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Atomic Structure
Chapter 2
Knowing the Atom
Composed of subatomic particles
Protons (p+)
Neutrons (n0)
Electrons (e-)
Consists of 2 regions
Nucleus
Electron shells
History of Chemistry
The atomism concept is thousands of years
old and was originally founded for
philosophical reasoning
The earliest references to atomism date back
to sixth century BC India
Later emerged with Leucippus and
Democritus in Ancient Greece
History Cont.
John Dalton developed the atomic theory
in 1803 explaining why chemical elements
react in simple proportions by mass
He proposed that each element had its
own set of atoms that bond to form
compounds
Did not prove atoms existed, but he did
show his atomic theory was consistent
History Cont.
Avogadro made the distinction between
atoms and molecules (O vs. O2) in 1811
Was not clarified and accepted until 1858 due
to the work of Cannizzaro
Dalton’s atomic theory was modified
throughout the twentieth century and his
original symbols for the elements replaced by
our modern symbols developed by Berzelius,
a Swedish chemist
Relative mass and charges
of subatomic particles
Protons have positive charge
Neutrons have no charge
These two have almost the same mass
(relative mass= 1)
Electrons have a negative charge and very
small relative mass (5E-4)
The opposite electrical charges hold the atom
together
Relative mass and relative
charges of subatomic
particles
Neutron help stabilize the nucleus by separating
the protons and reducing electrostatic repulsion
If too many or too few neutrons are present the
nucleus will undergo radioactive decay
Electrons control the chemical properties of the
elements
Each element has a unique set of chemical properties
History of Chemistry
Electron discovered in 1897 by J.J. Thompson,
a British physicist
He did a series of experiments with cathode
ray tubes and cathode rays to discover them
He found electron mass to be 1000 times less
than that of a hydrogen ion
He also demonstrated that hydrogen has only
a single electron
History of Chemistry
Thompson also developed a model of the
atom known as the “Plum Pudding Model”
This model was later changed to the atomic
model we rely on now by Ernest Rutherford
Mass Number, Atomic
Number and Isotopes
Mass Number The total number of
protons and neutrons (Symbol A)
Atomic Number The number of
protons in the nucleus (Symbol Z)
Isotope Atoms of the same element
that have different mass numbers due
to a different amount of neutrons
AZX Notation
Notation practice
Write the following elements in AZX
notation
Sodium
Hydrogen
Aluminum
Oxygen
Ions
An element forms an ion when one or more
electrons are added or removed from the
atom
A positive ion is formed with electrons are
removed (cation)
A negative ion is formed when electrons are
added (anion)
To determine the number of electrons in an
atom, add or subtract from the number of
protons
Isotopes
Many elements exits as a mixture of isotope
Each has a unique mass and abundance
Used to calculate average atomic mass
Isotopes have identical chemical properties,
but slightly different physical properties
Ex. Uranium-235 and uranium-238 (forms
99.3% natural uranium)
History of Chemistry
French chemist Henri Becquerel
discovered uranium salts (released
radiation which passed through the
wrapping paper around a photographic
plate)
Further experimented by Pierre and
Marie Curie, who named it radioactivity
The Curies also discovered the elements
radium and polonium
Radioactivity
Many elements contain unstable
nuclides which break up spontaneously
with the emissions of ionizing radiation
The unstable nuclides are described as
being radioactive and are called
radioisotopes
Radioisotopes
Radiation
Relative
Charge
Relative
Mass
Nature
Penetration
Deflection
by electric
field
Alpha
particles
+2
4
2 protons
and 2
neutrons
(He2+ ion)
Stopped by
a few
sheets of
paper
Low
Beta
particles
-1
1/1837
Electron
Stopped by
a few mm
of plastic or
Al
High
Gamma
rays
0
0
Electromag Stopped by
netic
a few cm of
radiation of
lead
very high
frequency
None
Radioactive Decay
The rate which nuclei undergo radioactive
decay varies between elements
Radioactive decay is an exponential process
Rates are compared using half-life (the time
taken for half of the radioactive nuclei to
undergo decay)
During alpha and beta decay a more stable
isotope forms
Radioactive Decay
Each radioisotope has a unique half-life
Unaffected by temperature or pressure
Ex. Iodine-131 has a half life of 8 days
Due to half-life, different radioisotopes are used to
diagnose and treat diseases
Ex. Cobalt-60, a powerful gamma emitter, used to
treat different types of cancer (Radiotherapy)
The Mass Spectrometer
Allows chemists to accurately determine the
relative atomic masses of atoms
Can also be used to determine relative molecular
masses and structures of compounds
Look at Interactive syllabus
Determining Relative
Atomic Mass
Relative atomic mass=(% isotope A * mass isotope A)+ (%
isotope B *mass isotope B)…
Ex. Chlorine- 35 and chlorine -37 have percent
abundances of 75% and 25%
Rubidium exists as a mixture of two isotopes 85Rb and
87Rb. The percentage abundances are 72.1% and 27.9%.
