Chapter 3
ATOMS: THE BUILDING BLOCKS OF MATTER
Section 1
FROM PHILOSOPHICAL IDEA TO SCIENTIFIC THEORY
Foundations of Atomic Theory
 Particle
Theory of Matter
Democritus
in 400 B.C.
Stated
that nature’s basic particle was the atom
(“indivisible” in Greek).
Aristotle
Believed
all matter was continuous (could be divided
forever), and did not believe in atoms.
Neither
had experimental evidence to support
their claims.
Fast Forward to

th
the18
century
Three laws discovered due to improved instrumentation and carefully observed
chemical reactions

Law of conservation of mass—Mass is neither created nor destroyed during an ordinary
chemical reaction or physical change.

Law of definite proportions—A chemical compound contains the same elements in exactly
the same proportions by mass regardless of the size of the sample or source of the
compound.


Example: Table salt will always consist of 39.34 % Na and 60.66% Cl.
Law of multiple proportions—If 2 or more compounds are composed of the same two
elements, 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.

Example: Carbon and Oxygen can combine to form CO2 or CO.
Practice (law of conservation of mass)

If 3.5 g of X reacts with 10.5 g of Y to form the compound
XY, what is the percent by mass of X in the compound?

If 40 g of X reacts with 35 g of Y, what is the mass of the
product XY?

If 3.5 g of X reacts with 10.5 g of Y to form the compound
XY, how many grams of Y would react to form XY2? What
would be the final mass of XY2?
Practice (law of definite proportions)

2 unknown compounds are tested. Compound 1 contains 15.0g of
hydrogen and 120.0g oxygen. Compound 2 contains 2.0g of
hydrogen and 32.0g oxygen. Are the compounds the same?
Practice with Multiple Proportions

Three compounds containing K and O are compared. Analysis
shows that for each 1.00 g of O, the compounds have 1.22 g, 2.44
g, and 4.89 g of K, respectively. Show how these data support the
law of multiple proportions.

First: Find ratios of K for the different compounds.

If the ratios are whole numbers, it supports the law.
More practice with Multiple
proportions…

In 100 g of the compound A there are 57.1 g O and 42.9 g C. In 100 g of the compound
B, there are 72.7 g O and 27.3 g C. Show how this data supports the law of multiple
proportions.
Recall

1.) If 13 g of X reacts with 45 g of Y to form the compound XY, what is the percent by
mass of X in the compound?

2.) Two unknown compounds are tested. Compound 1 contains 32.6 g of hydrogen
and 167.4 g of Carbon. Compound 2 contains 8.0 g of hydrogen and 24.0 g of
carbon. Are the compounds the same?
Dalton’s Atomic Theory

John Dalton—proposed Atomic Theory
 Accounted
for all three laws: conservation of mass, definite proportions, and
multiple proportions.
 Dalton’s
 All
Atomic Theory consists of the following statements:
matter is composed of atoms.
 Atoms
of a given element are identical in size, mass, and all other properties.
 Atoms
cannot be subdivided, created, or destroyed.
 Atoms
of different elements combine in simple whole-number ratios to form
chemical compounds.
 In
chemical reactions, atoms are separated, combined, or rearranged.
Modern Atomic Theory

Advances in instrumentation have allowed some aspects of
Dalton’s theory to be proven incorrect.
 Example:
The thoughts that atoms are not divisible into smaller particles
and have the exact same mass are incorrect.

The two most important aspects of Dalton’s theory still hold true
 All
matter is composed of atoms.
 Atoms
of any one element differ in properties than atoms of another
element.
Section 2
THE STRUCTURE OF AN ATOM
Discovery of the Electron
 Scientific
advances allowed for the discovery of smaller
particles that made up atoms.
 William
Crookes—Passed and electric current through a cathode
ray tube.

Noticed that current passing through the tube produced a stream of
glowing particles (particles were traveling from the negative end of the tube
(cathode) to the positively charged end (anode).
Two conclusions resulted from experiments:
1.) Cathode rays were deflected by a magnetic field
in the same manner as wire carrying an electric current
(negatively charged).
2.) The rays deflected away from a negatively
charged object.
Charge of Cathode Ray
JJ Thomson, 1900

Proved Dalton’s solid atom wrong when he discovered the
particles had a mass and it was much less than the mass of a
Hydrogen atom.

Concluded that the electron has a very large charge-to-mass
ratio

Conducted Cathode-ray experiments to prove that atoms are
divisible.

Proposed plum pudding model for the atom.
Milikan and the oil drop experiment
• Meaured mass of droplets
• Applied charge to droplets through
x-rays.
• Applied electric voltage to top and
bottom plates
• Measured the amount of voltage it
took to keep the oil droplets
suspended.
• Was able to measure the charge of
a single electron!
http://www.youtube.com/watch?v=EV1owO1H2dA
Ernest Rutherford

Discovered the nucleus of an atom with his gold-foil experiment.

http://www.youtube.com/watch?v=XBqHkraf8iE

Concluded that the volume of the nucleus was very small compared to the total volume
of the atom.
Sizes of Atoms

Expressed in picometers (pm).

1 pm = 1*10-12 m = 0.000000000001m
Section 3
COUNTING ATOMS
Atomic Number
The number of protons in each atom of a given element
• Unique to each element and does not change
• Used to identify elements.
Practice
1. How many protons and electrons are in each atom?
 Radon,
Rn
 Titanium,
Ti
2. An atom of an element contains 66 electrons. Which
element is it?
3. An atom of an element contains 14 protons. Which element
is it?
Isotopes and Mass Number

The amount of protons and electrons are constant in all neutral
atoms of an element, but the number of neutrons can vary

Isotopes– atoms of the same element that have different numbers
of neutrons

Example: Hydrogen atoms have three different isotopes.

