Atomic Theories
Learning Goals:
• Describe the research of
various scientists and explain
how they reached their
conclusions
• Define isotopes and solve
problems using atomic # and
mass #
Summary of Atomic Theories
• The first mention of what makes up matter was from the
Greeks in 600BC.
• The word “atom” comes from the Greek word meaning
uncuttable or indivisible. They believed that matter could be
cut into smaller and smaller units until eventually you would
get the smallest unit of matter.
• In the early 1800s, an English scientist named Dalton
investigated these ideas further.
History of the Atom
• Dalton (1808) – suggested that atoms are solid spheres, like
pool balls. He proposed that;
• All matter is made up of tiny, indivisible particles called
atoms
• Atoms cannot be sub-divided or changed into other particles.
• All atoms of an element are identical (size, shape, mass, etc.)
• Atoms of different elements are different
• Atoms are rearranged to form new substances in chemical
reactions, but they are never created or destroyed
• Atoms combine with each other in a ratio of small whole
numbers (1:1, 1:2, 2:2, 2:3)
History of the Atom
• Thomson (1897) – used a cathode ray tube to discover a new
particle inside the atom. He suggested that these particles are
small and negatively charged. He called them ‘electrons.’
The history of the atom
Pudding
POSITIVE charge
ELECTRONS
Negative charge
History of the Atom
• Rutherford (1909) – conducted gold foil experiments and
found that some positively charged particles were deflected.
He hypothesized that the particles must have come into
contact with a small, dense, positively charged nucleus.
• In 1920, he proposed that the nucleus must contain positively
charged ‘protons.’
• Rutherford Experiment
History of the Atom
• Bohr (1913) – experimented
with hydrogen. When
applying electricity to
hydrogen, he saw that the
electrons became ‘excited’
and they released specific
colours of light. He
proposed that electrons
must ‘orbit’ the nucleus in
different energy levels.
When electrons jump from
a higher to a lower energy
level, they emit light energy.
Bohr
History of the Atom
• Bohr postulated that;
• Electrons orbit at different radii from the nucleus and have
different energies
• There is a maximum number of e- allowed in each orbit
• The energy of e- increases as the distance from the nucleus
increases
• Lower energy orbits need to be filled up before higher
energy orbits.
• It was later discovered that the number of e- allowed in
each shell is calculated by 2n2:
• 2, 8, 18, 32 (simplified to 2, 8, 8, 8)
History of the Atom
• Schrödinger (1926) – published the Schrödinger wave equation to
describe electrons as particles AND waves. He proposed that
electrons do not exist in exact locations in orbits but instead exist in
defined regions of space called ‘electron clouds’ or ‘orbitals.’
The history of the atom
Dalton
Thomson
Rutherford
Bohr
Wave
Mechanical
Atoms and Atomic
Structure
The Nucleus
• The nucleus was found to contain 2 types of particles:
• Protons (p+): actual mass of 1.67x10-24g, +1 charge
• Neutrons (n0): actual mass of 1.67x10-24g, neutral charge
• In order to simplify calculations, protons and neutrons were given
a relative atomic mass of 1 amu (atomic mass unit.)
• The number of protons an element has is unique to that element
and is represented by the atomic number on the Periodic Table.
Bohr-Rutherford Diagrams
• Bohr-Rutherford diagrams combine the work of all
these scientists to show the arrangement of subatomic
particles (protons, neutrons and electrons) in an atom.
• In the centre, are the
number of protons and
neutrons.
• The number of electrons
are drawn on the outer
shells.
• The first shell holds 2
electrons, the second holds
8 and the third holds 8.
Lewis Diagrams
• Dots around the chemical symbol represent the valence
electrons
• When there are two dots beside each other, they are called
lone pairs of e-.
• When there is only 1 dot, it is called an unpaired e-. These are
more likely to be involved in bonding.
Atomic Number & Mass Number
• Atomic Number – the number of protons in one atom of an element.
In a neutral element, this will be the same as the number of
electrons.
• Mass Number (a.k.a. Atomic Mass) – the sum of the protons and
neutrons in one atom of an element.
• # of neutrons = Mass number – Atomic number
Practice
1.
2.
3.
4.
5.
6.
How many electrons are in 1 atom of F? Ans: 9
How many protons are in 1 atom of Na? Ans: 11
How many protons & neutrons are in 1 atom of N? Ans: 14
How many neutrons are in 1 atom of Al? Ans: 14
How many neutrons are in 2 atoms of C? Ans: 12
Draw a Bohr-Rutherford diagram for Neon.
Neutrons in the atom
– weight of the nucleus in amus.
Protons + Neutrons = Mass Number
Remember, electrons weigh nothing, so they do NOT contribute to
the Mass Number
+++
+++
24
6
18
12
Neutrons in the atom
Elements have different “versions” of themselves called
These will always have the same number of protons, but have
varying numbers of neutrons…thus, different mass numbers.
+++
+++
+++
+++
+++
+++
Carbon-12
Carbon-13
Carbon-14
12C
13C
14C
C-12
C-13
C-14
What’s happening to the proton count? The neutron count?
.
Isotopes
• Isotopes are atoms of the same element that have different
numbers of neutrons.
• Remember, they still have equal numbers of protons and
electrons so before they combine or bond with any other
atoms, they will be neutral.
• A different number of neutrons will affect the mass of the
atom.
