Chapter 4
Atomic Structure
In order to understand matter, we must consider its structure.
4.1 Defining the Atom
400 BC
Democritus - a Greek philosopher
·
·
·
late 1700's
John Dalton - an English scientist
·
·
Conservation of Matter
The Law of Constant Composition
states:
Example:
A
5.00g
10.0g
+
B
2.00g

AB
AB
21.0g

A
+
B
also:
In terms of percent:
The compound AB is composed of:
Example:
water
100. g

hydrogen
11.2 g
+
oxygen
Once the composition of a compound is known, any combination of reactants may be predicted.
How much oxygen is needed to react with 25.0 g of hydrogen to produce water?
How much oxygen is needed to form 38.0 g of water?
111.1 g of calcium chloride contains 71.0 g of chlorine. How much calcium and how much chlorine are needed to
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form 25.0 g of calcium chloride?
How much calcium chloride will be formed by the combination of 10.0 g of chlorine and 10.0 g of calcium?
Dalton’s Atomic Theory




All elements are composed of tiny indivisible particles called atoms
Atoms of the same element are identical. The atoms of any one element are different from those of any
other element.
Atoms of different elements can combine with one another in simple whole number ratios to form
compounds.
Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element
are not changed into atoms of another element by chemical reaction.
4.2 Structure of the Atom
Subatomic Particles

The idea that atoms were indivisible lasted for about 100 years.

Today we know that all matter can be broken down into particles more fundamental than atoms

Dozens of subatomic particles have been discovered

We will be concerned with three of these:
Electrons (e-)
·
negatively charged subatomic particles
·
first subatomic particle to be discovered
·
discovered in 1897 by the English physicist Joseph J. Thomson (cathode rays)

found that they were present in all types of matter
·
Robert Millikin determined the mass and charge of the electron using oil drops and an electric field

had a mass almost two thousand times lighter than a hydrogen atom

mass of electron = 9.11 x 10-28 g
p. 104 Fig.4.4
+
Protons (p )
·
discovered in 1886 by Eugen Goldstein

positively charged subatomic particle

1840 times heavier than the electron

mass of proton = 1.67 x 10-24 g
o
Neutron (n )

subatomic particle with no charge

discovered in 1932 by James Chadwick

same mass as the proton
see summary table 4.1 p.106
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Atomic Nucleus
1911 - Ernest Rutherford
·

The nucleus contains the protons and neutrons
Rutherford’s picture of the atom:
Atoms differ from one another in the number of electrons, protons and neutrons
4.3 Distinguishing Among Atoms
Three Principle of Atomic Structure
Atomic Number

The characteristic of an atom of an element is the number of protons in its nucleus.
The atoms of every element have a different number of protons
All H atoms have one proton
All O atoms have eight protons
This number is called the atomic number
Defined:
Symbol is:

Isolated atoms are electrically neutral.
number of protons = number of electrons
Mass Number

Different atoms of the same element may have different numbers of neutrons
mass number
Symbol for mass number is:
·
is a number not a mass
The composition of atoms are indicated in two ways:
Nuclear Symbols:
A
ZX
30
8
5Br
Mass Number:
hydrogen-1
3
helium-4
silver-108
Isotopes of Elements
Atoms of the same element may be different

must have the same number of protons and electrons

neutrons may be different
·
mass my be different
Examples
All oxygen atoms have 8 protons
Some oxygen atoms have 8 neutrons: A = 16
Some oxygen atoms have 9 neutrons: A = 17
Some oxygen atoms have 10 neutrons: A = 18

all have different masses!

the distribution of these atoms is not even, but is consistent from one sample to another
O-16 is the most common (99.759%) O-17 (.037%)
O-18 (.204%)
Atoms that have the same number of protons, but different numbers of neutrons are called isotopes
Nuclear symbols for the three isotopes of oxygen:
Hydrogen also exits in three common isotopes. These are the only isotopes which have their own unique names.
1H – hydrogen
21H
31H
Atomic Masses
·
measured in amu’s (atomic mass units)
The standard for atomic masses has changed over the years
1.
Hydrogen - used by Dalton
2.
Oxygen
3.
Carbon-12
One amu is defined as: “one twelfth the mass of a carbon-12 atom”
A carbon-12 atom:
One would expect that atomic masses to be whole numbers.
If all atoms are made up the same particles, their masses should be simple multiples of each other.
But, elements are made up of mixtures of two or more isotopes.
The RAM reflects this.
The RAM is the “weighted” average of the masses of the isotopes of an element.
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Calculating the Atomic Mass of an Element
The weighted average reflects the abundances
Example:
Element X exists as two isotopes
mass = 10.012 amu
19.91%
mass = 11.009 amu
80.09%
To find the weighted average: multiply the mass of each isotope by its natural abundance-add products
The Periodic Table-A Preview
Periodic Table:
·
·
·
·
arrangement of elements in which elements are separated into groups
these groups are based on a set of repeating properties
horizontal rows are called periods (7)
vertical columns are called groups/families (18)
Homework: 4,7,17,18,20,23,25,26,27,32,34,36,38,39,41,42,43,44,45,46,47,48,49,50,51,52,54,61,63,67,72,73,81
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