 0.25 mg

The Mole

mole (mol)- SI Unit for the amount of a substance that contains as many particles as there are atoms in exactly 12g of carbon-12.

A unit of counting, like the dozen.

Avogadro’s Number

Avogadro’s Number - the number of particles in exactly one mole of a pure substance.

1 mole = 6.0221415 X 1023

1 mol = 6.02 x 1023

Amedeo Avogadro

Atomic Mass

atomic mass - the mass of one mole of an atom

Atomic mass is expressed in atomic mass units (amu) or (u) or g/mol.

Can be found on the periodic table. All atomic masses are based on the

atomic mass of carbon-12 being 12 amu.

Molar Mass

molar mass - the mass of one mole of a pure substance.

Molar mass is written in units of amu or g/mol.

Atomic mass vs. Molar mass

atomic mass - the mass of one mole of an atom.

molar mass - the mass of one mole of a pure substance.

Atomic Mass vs. Molar Mass

Example Atomic Mass

Na

Ag

C

O

22.99 g/mol

107.87 g/mol

12.01 g/mol

16.00 g/mol

Molar Mass of Compounds

Compound Molar Mass

H2O

C6H12O6

NaCl

Cl2

(NH4)3PO4

CuSO4·5H2O

18.02 g/mol

180.18 g/mol

58.44 g/mol

70.90 g/mol

149.12 g/mol

249.72 g/mol

Introduction to Molar Conversions

Amount Mass

1 mol O2

½ mol O2

2 mol O2

3 mol O2

32.00 g

16.00 g

64.00 g

96.00 g

Grams to MolesConverting grams to moles: divide by molar mass.

1. How many moles of Ca are in 5.00g of Ca?

2. How many moles of H2O are in 36.0g of H2O?

3. How many moles of AgNO3 are in 124.5g of AgNO3?

1 mol Ca5.00g Ca x =

40.08 g Ca0.125 mol Ca

22

2

1 mol H O36.0 g H O x =

18.02 g H O22.00 mol H O

33

3

1 mol AgNO124.5 g AgNO x =

169.88 g AgNO0.7329

3mol AgNO

Moles to Grams

Converting from moles to grams: multiply by molar mass

1. What is the mass in grams of 2.25 moles of Fe?

2. What is the mass in grams of 0.896 moles of BaCl2?

55.85 g Fe2.25 mol Fe x =

1 mole Fe126 g Fe

2187 g BaCl22

2

208.23 g BaCl0.896 mol BaCl x =

1 mole BaCl

Types of Particles

Atoms – C, Cu, He

Molecules – O2, C12H22O11, CO2 (all nonmetals in the formula)

Formula units – NaCl, CaCl2, Mg(NO3)2 (includes a metal in the formula)

1 mole = 6.02 x 1023 particles

Particles to MolesConverting particles to moles: divide by

Avogadro's Number.

1. How many moles of Pb are in 1.50 X 1025 atoms of Pb?

2. How many moles of CO2 are in 6.78 X 1021 molecules of CO2?

2.49 x 101 moles Pb

1.13 x 10-2 moles CO2

Moles to ParticlesConverting moles to atoms: multiply by

Avogadro's Number.

1. How many molecules of NO are in 0.87 moles of NO?

2. How many formula units of NaI are in 2.50 moles of NaI?

5.2 x 1023 molecules NO

1.51 x 1024 formula units NaI

Grams to Moles to Particles

Example: How many molecules of N2 are in 57.1g of N2?

257.1 g N x 2

2

1 mol N

28.02 g N

232

2

6.02 x 10 molecules Nx

1 mol N

242= 1.23 x 10 molecules N

Particles to Moles to Grams

Example: How many grams of NaF are in 7.89 X 1024 formula units of NaF?

247.89 x 10 f.un. NaF x 23

1 mol NaF

6.02 x 10 f.un. NaF

41.99 g NaFx

1 mol NaF

= 550. g NaF

Atoms to Moles to Grams

Tough Example:

How many total atoms are in 235 g of CO2?

9.64 x 1024 total atoms

The Mole Bridge

Atomic Mass Determination

Average Atomic Mass - the weighted average of atomic masses of the naturally occurring isotopes of an element.

The atomic mass is expressed relative to the value of exactly 12u for a carbon-12 atom.

Atomic Mass Unit – amu or u

What is a weighted average?

Example: Your grade in math might be 75% tests and 25% homework. What would your grade be if you had a test average of 80% and a homework average of 100%?

Normally, you would average 80% & 100% to get 90%.

However, with a weighted average, it is 85%.

80%(.75) 100%(.25) 85%

Calculating Average Atomic Mass

Average atomic mass =

atomic mass of each isotope X percent natural abundance (in decimal form)

average atomic mass =

(atomic mass of each isotope x % abundance of each isotope in decimal form)

Percent Natural Abundance

Percent Natural Abundance- the relative proportions expressed as percentages, in which isotopes of an element are found in nature.