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DEMOCRITUS
DALTON
THOMSON
RUTHERFORD
BOHR
DeBROGLIE
The Greeks - Democritus
• 450 B.C.
• Believes there is a
fundamental particle that
cannot be broken. ( won’t
subdivide to smaller
components)
• The particle is invisible /
indivisible
• Coins the term “atomos”
John Dalton
• 1810
• Believes atoms are invisible
and indivisible.
• All atoms of the same element
are alike in every respect –
especially mass.
• Atoms of different elements
are different in every respect –
especially mass.
• -- combine in whole number
ratios
• --Billiard Ball Theory
Early symbols representing
compounds in various simple
ratios. Dalton
J.J. Thomson
• 1897
• An atom consists of a positively charged globule.
• Negative particles, called electrons, are embedded in the
globule.
• These electrons move about in the globule.
• Jellied salad or raisin bun theory.
(Cathode Ray Tube)
JJ Thomson’s model of the atom
Positive charged mass with embedded electron particles
electrons
Positive mass
RUTHERFORD EXPT
Ernest Rutherford / Nuclear model
• 1908
• Performs his famous scattering experiment
HYPOTHETICAL
Observations
- Most alpha particles pass through the foil as if there
was nothing there
- A few veer off as they are slightly deflected
- A very few bounce back at the alpha gun
Conclusions
- Atoms are mainly empty space (over 99.9 %)
-The mass of the atom is concentrated in a core or
nucleus which is positively charged.
- Electrons, which are negative move about in a large
region of space outside the nucleus
Rutherford Model
Negative Electron
Positive Nucleus
Quantum Mechanical Model OR the Nuclear Atom
I. General Information
1. Atoms have a very dense core called the nucleus.The
nucleus contains two sub-atomic particles:
A) Protons which have a positive charge
B) Neutrons which do not carry a charge (neutral)
2. The electrons, which carry a negative charge, are in a
state of constant motion in a relatively large region of
space away from the nucleus.
3. Atoms are mainly empty space. (over 99.9 %)
NUCLEAR MODEL
RUTHERFORD
- DISCOVERED NUCLEUS
-POSITIVE CHARGED PROTON
-ELECTRONS ORBIT OUTSIDE
Described as - - SOLAR MODEL
Neils Bohr
Energy Levels
Electrons are located
at different energy
levels due to energy of
the electrons
Quantum model
Orbitals s p d f
s – 2 electrons
p – 6 electrons
d – 10 electrons
F – 14 electrons
[ Electron configuration ]
Neils Bohr
ENERGY LEVELS - 7 Levels for 7 periods
Electron orbits / electron configuration
s p d f orbits / electron location
ELECTRON LOCATION
• Electrons are located in specific energy levels
• These energy levels can be thought of as globe shaped
regions of space surrounding the nucleus.
• Electrons in the inner level have the lowest amount of
energy and energy increases as the levels move
outward.
• Electrons enter and fill these levels in order and each
level has a limit to how many electrons it can hold.
• Level 1 holds 2 electrons
• Level 2 holds 8 electrons
• Level 3 holds 8 electrons
• Level 4 holds 18 electrons
• Level 5 hold 18 electrons
• Level 6 holds 32 electrons
• Level 7 holds 32 electrons
Shows the 3 orbitals and how they overlap
DUST CLOUD MODEL
De Broglie - Model
Probability of finding electrons at
different levels
Electrons are particles / bundles of
Energy.
Dual principle theory – wave
mechanics / particle theory
PROBABILITY OF ELECTRON LOCATION (DOTS)
II. Summary of Atomic Particles
Particle
Symbol
Relative Relative Location
Mass
charge
Proton
H or p+
1
+1
Nucleus
neutron
n or n
1
0
Nucleus
electron
e or e-1
0
-1
Constant motion
outside nucleus
Relative => means with respect to each other
About 1864 ( 2000) electrons = 1 proton
III. Atomic Number
• This is used to identify the element(name and/or symbol)
• This number ALWAYS tells us the how many protons are
in the nucleus of atoms of that element.
• NO EXCEPTIONS
• Examples:
Carbon atoms (atomic # of 6) => 6 protons
Sodium atoms ( atomic # 11 ) => 11 protons
Osmium atoms ( ?
) => ? Protons
?????? Atoms (atomic # 47) => 47 protons
IV. Nuclear Charge
• This is the total positive charge present on the
nucleus.