Protium—1 proton/0 neutrons (accounts for 99.985% of all hydrogen atoms).

Deuterium—1 proton/1 neutrons (0.015%)

Tritium—1 proton/2 neutrons (radioactive, and can be produced artificially)

Most elements contain mixtures of isotopes. Tin (Sn) has 10 stable
isotopes, the most of any element.

Mass Number--The number of protons and neutrons that make up
the nucleus of an isotope. Since electrons are so small, their mass is
considered insignificant.
Designating Isotopes
Isotopes are usually identified by their mass number.
 Two
Methods:
 Hyphen
notation—has hyphen after name of element,
followed by the mass number.
 Example:
Tritium is hydrogen-3 since it has 2 neutrons and 1
 Example:
Uranium-235
proton.

Nuclear symbol
 Mass
number written as superscript
 Atomic
number written as subscript
 Followed
by element symbol
In both methods…
 The
number of neutrons is found by subtracting the
atomic number from the mass number.
 Example:
 (Mass
 235
Uranium-235 or 23592U
#) – (Atomic #) = (# of neutrons)
(protons + neutrons) – 92 protons= 143 neutrons
Notation Practice
Determine the number of protons, neutrons, and electrons in
each of the following:
1.
2.
3.
1.
Neon-22
2.
Iron-57
3.
64
30𝑍𝑛
Write each of the following in hyphen notation:
1.
p = 12
n = 14
e = 12
2.
p = 28
n = 25
e = 28
Write each of the following in nuclear symbol notation:
1.
p = 38
n = 39
e = 38
2.
p=4
n=4
e=4
Pop quiz

What element contains 48 protons and 48 electrons?

How many protons does Krypton, Kr, contain?

How many neutrons are in the following isotopes:
 194
78𝑃𝑡
12
 6𝐶
Nitrogen-15
Relative Atomic Masses

Masses of subatomic particles are extremely small and hard to
work with—scientists created a new unit

Atomic mass units (amu) – based on a standard of carbon-12
that has a mass of 12 amu
 Neutron
 Proton
– 1.008665 amu = 1.66x10-24g
– 1.007276 amu
 Electron
– 0.0005486 amu
Average Atomic Mass

Avg. atomic mass – the weighted average of the masses of the isotopes
of that element

Avg. atomic mass = (mass x abundance)isotope1 + (mass x
abundance)isotope2 + …

To find abundance – divide the percentage by 100 (all abundances
must be decimals)

Atomic mass can help you determine which isotopes of that element is
the most abundant.
Atomic Mass Practice
1.
Boron has two naturally occurring isotopes: boron-10
(abundance 19.8%, mass 10.013 amu) and boron-11
(abundance 80.2%, mass 11.009 amu). Calculate the
atomic mass of boron.
2.
Nitrogen has two naturally occurring isotopes, N-14 and
N-15. Its atomic mass is 14.007. Which isotope is more
abundant? Explain.
Calculate the average atomic mass of lithium, which occurs as
two isotopes that have the following atomic masses and
abundances in nature: 6.017 amu, 7.30% and 7.018 amu,
92.70%.
 What
is the atomic mass of hafnium if, out of every 100
atoms, 5 have a mass of 176, 19 have a mass of 177, 27
have a mass of 178, 14 have a mass of 179, and 35
have a mass of 180.0?
 Iodine
is 80% 127I, 17% 126I, and 3% 128I. Calculate the
average atomic mass of iodine.

Magnesium has three naturally occuring isotopes. 78.70% of
Magnesium atoms exist as Magnesium-24 (23.9850 g/mol),
10.03% exist as Magnesium-25 (24.9858 g/mol) and 11.17% exist
as Magnesium-26 (25.9826 g/mol). What is the average atomic
mass of Magnesium?
Relating Mass to Number of Atoms
 The
Mole (mol)
 Since
atoms are so tiny, we count them in groups
 Same
as 1 dozen = 12 donuts
 Mole
= SI unit for amount of substance.
 The
amount of substance that contains as many particles as there
are atoms in exactly 12 g of carbon-12.
 Avogadro’s
number
 The
number of particles in exactly 1 mol of a pure
substance.
 6.022
× 1023 atoms in 1 mol of any substance.
Practice

How many atoms are in 5.76 moles of a pure substance?

How many moles are in 8.56 × 1026 atoms?
Molar Mass

The mass (in grams) of 1 mol of a pure substance.

Units are g/mol

Equal to the sum of the atomic masses of each element of
a compound.

Example:
 K2CO3

2(39.10 g/mol) + 1(12.01 g/mol) + 3(16.00 g/mol) = 138.21 g/mol
K
C
O
Gram/Mole Conversions
 Use
molar mass to convert moles to grams
and vice versa.
 Practice:
 Convert
7.00 g of He to moles.
 Convert
5.67 mol of Xe to grams.
 Convert
65.4 g of K2CO3 to mol.
 Convert
5.21 mol of NaCl to grams.
Recall

Find the molar masses of the following compounds:
 SrCO3
 C13H18O2
 Sn2(SO4)4

A sample of caffeine contains 96.1 g of carbon for every
10.1 g of hydrogen. If another sample of caffeine contains
30.0 g of carbon, how many g of hydrogen does it contain?
More Practice

How many atoms are in 4.56 g of Na2CO3?

What is the mass of 5.67 × 1032 atoms of NaOH?
Chapter 3 Vocabulary

Law of Conservation of Mass

Mass Number

Law of Definite Proportions

Nuclide

Law of Multiple Proportions

Atomic Mass Unit

Atom

Average Atomic Mass

Nuclear Forces

Mole

Atomic Number

Avogadro’s number

Isotope

Molar Mass