Chapter 3

Atoms: the smallest particle of an element

that retains the properties of that element.

(Greek: atomos = indivisible)

Democritus (Greek teacher in the 4th

century BC)

First suggested the idea that atoms

existed

Section 3.1: The Atom: From

Philosophical Idea to Scientific Theory

1700’s – chemists were able to relate

changes to individual atoms

Average atom size:

Mass = 1 x 10 –23 g

Diameter = 1 x 10-8 cm

How small is that?100,000,000 copper

atoms in a row would = 1 cm in length!

The Atom

Law of Conservation of Mass

Definition: mass cannot be created or destroyed

only transformed

Law of Definite Proportions

Definition: a chemical compound contains the

same elements in exactly the same proportions

by mass regardless of the size of the sample or

source

Example:

▪ Sodium chloride: NaCl always consists of exactly

39.34% sodium & 60.66%chlorine by mass

▪ Water: H2O always consists of exactly 11.18%

hydrogen & 88.82% oxygen by mass

Law of Multiple Proportions

Definition: if two or more different compounds

are composed of the same two elements, then

the ratio of the masses of the second element

combined with a certain mass of the first

element is always a ratio of small whole

numbers

Examples:

CO & CO2 : 1:1 ratio & a 1:2 ratio

H2O & H2O2: 2:1 ratio & a 2:2 ratio

John Dalton

English school teacher

Proposed an explanation for the 3 laws

Established in 1808

1. All elements are composed of tiny

indivisible particles called atoms.

2. Atoms of the same element are

identical. The atoms of any one

element are different from those of

any other element.

Dalton’s Atomic Theory

3. Atoms of different elements can combine

with one another in simple whole number

ratios to form compounds.

H2O C12H22O11 NOT H2.5O¾

4. Chemical reactions occur when atoms are

separated, joined or rearranged. Atoms of

one element are not changed into atoms of

another by chemical means!

5. Atoms can not be subdivided.

Most of Dalton’s Atomic Theory is accepted

One major revision includes that idea that

atoms are indivisible….

There are 3 parts to an atom….

1. electrons

2. protons

3. neutrons

Section 3.2 – The Structure of the Atom

Negatively charged subatomic particles J.J. Thomson discovered in 1897

Passed an electric current through gases at low pressures called a “Cathode Ray Tube”

Noticed the surface of the tube directly opposite the cathode glowed.

Why? Opposites attract and the electrons were attracted to the positive ends and lights up!

Cathode Ray Tube

Discovery of the Electron

Cathode Ray Tube

Cathode rays are identical regardless of the element

Therefore all elements must have electrons!

Other important findings:

Atoms are electrically neutral, so they must contain a positive charge to cancel it out

Since electrons are so small, atoms must contain other particles that account for their mass

J.J. Thomson – plum-pudding model

e- are spread evenly throughout the positive

charge of the rest of the atom

Ms. Agostine’s “mint chocolate chip ice

cream model”

Robert Millikan (1868-1953)

Found quantity of charge in 1

electron (e-)

Also determined the ratio of the

charge to the mass of 1 e-

Calculated the mass of 1 e-

Electrons weigh 9.109 x10-31 kg

Ernest Rutherford (1911)

nucleus of the atom is positively

charged

Gold Foil Experiment

Most particles go straight through Positively charged particles deflect off of the

positively charged nucleus(~1/8,000)

Gold Foil Experiment

“…it was as if you fired a 15-inch [artillery]

shell at a piece of tissue paper and it came

back and hit you.”

Nucleus was very small

If a nucleus were a marble

the atom would be a football field

Protons (p+)

Positively charged

particles

Mass = 1.673 x 10-27

kg

1,836 times heavier

then an electron

Neutrons (no)

Subatomic particles

with no charge

Discovered by Sir

James Chadwick

Mass is nearly the

same as a proton

Mass = 1.675x10-27 kg

Particle

Symbol

Relative

Charge

Relative

Mass

(amu)

Actual

Mass (kg)

Electron e- 1- 1/1836 9.11x10-31

Proton p+ 1+ 1 1.67x10-27

Neutron no 0 1 1.68x10-27

Particle

Symbol

Relative

Charge

Relative

Mass

(amu)

Electron e- 1- 0

Proton p+ 1+ 1

Neutron no 0 1

Atomic Number : the number of protons in the

nucleus of an atom of an element

Atoms are electrically neutral

Tells how many electrons there are also!