• It is always positive because the protons are
the only charged particle in the nucleus.
• Therefore, the nuclear charge is equal to the
number of protons in the atom or ion. (equal to
the atomic number)
• Examples:
The nuclear charge on K is
+ 19
The nuclear charge on Ni
+ 28
V. ATOMIC MASS
• The mass number is the total mass in grams for a
constant and specific number of atoms of each element.
• This number is very large (because atoms are very
small) and is known as the mole number.
• 1.0 mole is equal to 6.02 x 1023 particles.
Examples:
The mass of one mole of Hydrogen atoms is 1.01 g.
Therefore 1.01 g/mol is the mass number for Hydrogen.
ATOMIC MASS OF
Zr
91.22 g / mol
MASS NUMBER -
ATOMIC MASS ROUNDED TO A
WHOLE NUMBER.
ATOMIC MASS ROUNDED 
MASS NUMBER
CARBON -
ATOMIC MASS OF
12.01 => 12 amu
SILVER -
ATOMIC MASS OF
107.87 => 108 amu
ANTIMONY - ATOMIC MASS OF
121.75 => 122 amu
THE NUMBER OF NEUTRONS IS CALCULATED AS:
MASS NUMBER - ATOMIC NUMBER = NEUTRONS
VI. Number of Neutrons
•
Since only protons and neutrons contribute to the mass
of an atom, the sum of these two particles must be
equal to the mass number.
•
To calculate the number of neutrons in a given atom:
MASS NUMBER – # OF PROTONS = # OF NEUTRONS
Examples:How many neutrons are in each of the following?
1. A carbon atom?
12 – 6 = 6
2. A zinc atom?
65 – 30 = 35
3. 44.96 Sc 21
Scandium has 45 – 21 = 24 neutrons
VII. ISOTOPES
All atoms of the same element are not identical.
Isotopes - 2 or more atoms of the same element that have
(same atomic number and same number of protons)
Different number of neutrons ===> Different mass
Hydrogen 1
1
1
H
1 proton
0 neutrons
Protium
Occur naturally
Hydrogen 2
2
1
H
1 proton
1 neutron
Deuterium
Hydrogen 3
3
1
H
1 proton
2 neutrons
Tritium
Man-made / synthetic
Please do these in your NOTE BOOK !!
A carbon isotope has 8 neutrons. What is its mass number?
Answer: 14
An isotope of zinc has a mass number of 62. How many
neutrons does it have?
Answer: 32
A given isotope has a mass number of 52 and contains
26 neutrons. What element is it?
Answer: Iron (Fe)
VIII. WRITING ISOTOPES
2 methods to represent the different isotopes of an element
1.) X(Y) where ‘X’ is the element symbol and ‘Y’ is
the mass number of that particular isotope.
Examples: H(1)
H(2)
H(3)
C(12) C(14) U(238) Zn(64)
A
Z
X where ‘X’ is the element symbol, ‘A’ is the
2.)
mass number of the isotope and ‘Z’ is the atomic
number
Examples:
14
6
65
30
C , Zn ,
238
92
90
38
U , Sr etc.
IX. CALCULATING AVERAGE MASS
• An example: A theoretical element
• Element J has 4 known isotopes. Their mass distribution
in nature is as follows:
1) J(40) = 30.0%
2) J(43) = 40.0%
( largest % in nature)
3) J(44) = 10.0%
( smallest % in nature)
4) J(50) = 20.0%
Calculate the average mass of this element ( J )
Take the % of each mass and add the results
together to get a total.
Each Isotope may occur in different % in nature
30/100 x 40g = 12.0 g
40/100 x 43g = 17.2 g
10/100 x 44g = 4.4g
20/100 x 50.0g = 10.0g
43.6 g
The average mass of element J is 43.6 g / mol
Remember – the total percent of any element must be
100%. If you are given numbers like 20% and 40% for
2 of the 3 isotopes… 100 – ( 20 + 50) = 30 % for the
third isotope.
IMPORTANT
****when completing tables and answering
questions on atomic theory,
ALWAYS USE THE GIVEN DATA FROM THE
QUESTION.
The values from the periodic table are used if
no other information is available.
X. NUMBER OF ELECTRONS
• In any neutral element , the number of electrons is equal
to the number of protons. (equal to the atomic number)
• That is:
[ p+
=
e-]
XI. IONS
Species where the number of protons is NOT equal
to the number of electrons.