The relative atomic mass of gallium is 69.7. gallium is
composed of two isotopes: gallium-69 and gallium-71.
Calculate the percentage abundance of gallium-69.
Light as Waves and
Particles
Light is emitted in the
form of electromagnetic
waves
These consist of an
oscillating electric wave
and magnetic wave which
travel in a sinusoidal
pattern
The two waves are
arrange perpendicularly
Light Wave
Terminology
Wavelength (λ) The distance between two
neighboring crests or troughs of a wave
Frequency (ν) The number of waves which pass a
point in one second. Units are hertz (Hz); 1/sec.
Speed (c) The distance travelled by a wave in one
second. Units are m s-1.
Wave equation c=λν where c= 3.0E8 m s-1
Terminology
Photons “packets” of light energy. The term used
when light is described by a particle model
Planck’s constant (h) A value of 6.63E-34 J s. Used
in the equation: E=hν.
Continuous spectrum The spectrum of colours
formed from white light which is composed of
visible light of a certain range of wavelengths
History of Chemistry
Issac Newton was the first western scientist to show that
sunlight is composed of many colours. He used a prism to split
the light into a spectrum then re-formed it into white light (Late
1600s- early 1700s)
A Muslim scientist, Ibn al-Haytham, performed similar light
dispersion experiments during the Golden Age (965-1040)
Electron Excitation
When gaseous atoms are excited, they emit light of
certain wavelengths
Excitation is when an electron is raised to a higher level,
then light is emitted when it returns to the unexcited
state
Could be a thermal or electrical process
Thermal a flame is formed
Ex. Fireworks exploding
Electrical a high voltage is passed across a gaseous sample
of an element at low pressure
Ex. Neon signs
Emission Spectra
When light from atoms with excited
electrons are passed through a prism,
emission spectra are formed
These spectra are called line spectra
Each chemical element has its own
unique spectrum which could be used
to identify that element
Line Spectra
Line spectra differ from continuous
spectra in two ways:
1. An emission (line) spectrum is made up
of separate lines (discontinuous)
2. the lines converge, becoming
progressively closer as the frequency or
energy of the emission lines increases
History of Chemistry
Johann Jakob Balmer (1825-1898), Swiss
mathematics teacher who studied visible lines of
the hydrogen emissions spectrum
Developed a mathematical formula to calculate
wavelengths
His formula allowed for discovery of additional lines
on the hydrogen spectrum and of white stars
Neils Bohr proposed a model in 1913 to explain
why Balmer’s formula gave correct wavelengths
Energy Levels and
Spectra
Bohr suggested a theory to account for the
complete emission spectrum of helium
He thought electrons moved in orbits where they
had fixed amounts of potential energy
The further an electron from the nucleus, the
greater the energy
Bohr’s theory An electron moving in an
orbit does not emit energy
Energy Levels
In order to move into an orbit further from
the nucleus:
The electron must ABSORB electric or thermal
energy
Do work against the attraction of the positive
nucleus
The electron is then in the excited state
Electron Excitation and
Spectra
Electrons dropping back from higher energy
orbits for emission spectra
The energy of the light is equal to the
difference between the two energy levels:
ΔE=hf
Energy Levels
Bohr labeled his energy levels with n and a number
n=1 is the lowest energy level (nearest the nucleus)
Considered “Ground state”
The most stable state for a hydrogen atom
The next highest energy state is n=2, etc.