Periodic Table

#1 – Hydrogen: has 1 p+ and 1 e-

#6 – Carbon: has 6 p+ and 6 e-

Mass Number – total number of protons and neutrons in a nucleus

# of neutrons = mass # - atomic # = (# p+ + # no) - (# p+) Ex) Beryllium – 9 Hyphen notation: The number “9” is the

mass number # of p+? # of no? # of e-?

Definition – atoms that have the same number

of protons but different numbers of neutrons

Different types of the same element

Ex) Carbon – has 3 isotopes

Carbon – 12

Carbon – 13

Carbon – 14

Differ by # of no

All have the same # of p+

If not, it would be a different element

All have 6 protons

Carbon – 12

Has 6 neutrons

Carbon – 13

Has 7 neutrons

Carbon – 14

Has 8 neutrons

Hydrogen-1: 1 p+ and 0 no

Relative abundance = 99.985 %

Commonly called normal “hydrogen”

Hydrogen-2: 1 p+ and 1 no

Relative abundance = 0.013%

Commonly called heavy hydrogen or “deuterium”

Hydrogen-3: 1 p+ and 2 no

Relative abundance = 0.002%

Commonly called “tritium”

Definition – weighted average mass of the

atoms in a naturally occurring sample of the

element

Carbon-12 = 98.89 % abundant

Carbon-13 = 1.11% abundant

Carbon-14 = ~0.0000001% abundant

Formula:

Atomic = relative • mass # + relative • mass # +

mass abund. abund.

Repeats for as many isotopes as exist for that element….

Units: atomic mass unit (amu): defined as

exactly 1/12 the mass of a carbon-12 atom

1 amu = approximately the mass of 1 proton

amu’s are used so you don’t have to use scientific notation

when talking about such small masses

Remember that the mass of 1 proton ≈ mass of 1 neutron

Sample Problem:

Chlorine has 2 isotopes:

chlorine-35 which is 75.77%

abundant and chlorine-37

which is 24.33% abundant.

What is the atomic mass of

chlorine?

35 Cl = 75.77% abundant = 0.7577 rel. abund. 37 Cl = 24.33% abundant = 0.2433 rel. abund.

Atomic mass =

= (35 amu x 0.7577) + (37 amu x 0.2433)

= (26.5195 amu) + (9.0021 amu)

= 35.5216 amu

Compare to value on Periodic Table = 35.453 amu

which rounds to 35.45 amu

We must have a way to relate masses in grams

to numbers of atoms…

The amount of a substance that

contains as many atoms, molecules,

ions, or other elementary units as the

number of atoms in 12.01 g C.

That number is equal to 6.02 × 1023

It is called Avogadro's number.

Named in honor of

Amedeo Avogadro

di Quaregna

(1776-1856)

Proposed:

Equal volumes of different gases at

the same temperature and pressure,

contain the same number of

molecules.

How do you buy donuts?

How do you buy computer paper?

How do you buy pencils?

How do you buy soda?

You use a counting/measuring

amount!

1 dozen donuts = 12 donuts

1 ream of paper = 500 sheets

1 gross of pencils = 144 pieces.

1 case of soda = 24 cans

602,000,000,000,000,000,000,000

1 mole of paperclips: goes around earth 4

trillion times

1 mole of large marshmallows: covers the

continental USA 650 miles deep

1 mole of marbles: 116 times the size of Mt.

Everest

1 mole of pennies stacked up: distance to

the moon 1.55x1012 times

Is equal to 18 mL

…that’s how small

molecules really are!

Mole of iron = 55.85 g Mole of sulfur = 32.07 g It’s not individual grains of iron or

sulfur! Where are these numbers coming

from? …Periodic table!!!