Radioisotopes
• Isotopes that have unstable nuclei decay into more stable isotopes
• Instability is caused by many positively charged protons in a
densely packed nucleus (very high repulsive force)
• Neutrons help to attract protons and stabilize atoms with large
atomic mass
• Neutrons can decay into an electron and a proton, the electron is
‘emitted’ and the new element contains one extra proton
Calculating Average Atomic Mass
• Look at the periodic table and find the atomic mass for
Mg.
• Why isn’t it a whole number?
• In a sample of natural magnesium, 78.7% will be
magnesium-24, 10.1% will be magnesium-25 and 11.2%
will be magnesium-26.
• To find the average atomic mass of any one atom in the
sample, we must calculate the average, taking into
account the percentage of each isotope in the sample.
Average relative atomic mass
• Formula =
(Isotopic abundance x relative atomic mass of isotope 1) +
(isotopic abundance x relative atomic mass of isotope 2)
+(isotopic abundance x relative atomic mass of isotope 3)….
• Isotopic abundance is taken from percent form and turned
into a decimal
• Eg) 99.757% = 0.99757
Example Problem
• Calculate the average relative atomic mass of oxygen:
Isotope
Isotopic Abundance
Atomic Mass (u)
16O
0.99757
16.000
17O
0.00038
17.000
18O
0.00205
18.000
• Average relative atomic mass
=(0.99757)(16.000) + (0.00038)(17.000) + (0.00205)(18.000)
=15.961 + 0.0065 + 0.0369
=16.000 u
*last step is addition, therefore use addition rule for sig figs
Practice Problem
• Naturally occurring silver has two isotopes. Calculate the
average atomic mass of silver if 47Ag107 has a mass of
106.9u and a relative abundance of 51.8% and 47Ag109
has a mass of 108.9u and a relative abundance of 48.2%.
ANSWER:
AAM = 0.518(106.9) + 0.482(108.9) = 107.8638u
The atomic mass of silver is 107.9u
Practice Problem
Given the following abundances for various magnesium
isotopes, calculate the average relative atomic mass; 78.7%
magnesium-24, 10.1% magnesium-25 and 11.2% magnesium26.
Answer:
Solve: 0.787(24) + 0.101(25) + 0.112(26) = 24.325
Statement: The atomic mass of magnesium is 24.3 u.
Isotopic Abundance
• Isotopic Abundance is the amount of a given isotope of an
element that exists in nature, expressed as a percentage of
the total amount of this element.
• For example, in a sample of copper metal:
Isotope
Mass (amu)
Isotopic
Abundance
Cu-63
62.93
69.2%
Cu-65
64.93
30.8%
• Which isotope is more common?
Example Problem
• Boron exists as two naturally occurring isotopes: 5B10 (10.01u)
and 5B11 (11.01u). Calculate the relative abundance of each
isotope.
*use the Periodic Table to find the AAM for boron
AAM = x(isotopic abundance B10) + (1-x)(isotopic abundance B11)
10.81 = x(10.01) + (1-x)(11.01)
10.81 = 10.01x + 11.01 – 11.01x
11.01x – 10.01x = 11.01 – 10.81
x = 0.20 = 20% B10
(1-x) = 0.8 = 80% B11
Therefore the isotopic abundance of B10 is 20% and the isotopic abundance
of B11 is 80%.
Atomic Structure &
Chemical Bonding
The Octet Rule
• When atoms combine, they achieve a stable
arrangement with a full valence shell or 8 outer electrons
(2 for hydrogen.) This is called the ‘octet rule.’
• When atoms loose or gain electrons, they become ions.
• ION – A charged particle
The Formation of Ions
Cations
• Formed when an atom
looses an electron
(metals)
• They have the same
name as their element
• A sodium atom becomes a
sodium ion.
• An aluminum atom
becomes an aluminum ion.
Anions
• Formed when an atom
gains an electron (nonmetals)
• They are named by
replacing the end of the
elements name with –
ide.
• A chlorine atom becomes a
chloride ion.
• A sulfur atom becomes a
sulfide ion.
Example: Indicate the number of protons, neutrons
and electrons in each atom or ion
Element
Atomic Number
of atom
Atomic Mass # of protons
of atom
N
7
14
Ca2+
20
40
Br1-
35
80
# of
neutrons
# of
electrons
Example: Indicate the number of protons, neutrons
and electrons in each atom or ion
Element
Atomic Number
of atom
Atomic Mass # of protons
of atom
N
7
14
Ca2+
20
40
Br1-
35
80
7
# of
neutrons
14 – 7 =
7
# of
electrons
7
Example: Indicate the number of protons, neutrons
and electrons in each atom or ion
Element
Atomic Number
of atom
Atomic Mass # of protons
of atom
N
7
14
7
Ca2+
20
40
20
Br1-
35
80
# of
neutrons
# of
electrons
14 – 7 =
7
7
40 – 20 20 – 2 =
= 20
18
Example: Indicate the number of protons, neutrons
and electrons in each atom or ion
Element
Atomic Number
of atom
Atomic Mass # of protons
of atom
N
7
14
7
Ca2+
20
40
20
Br1-
35
80
35
# of
neutrons
# of
electrons
14 – 7 =
7
7
40 – 20 20 – 2 =
= 20
18
80 – 35 35 + 1 =
= 45
36
Game Time:
• “Build an Atom”
Practice Questions
• Check Unit Outline:
• Page 14, Q’s 1, 2, 5, 6
• Page 19, Q’s 1, 2, 8-10
• Page 21, Q’s 3-6, 9-11
• WS: Isotopes Practice
• WS: Average Atomic Mass