That is:
[ p+
=
e- ]
These are charged particles known as ions
Type: II
It is also possible to remove one or more electrons
from an atom or ion. This will result in the formation of
a positive ion.
Energy must be supplied in order to remove each
electron.
REACTANTS
Na + energy →
Al + energy →
PRODUCTS
Na+ + eAl 3+ +
Mg + energy → Mg 2+ +
OIL <-> RIG
3e-
2e-
This kind of reaction is known as an OXIDATION reaction
**All ions are formed by the addition or
removal of electrons. NEVER CHANGE THE
NUMBER OF PROTONS.**
Type I:
It is possible for an atom to gain one or more electrons.
This results in the formation of a negative ion.
Energy is released as each electron is
accepted.
Examples:
F + e- → F- + energy
S + 2e- → S2- + energy
N + 3e- → N3- + energy
This kind of reaction is known as a REDUCTION reaction.
To help remember oxidation and
reduction !! (1 WAY)
OIL
OXIDATION IS LOSS
OF ELECTRONS
RIG
REDUCTION IS GAIN
OF ELECTRONS
2 ND
WAY
My name is
LEO the lion says GER
Oxidation - Reduction
Left side is reactants 
Right side is products
Na +
Na becomes what type of ion ? ( + or - )
That indicates that it lost electrons….which are the products
Losing electrons means gaining Energy
Na +
Energy
== >
Na +
+
Energy is a reactant
1e-
Sodium plus energy == > sodium ion and 1 electron
Predicting Ion Formation
• Stability in ions is represented by the total
number of electrons present.
• The stable numbers are those found in atoms of
the inert gases. (2, 10, 18, 36, 54, 86) These are
unreactive – that is they are stable.
• All other atoms will gain (become reduced) or
lose (undergo oxidation) in order to achieve one
of these numbers.
Predict the ion that each of the following atoms would form
in order to become stable. Write out the oxidation or
reduction reaction.
•
•
•
•
•
•
•
•
•
K
Cl
O
P
Ca
Ga
Cs
N
Br
ANSWERS
K + energy → K+ + eCl + e- → Cl- + energy
O + 2e- → O2- + energy
P + 3e- → P3- + energy
Ca + energy → Ca2+ + 2eGa + energy → Ga3+ + 3eCs + energy → Cs+ + eN + 3e- → N3- + energy
Br + e- → Br- + energy
Dmitri Mendeleyev
given credit for the first working
periodic table
1. The periodic table is arranged in horizontal
rows and vertical columns.
2. Each horizontal row is called a PERIOD or
SERIES.
3. Each vertical column is known as a GROUP
or FAMILY. ( same chemical characteristics)
Label the following on your blank table:
- Families (show numbers & Roman numerals)
- Periods (show number of each period)
- Alkali metals ------ - -- - ( group 1A )
- Alkali earth metals - - - ( group 2A)
- Transition elements - - - ( group B )
- Halogens ------ -- -- -- -- -- - - - (group 7A )
- Inert gases (noble gases)- - (group 8A )
VALENCE ELECTRONS
• All of the electrons located in the outermost (highest)
energy level are called valence electrons.
• For all ‘A’ elements, the group number equals the
number of valence electrons. Group I A to VIII A
•
Carbon is in group IVA and therefore has 4 valence
electrons.
• All transition (B family) elements have 2 valence
electrons.
Metallic Properties
Metallic properties decrease as we move left to right across
the table.
Metallic properties increase as we move from top to bottom
in the table.
The most metallic element is in the bottom – left of the
table and the least metallic element is in the top right.
ATOMIC SIZE
IN ANY FAMILY
As atomic number increases – Atomic size increases
This is due to more energy levels holding
electrons. Eg. Cs atoms are larger than Na
atoms because Na atoms have electrons in 3
energy levels holding electrons and Cs has 6
energy levels holding electrons
IN ANY PERIOD
As atomic number increases – Atomic size decreases
This is due to increased nuclear attraction on
electrons in the same energy level. Eg. A Br atom is
smaller than a Ca atom. 35 protons pulling on outer
electrons vs 20 protons pulling on outer electrons.