If an electron receives enough energy, it is completely
removed from the atom; forming an ion
The amount of energy required to remove the electron is the
ionization energy (transition from n=1 to n=∞)
Hydrogen Atom
H Absorption and
Emission
Hydrogen
Photon Energy
Calculate the energy of photons that five rise
to the red line of the Balmer series. The
wavelength is 656.3 nm and the speed of light
in a vacuum is 3.0 x 108 m s-1.
f= c/λ
E=hf (h=6.63 x10-34 J s)
History of Chemistry
Neils Bohr (1885- 1962) a Danish physicist awarded
Nobel Prize in physics (1922) for atomic structure
contributions
Also part of Manhattan Project; developed first atomic
bomb
Studied with Rutherford and J.J. Thompson
He determined that electrons can only have certain
values or quantities of energy
His theory could only account for the spectra of
hydrogen, thus his theory was flawed and later replaced
by Shrödinger’s quantum mechanical model
Electron Arrangement
Electrons are arranged in electron shells
Hydrogen has only a single electron and it is located in
the shell nearest the nucleus
The first shell (first energy level) can hold a maximum
of 2 electrons
The second shell (second energy level) can hold a
maximum of 8 electrons
The third shell (third energy level) can hold a maximum
of 18 electrons
Electron Arrangement
Indicates the number of electrons without drawing
the shells
Ex. Lithium has 3 electrons and an arrangement of 2,1
or 2.1
Ex. Sodium has an arrangement of 2,8,1 or 2.8.1
Ionization Energy (HL)
This force is dependent on two main
factors and is fine tuned by a third factor.
Size of the atom
The nuclear charge
Shielding effectElectrons between the
outer electron and the nucleus (fine
tuning)
Atomic Radius
Distance of the outer electrons from
nucleus
As distance between electron and
nucleus increases, the attraction
decreases
Causes ionisation energy to decrease
Nuclear Charge
As nuclear charge becomes more
positive, the attraction to electrons
becomes stronger
Increases ionisation energy
Shielding Effect
Outer valence electrons repelled by inner electrons
and kept from the pull of the nucleus
Mostly effective if electrons are close to nucleus
Electrons in the first shell have the highest shielding
effect
The more shielded from the nucleus, the lower the
ionization energy
HL Ionizing Energies
Comparing the I.E. of
H and He
If we compare the energy required to remove the
outermost electron of hydrogen (element number 1)
and helium element number 2) we see that
approximately twice the energy is needed for
helium. As the nuclear charge of helium is twice that
of hydrogen we can conclude that the electron is
roughly the same distance from the nucleus in both
cases.
The first ionisation
energy
The first ionisation energy is defined as the
energy required to remove one mole of
electrons from one mole of gaseous atoms to
provide one mole of gaseous single charged
ions.
Na(g) --> Na+(g) + 1e
Subsequent ionisation energies are defined in a
similar way only by removing electrons from
already charged ions.
The second ionisation
energy
Na+(g) --> Na2+(g) + 1e
Successive electrons can be stripped from an
atom until there is only the nucleus left. If the
energy required to achieve this for each
ionisation is plotted on a graph (with a log scale)
against the ionisation number, the 'jumps' in the
required energy clearly show the main and sub
energy levels.
Ionization energy trend
Drops in I.E.
