Unit 2
Name
Symbol
Charge Relative Mass Actual Mass (g)
(amu)
Electron
e-
-1
1/1840
9.11x10-28
Proton
p+
+1
1
1.67x10-24
Neutron
no
0
1
1.67x10-24
•Atoms are measured in picometers, 10-12 meters
Hydrogen atom, 32 pm radius
• Nucleus tiny compared to atom
If the atom were a stadium, the nucleus would be a marble
• Radius of the nucleus is on the order of 10-15 m
• Density within the atom is near 1014 g/cm3
•Atomic Number (Z) = number of protons (p+) in the nucleus
Determines the type of atom
• Li atoms always have 3 protons in the nucleus, Hg always 80
• Mass Number (A) = number of protons + neutrons [Sum of p+
and nº]
Electrons have a negligible contribution to overall mass
• In a neutral atom there is the same number of electrons (e-)
and protons (atomic number)
•Every element is given a corresponding symbol which is
composed of 1 or 2 letters (first letter upper case, second lower),
as well as the mass number and atomic number
mass number
A
elemental symbol
atomic number
Z
E
ATOMIC NUMBER AND MASS NUMBER
4
2
He
Mass Number
the number of protons and
neutrons in an atom
Atomic Number
the number of protons in an
atom
Number of electrons = Number of protons
in a neutral atom
5
•Find the
number of protons
number of neutrons
number of electrons
atomic number
mass number
19
9
F
80
35
Br
184
74
W
Ions
 Cation is a positively charged particle. Electrons have
been removed from the element to form the + charge.
ex: Na has 11 e-, Na+ has 10 e-
 Anion is a negatively charged particle. Electrons have
been added to the atom to form the – charge.
ex: F has 9 e-, F- has 10 e-
•Atoms of the same element can have different numbers of
neutrons and therefore have different mass numbers
• The atoms of the same element that differ in the number of
neutrons are called isotopes of that element
1
1
H
2
1
H
3
1
H
Hydrogen-1
Hydrogen-2
Hydrogen-3
• When
naming, write the
mass number after the
name of the
element
Average = (% as decimal) x (mass1) + (% as decimal) x (mass2) +
(% as decimal) x (mass3) + …
Problem:
Silver has two naturally occurring isotopes, 107Ag with a
mass of 106.90509 u and abundance of 51.84 % ,and 109Ag
with a mass of 108.90476 u and abundance of 48.16 %
What is the average atomic mass?
Average = (0.5184)(106.90509 u) + (0.4816)(108.90476 u)
= 107.87 amu
Average Atomic Masses
• If not told otherwise, the mass of the isotope is the mass
number in ‘u’
• The average atomic masses are not whole numbers because
they are an average mass value
• Remember, the atomic masses are the decimal numbers on
the periodic table
More Practice Calculating Averages
• Calculate the atomic mass of copper if copper has two
isotopes


69.1% has a mass of 62.93 amu
The rest (30.9%) has a mass of 64.93 amu
• Magnesium has three isotopes




78.99% magnesium 24 with a mass of 23.9850 amu
10.00% magnesium 25 with a mass of 24.9858 amu
The rest magnesium 26 with a mass of 25.9826 amu
What is the atomic mass of magnesium?
Bohr
 Proposed electrons (e-) orbit around the nucleus in circular




paths
Said e- in a particular path have a fixed energy (energy
levels)
e- can go from any energy level to another by gaining or
losing a specific amount of energy = a “quantum of energy”
When e- absorbs a quantum of energy, it goes from it’s
ground state (where it’s normally found) to an excited state
The excited state is at a higher energy level
Bohr postulated that:





Fixed energy related to the orbit
Electrons cannot exist between orbits
The higher the energy level, the further it is away
from the nucleus
An atom with maximum number of electrons in the
outermost orbital energy level is stable (unreactive)
Think of Noble gases
Atomic Line Emission
Spectra and Niels Bohr
Niels Bohr
(1885-1962)
Bohr’s greatest contribution to science
was in building a simple model of the
atom. It was based on an
understanding of the LINE
EMISSION SPECTRA of excited
atoms.
Problem is that the model only
works for Hydrogen
Spectrum of White Light
Spectrum of
Excited Hydrogen Gas
Line Emission Spectra
of Excited Atoms
 Excited atoms emit light of only certain
wavelengths
 The wavelengths of emitted light depend
on the element.
Drawback to Bohr
 Bohr’s theory did not
explain or show the
shape or the path
traveled by the electrons.
 His theory could only
explain hydrogen and
not the more complex
atoms
Energy level populations (Science10)
 Electrons found per energy level of the atom.
 The first energy level holds 2 electrons
 The second energy level holds 8 electrons
 The third energy level holds 18 electrons
Examples for group 1
 Li
 Na
K
2.1
2.8.1
2.8.8.1
The Quantum Mechanical Model
 Energy is quantized. It comes in chunks.
 A quanta is the amount of energy needed to move
from one energy level to another.
 Since the energy of an atom is never “in between” there
must be a quantum leap in energy.
 Schrödinger derived an equation that described the
energy and position of the electrons in an atom – an
ORBITAL
Orbits (Bohr) vs Orbitals (Quantum
Mechanics)
Bohr said electrons travel in an orbit – can
predict exact location of electron at any point in
time.
Schrodinger used mathematics (calculus) to
find the region in space where an electron will
be found 90% of the time - this region is called
an orbital. There are 4 main types of orbitals –
s, p, d, and f.
Modern View of the Atom
The modern view of the atom suggests
that the atom is more like a cloud.
Atomic orbitals around the nucleus
define the places where electrons are most
likely to be found.
23
s orbitals
 1 s orbital for
every energy level
1s
2s
3s
 Spherical shaped
 Each s orbital can hold 2 electrons
 Called the 1s, 2s, 3s, etc.. orbitals
p orbitals
 Start at the second energy level
 3 different directions
 3 different shapes
 Each orbital can hold 2 electrons
The d sublevel contains 5 d orbitals
 The d sublevel starts in the 3rd energy level
 5 different shapes (orbitals)
 Each orbital can hold 2 electrons
The f sublevel has 7 f orbitals
 The f sublevel starts in the fourth energy level
 The f sublevel has seven different shapes (orbitals)
 2 electrons per orbital
Electron Configuration
 We use e- configuration as a shorthand to
show how e- are arranged around a nucleus
 Example: Carbon is …
Electron Configurations
 The way electrons are arranged in atoms.
 Aufbau principle- electrons enter the lowest
energy first.
 This causes difficulties because of the overlap of
orbitals of different energies.
 Pauli Exclusion Principle- at most 2 electrons
per orbital - different spins
 Hund’s Rule- When electrons occupy orbitals of
equal energy they don’t pair up until they have to .
Summary
# of
Max
Starts at
Sublevel shapes number of energy
(Orbitals)
elevel
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
Electron Arrangement
1st Rule: The Aufbau Principle
 e- fill orbitals of the lowest energy first
 We can use the periodic table to help us!
The Diagonal Rule
Example #1
1s
2s
 Oxygen
2p
Example #2
1s
2s
3s
 Magnesium
2p
Example #3
1s
2s
3s
4s
 Iron
2p
3d
3p
Practice
 Boron
 Argon
 Calcium
 Iodine
 Sodium
 Zinc
 Lead
Abbreviations
 We can abbreviate electron configurations
using the Noble Gases
 Ex: Sulfur
 1s2 2s2 2p6 3s2 3p4
 [Ne] 3s2 3p4
 Ex: Lead
 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10
6p2
 [Xe] 6s2 4f14 5d10 6p2
2nd Rule: Pauli Exclusion Principle
 Each orbital orientation can hold up to 2 e e- must have opposite spins (up/clockwise or
down/counter clockwise)
 Therefore:
 s has up to 2 e- (1 orientation)
 p has up to 6 e- (3 orientations)
 d has up to 10 e- (5 orientations)
 f has up to 14 e- (7 orientations)
 We can use the 2nd rule to draw
Orbital Diagrams
Example #1
 Oxygen: 1s2 2s2 2p4
Example #2
 Magnesium: 1s2 2s2 2p6 3s2
Example #3
Iron
3rd Rule: Hund’s Rule
 e- will not pair up until each orbital
orientation has 1 e- in it
 The first e- in a pair will spin up, the second
will spin down
 Example: Oxygen is 1s2 2s2 2p4
Orbital Notation
 Orbital Notation shows us visually the
arrangement and spin of electrons
 Example: Carbon is 1s2 2s2 2p2
Energy Level Diagrams
 Energy Level
Diagrams give us the
same information as
orbital diagrams, plus
they show us the
different energy levels
of each orbital
 Example:
Carbon is 1s2 2s2 2p2
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
3d
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p  Phosphorous, 15 e- to
place
2p  The first to electrons go
into the 1s orbital
 Notice the opposite
spins
 only 13 more
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p  The next electrons go
into the 2s orbital
2p
 only 11 more
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons go
into the 2p orbital
2p
• only 5 more
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons go
into the 3s orbital
2p
• only 3 more
5f
4f
Increasing energy
7s
6s
5s
4s
7p
6p
6d
5d
5p
4d
4p
3p •
3s
2s
1s
2p •
•
•
5f
4f
3d
The last three electrons
go into the 3p orbitals.
They each go into
separate shapes
3 unpaired electrons
1s22s22p63s23p3
Orbitals fill in order
 Lowest energy to higher energy.
 Adding electrons can change the energy of the orbital.
 Half filled orbitals have a lower energy.
 Makes them more stable.
 Changes the filling order
Write these electron
configurations
 Titanium - 22 electrons
 1s22s22p63s23p64s23d2
 Vanadium - 23 electrons
 1s22s22p63s23p64s23d3
 Chromium - 24 electrons
 1s22s22p63s23p64s23d4
Electronic Structure - Questions
 Copy and complete the following table:
Atomic
no.
Mass
no.
No. of
No. of
No. of
protons neutrons electrons
Mg
12
Al3+
27
S2Sc3+
Ni2+
21
1s2 2s2 2p6 3s2
10
16
16
45
30
Electronic
structure
26