Definition: the mass of one mole of a substance (Units: g/mol)

AKA the atomic mass on the periodic table

Round all elements masses to two decimal places

Exception: Hydrogen gets 3 decimal places

12.01 grams of Carbon and 1.008 gram of Hydrogen contain the same number of atoms

6.022 x 1023 atoms

The gram atomic mass of any two elements must contain the same number of atoms

Carbon = 12.0107 g/mol

Rounds to 12.01 g/mol

Oxygen = 15.9994 g/mol

Rounds to 16.00 g/mol

Hydrogen = 1.00794 g/mol

Rounds to 1.008 g/mol

What if you have a compound?

1 mole NaCl = 58.44 g/mol

NaCl

1 Na = 22.99 g/mol

+ 1 Cl = 35.45 g/mol

58.44 g/mol

What is the molar mass of sugar?

C12H22O11?

C12H22O11

12-C x 12.01g/mol = 144.12 g/mol

22-H x 1.008 g/mol = 22.176 g/mol

+ 11-O x 16.00 g/mol = 176.0 g/mol

342.30 g/mol

What is the molar mass of aluminum

sulfate?

Al2(SO4)3

Count the atoms!

Al = 2, S = 3, O = 12

Al2(SO4)3

2-Al x 26.98 g/mol = 53.96 g/mol

3-S x 32.07 g/mol = 96.21 g/mol

+ 12-O x 16.00 g/mol = 192.0 g/mol

342.17 g/mol

1. Gram Atomic Mass (gam)

2. Gram Molecular Mass (gmm)

3. Gram Formula Mass (gfm)

GAM: molar mass of an atom or element

Ex) Fe, Cu, P, S… GMM: molar mass of a molecule (nonmetal compound)

Ex) H2O, Cl2, O2, CO2, P2O5

GFM: molar mass of a formula

unit (ionic compounds)

Metal-nonmetal compounds

Ex) NaCl, FeBr3, Zn3(PO4)2

If one mole of water is 18.02 g, how much

would two moles weigh?

36.04 g

How much would 3 moles weigh?

54.06 g

If one mole of water is 18.02 g, how many

moles are in 90.1 g?

5 moles

Three types of Mole Conversions

Mole – mass (g)

Mole – representative particles

(molecules, atoms, formula units)

Mole – volume (L)

What is the mass of 4.00 moles of NaCl?

Given:

Unknown:

Conversion Factor:

Solve:

If one mole of Helium weighs 4.00

g/mol, how much does 2 moles weigh?

Can you weigh out something less dense

then air?

You CANNOT weigh a gas less dense

then air!

Gravity does not pull it down on the

scale!

The volume of a mole of a gas is much more

predictable than that of a liquid or solid

That is, under the same physical conditions

(STP)

Standard temperature and pressure

▪ Standard Temperature is 00C

▪ Standard Pressure is 1 atm

At STP, 1 mole of any gas will occupy a

volume of 22.4 L

22.4 L is known as the molar volume of a gas

What does it mean?

It means that 22.4 L of any gas at STP contains

6.02 X 1023 representative particles of that gas.

Standard Temperature and Pressure (STP): specific conditions that can be reached in the lab

Standard Pressure: 1 atmosphere (atm)

Standard Temperature: 0oC or 273 K

How many moles are in 145.6 L of O2

gas at STP?

Given:

Unknown:

Conversion Factor:

Solve:

How many moles is 2.107x1024 molecules

of O2?

Given:

Unknown:

Conversion Factor:

Solve:

How many liters of gas are 3.612x1024

atoms of Neon gas at STP?

Given:

Unknown:

Conversion Factor:

Solve:

Temperature = a measurement of the average kinetic energy of the particles of an object

Scales

Fahrenheit (oF)

Celsius (oC)

Kelvin (K)

Boiling

Point

Room

Temp

Freezing

Point

Absolute Zero – coldest possible temperature where all movement stops, theoretical value

0 Kelvin = -273.15 oC = -459 oF

Equations:

K = oC +273 oC = K – 273

Example:

What is 98 K = -175 oC?

What is 159 oC = 432 K?