THE ELECTRON AFFINITY
ADDITION OF AN ELECTRON TO AN ATOM
ELECTRON ADDED / ENERGY IS LOST
MOVIE 1
MOVIE 2
IONIZATION
ENERGY
This is the amount of energy required to remove an
electron(s) from an atom or ion. (See formation of a
positive ion..)
IN ANY FAMILY
As atomic number increases – ionization energy decreases
This is a function of size. The outermost
electrons are further and further from the
nucleus and are easier to remove.
IN ANY PERIOD
As atomic number increases – ionization energy increases
This is due to increased nuclear attraction
on electrons in the same energy level.
BACKGROUND / Electron Configuration
1] Electrons are located in specific energy levels surrounding
the nucleus. These are numbered consecutively from 1 – 7,
starting from the nucleus and working outward.
2] Each energy level has a specific maximum capacity for
holding electrons.
3] Energy levels can be sub-divided into sub-levels (or subshells). These are identified with letters: s, p, d, f…etc.
4] These sub-levels also have maximum capacities.
5] Each sub-level are, in turn, divided into orbitals that hold 2
electrons each.
6] Levels and sub-levels fill systematically with electrons
starting from the inside (lowest energy) and working
outward (highest energy).
ELECTRON CONFIGURATION
• That is: level one completely fills before level 2 starts
filling…level 2 before level 3….etc.
• Note: Not all energy levels contain all sub-levels.
• Capacities: s – 2 electrons – 1 orbital 1 (2electrons)
p – 6 electrons – 3 orbitals 3 (2electrons)
d – 10 electrons – 5 orbitals 5 (2electrons)
f – 14 electrons – 7 orbitals 7 (2electrons)
1s2 / 2s22p6 / 3s23p6 / 4s23d104p6 / 5s24d105p6 / 6s24f145d106p6 / 7s25f146d107p6
1
2
3
4
5
6
7
32
32
ENERGY LEVELS
2
8
8
18
18
Number of electrons in each energy level
Notation
Energy levels
1
2
3
1 s 2 2 s 2 2 p 6 3s 2 3p 6
electrons
Orbitals / sublevel
Prime quantum number/ energy level
(refer to period on periodic table)
Writing Electron Configurations / Review
H: 1s1
He: 1s2
Li: 1s2 2s1
Ca: 1s2 2s2 2p6 3s2 3p6 4s2
Ag: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9
U: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14
5d10 6p6 7s2 5f4
Your turn --- make my day
N:
1s2 2s2 2p3
Na:
1s2 2s22p6 3s1
Br:
1s2 2s22p6 3s23p6 4s23d104p5
VALENCE ELECTRONS
Electrons in the last energy level / max 8
Bonding electrons / octets - only 8
F - 1s2 2s2 2p5
look at the last energy level
F – 1s2 2s2 2p5
Filled up
[ count only the s and p electrons
2 electrons
+
5 electrons
=
Cl - has 7 valence electrons
Cl - will gain 1 electron to become stable
Cl
becomes
Cl 1-
[ valence of 1- ]
7
Other Topics
•
•
•
•
Last electron added (valence electrons)
How to remember the sequence (two methods)
Grouping quantum numbers
Configurations of ions
Short Cut Version / Noble Gas Method
(Only use when instructions allow it)
Write the symbol for the last complete inert gas and then
continue the configuration from the next energy level.
Example: Write the Noble Gas configuration for Ra.
Ra: 86Rn – [ 7s2 ]
Using this format, write configurations for:
– [ 6s2 4f14 5d9 ]
Au:
54Xe
Ga:
2 3d10 4p1 ]
Ar
–
[
4s
18
Cf:
Sc:
86Rn
– [ 7s2 5f10 ]
2 3d1 ]
Ar
–
[
4s
18
Questions for Periodic Table Review
Which Element?
•
•
•
•
•
•
•
•
•
•
•
•
•
Is an alkali metal?
Is chemically inert?
Is in group IV A?
Is in period 2?
Has the largest mass number?
Is a transition element?
Forms a 3 – ion?
Forms a 2+ ion?
Tends to lose one electron?
Tends to gain 2 electrons?
Is Lithium?
Has a nuclear charge of + 20?
has 2 valence electrons?
• Are chemically similar?
• Has the largest radius?
• Has the lowest ionization
energy?
• Has valence e- in period 3?
• Last e- added was d1?
• Is the most metallic?
• Configuration ends in 4p2?
• Has 22 protons in nucleus?
• Has e- in 3 energy levels only?