Chapter 5 Outline
5.1 Dalton’s Atomic Theory
1808 – John Dalton developed concept of the atom. Dalton’s theory is based on the Law of Definite
Composition (Section 2.6) & the Law of the Conservation of Mass (Section 2.9).
Dalton’s theory is as follows (see Summary Box and Figure 5.1 on page 124 & Figure 5.2 on Page 125)
1. Every element is made up of tiny particles called atoms.
2. Atoms cannot be broken down into smaller particles. They can’t be created or destroyed.
3. All atoms of the same element are identical.
4. Atoms of one element are different form atoms of another element.
5. Atoms of one element can combine with atoms of other elements to form compounds. This
usually occurs in ratios of whole numbers.
Dalton’s theory was not immediately accepted. Led to the Law of Multiple Proportions.
When atoms from 2 different elements combine to form more than one compound, the different
weights of one element that combine with the same weight of the other element are in ratio of whole
numbers.
 Example Carbon & Oxygen:
o Carbon monoxide always 1 carbon & 1 oxygen
o Carbon dioxide—Same elements different ratio: 1 carbon & 2 oxygen
5.2 Subatomic particles
Dalton’s theory was soon challenged. Work of Faraday & Crookes led to the discovery of the electron.
All atoms have subatomic particles (See Table 5.1 on Page 126) (More than 100—we are only going to
concentrate on 3)
Subatomic
particle
Symbol
Electron
e
Proton
p or p
Neutron
n or n
—
Charge
Mass (g)
Mass (amu)
Location
Discovered
1897 by
Thompson
1- (-1)
9.109 x 10
—28
~0
+
1+ (+1)
1.673 x 10
—24
~1 (1.00728)
Orbitals
outside
nucleus
Nucleus
0
0 (nc)
1.675 x 10
—24
~1 (1.00867)
Nucleus
1919 by
Rutherford
1932 by
Chadwick
5.3 The Nuclear Atom
1911—Rutherford performed an experiment and came up with the Nuclear Model of the Atom (See
Figures 5.3 Page 127)
Their conclusions were: (See Summary box Page 127)
1. All atoms have an extremely small, dense nucleus.
2. The nucleus contains all of the positive charge and the majority of the atom’s mass.
3. The nucleus is surrounded by a larger volume of almost empty space that makes up the
remainder of the atom.
4. The space outside of the nucleus contains the electrons, whose charge equally balances the
positive charge of the nucleus.
Putting the atom in perspective: See Figure 5.4 on page 128.
If the nucleus of an atom was that size (nearly 10—10 m) then the diameter of the atom (where
the electrons are) would be 100 yards or the length of a football field.
Early scientists thought electrons simply traveled in circular orbits around the nucleus (like the planets
around the sun… This would be the planetary model of the atom).
5.4 Isotopes
We now know that all atoms of the same element are identical. They can have different masses.
We know the following:
 All atoms of the SAME element have the same # of protons. This is called the atomic number &
is represented by the symbol Z.
 Atoms of the same element that have different masses are called isotopes. They have the same
number of protons, but DIFFERENT numbers of neutrons.
 The total number of protons and neutrons an atom has makes up its atomic mass, represented
by the symbol A.
Thus: Mass number = # protons + # neutrons AKA A = Z + # neutrons.
Naming Isotopes: Elemental name followed by the mass number.
Examples:
Oxygen (8 p+ and 8 n0) = Oxygen-16
Carbon (6 p+ and 6 n0) = Carbon-12 while (6p+ and 8 n0) = Carbon-14
The Nuclear symbol is the chemical symbol plus the mass number & atomic number. It is written as:
Where Sy = Chemical symbol of the element
Mass number
Atomic number
A
Sy
Sy
Z
16
O
8
You can find the number of protons or neutrons based on the mass # & atomic number by simple math.
See Example on Page 129-130.
5.5 Atomic Mass
The mass of an atom is too small to be measured on a balance. However the ratios of specific amounts
(mass/weights) between two elements is the same as the ratios between their atomic masses. This led
to the development of a relative scale of atomic masses (aka atomic weight).
Since elements have several different isotopes, the masses of atoms are expressed in atomic mass units
(amu). 1 amu is exactly 1/12 the mass of a carbon-12 atom.
Hence 1 atom of carbon-12 = 12 amu. AND
1 amu = 1.66 x 10—24 g
Since many elements have 2 or more isotopes, the atomic mass of an element = average mass of all
atoms of that element as they occur in nature.
To determine the (average) atomic mass of an element you need to know a few things:
 Atomic mass of each isotope
 Fraction of each isotope in a sample
o See Table 5.2 on Page 132
Example 5.3 – Atomic Mass of potassium (K) using the data in Table 5.2
Potassium has 3 isotopes: potassium-39, potassium-40 and potassium-41.
 Potassium-39 has an atomic mass of 38.9637074 and makes up 93.2581% of all of the isotopes
of potassium.
 Potassium-40 has an atomic mass of 39.9639992 and makes up 0.0117% of all of the isotopes of
potassium.
 Potassium-41 has an atomic mass of 40.9618254 and makes up 6.7302% of all of the isotopes of
potassium.
Calculate the atomic mass of potassium.
First, convert the percent abundance of the isotope to a decimal. Then multiply it by the amu for that
isotope. Do this for each isotope.
0.932581 x 38.9637074 amu
=
36.3368 amu
0.000117 x 39.9639992 amu
=
0.00468 amu
0.067302 x 40.9618254 amu
=
2.7568 amu
Add values together =
39.0983 amu = accepted atomic mass
Remember you need to use proper significant figures!!!
5.6 The Periodic Table
1869 Mendeleev & Meyer independently discover that elements are arranged based on their atomic
masses & that certain properties repeat are regular intervals.
They arranged elements into tables. Elements with similar properties are in the same column or row.
This led to the first periodic table. This led to some discrepancies. In order to have proper groups, they
switched the order & this disrupted the order of atomic masses. (It was assumed that there were errors
in some of the masses). It was later discovered that the proper order is by atomic number NOT atomic
mass.
Also allows you to predict the properties of undiscovered elements. See table 5.3 on Page 134.
Refer to Figure 5.7 on Page 135 for the following:
Each box has the information for 1 element.
The horizontal rows are called periods. They are numbered rom top tpo bottom (but you usually don’t
see this).
 Periods vary in length.
 Period 1 only has 2 elements, Periods 2 & 3 has 8 etc.
o How many elements are in Periods 4, 5, 6, 7, 8?
Vertical columns are called groups or chemical families. Groups are identified by numbers across the
top.
 Difference in numbering between US and Europe led to IUPAC numbering system (see Figure 5.6
on page 135). Your book will use both sets of numbers. Example of Carbon= Group is 4A/14.
Two other regions hold elements w/ special classifications.
 Elements in A Groups (1, 2, and 13 -18) are Main Group elements or representative elements.
 Elements in B groups (3 to 12) are tranisition elements or transistion metals.
 The stair-step line between Atomic Numbers 4&5 in Period 2 and end between 84 & 85 in
Period 6 separates metals on the left form the nonmetals on theright.
 An element’s location is given by period and group numbers.
o Example which element is found in the 3rd period, Group 6a/16? Sulfur
5.7 Elemental Symbols & the Periodic Table
The periodic table contains a large amount of information. You will see how useful it is as the course
progresses.
You will need to learn/memorize the names and chemical symbols of 35 common elements. Please
refer to Figure 5.8 on Page 138 to help you.
Please note that the “periodic table” with the elemental symbols highlighted in yellow/brown are the
ones you will need to know for this course!!! You will learn the chemical symbols for compounds in a
later chapter.