There are dips in ionization across periods 2 and
3
First decrease is because electron shielding is
changed due to loss of sub-level electron
(Between Be and B)
More E required for Be to lose a lower-level electron
Second decrease between N and O
Easier to remove electron from O because it is part
of a pair; N has a single electron in that orbital
Hund’s Rule
Every orbital in a sub-shell is single occupies
with one electron before any one orbital is
doubly occupied
Reason for the drop between N and O
Practice using Hund’s rule with:
Nitrogen
N+
Oxygen
O+
O2-
Relating I.E. and
Configuration
I.E. are always endothermic: energy absorbed
and work done to remove electron(s)
Second I.E. always larger than the first
Further successive ionization energies increase
Successive Ionisation E
of K
Interpreting the graph
Electron arrangement is 2.8.8.1
Which part of the graph shows the first electron(s)
removed?
After which removed electrons are the largest
increases in I.E.?
What do these large jumps correspond to?
The y-axis is measured in log10 to reduce the values of
the “jumps”
Construct a Graph
Sketch a graph of the successive ionization
energies of silicon against electrons removed
Start by determining the electron arrangement
for Si
Then, start from the outer electrons and work
your way to the inner
Working in Reverse
The first 8 experimentally determined I.E. for a
chemical element are as follows (kJ mol-1)
580 1800 2750 11580 14850 18400 23300 27500
Deduce the following:
Number of e- in outer shell of an atom of the element
The group in the periodic table which the element
belongs
Outer detailed configuration of the atom
Orbitals and Energy
Levels
Electrons are found in shells around nucleus, numbered 1, 2, 3 etc
Each shell has sub-shells, labeled s, p, d, or f
Number of sub-shells are equal to the number of the shell
Shell (energy level)
number
Number of sub-shells in
the shell
Sub-shells (sub-levels)
1
1
1s
2
2
2s and 2p
3
3
3s, 3p, 3d
4
4
4s, 4p, 4d, 4f
Sub- Shells
Each sub-shell contains a particular number of
orbitals, where the electrons are located
Number of orbitals depends on the type of sub-shell
Each orbital can hold a maximum of 2 electrons
Type of sub-shell
(sub-level)
Number of orbitals
Maximum number of
electrons in the subshell (sub-level)
s
1
2
p
3
6
d
5
10
f
7
14
Energy shells split into
sub-shells
Diagram
• Note that the 4s sub-shell has lower energy than 3d, so the
electrons will fill it before the 3d
Electron Configurations
Write a complete electron
configuration for the first 15 elements
on the periodic table (H P)
Shapes of Orbitals
The four types of orbitals all have different
shapes
s orbitals are spheres and differ in size and
energy
p have a dumb-bell shape and have different
orientations (px, py, pz) and do not differ in
energy (degenerate)
History of Chemistry
Erwin Schrödinger developed Schrödinger wave
equation to give energy levels of H atom and the
shapes and energies of orbitals; published in 1926
Werner Heisenberg proposed that it was impossible to
measure both location and momentum of an electron;
Heisenberg’s uncertainty principle published in the
mid-1900’s
A scientist can measure the position with some accuracy,
but the momentum has a large range and vice versa
Filling Atomic Orbitals
Pauli exclusion principle each orbital can hold
a maximum of two electrons
Aufbau principle electrons occupy an empty
atomic orbital with the lowest energy
Hund’s rule within a sub-shell electrons
experience repulsion, thus enter two different
orbitals of the same energy
Electrons can spin clockwise  or anticlockwise 
Single electrons in sub shells have the same spin
Electron Configurations
of Atoms
Hydrogen:
1s
1s1

Lithium: 1s
1s22s1

Helium:
1s
Carbon: 1s22s22p2
1s2
1s
2s
2s

2p





Electron Configurations
to Know! (HL)
1-38
49-54
Electron Configurations
for Ions
Electron configurations of ions are determined
based on the same principles as neutral atoms
The added or lost electrons are simply included
or excluded from the arrangement
Ex: Mg 1s22s22p63s2
Mg2+  1s22s22p6
Octet Rule
Atoms gain or lose electrons to achieve
a complete outer electron orbital with 8
electrons
The noble gases have a complete
orbital, thus are unreactive