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SCH4U CHEMISTRY
BOHR MODEL (approx 1913)
The Bohr model of the atom was based on the line spectra of the Hydrogen
atom. The model also incorporated the concept developed by Einstein
regarding the particle behaviour of light during emission or absorption (photon
or quanta of energy).
Postulates:
1. Energy Levels
 an electron can only have specific energy values in an atom
 the path followed by the electrons is a circular orbit (spherical in 3-D)
 these energies are called energy levels given the name Principal Quantum
Number, n
 the orbit closest to the nucleus is given n=1, with the lowest
energy
 as one moves outward the energies get larger and “n”
increases (n=1,2,3,4,...)
 an electron can only circle in one of the allowed orbits WITHOUT a loss of
energy
2.
Transition Between Energy Levels
 an electron in an atom can only change energy by going from one energy
level to another level, i.e. NOT in-between
 light is emitted when an electron falls from a higher energy level to a
lower energy level
 an electron generally remains in its ground state – the lowest energy
level possible (n=1 for hydrogen)
 when an electron is given energy it can be bumped to a higher energy
level called the excited state
 as the electron drops from the excited state down to lower energy
levels it emits light
 Bohr was able to show that his model was able to match the Balmer
series (Visible light, drop to n=2) and predict the Paschen series ( IR light
, drop to n=3) and the Lyman series (UV light, drop to n=1)
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Although the Bohr model could explain the behaviour of the hydrogen atom, it
could not fully explain the behaviour of other atoms
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QUANTUM NUMBERS
During and subsequent to Bohr’s work, others continued to probe the behaviour
of light emitted or absorbed by the atom. As a result, additional components
were added to Bohr’s model to improve its ability to explain the behaviours
observed. This led to the introduction of quantum numbers, in addition to
Bohr’s Principle Quantum Number.
Principal Quantum Number, n




Goes back to Bohr model which labeled the shells
Today called Principal Quantum Number, i.e. n=1,2,3……etc.
n=1 is closest to the nucleus with the lowest energy
Relates primarily to the main energy of an electron
e.g. Fig. 1 “energy staircase” where energy levels are like unequal steps
Secondary Quantum Number, l
 Michelson found that the spectrum of H was composed of more than one
line (experimental observation using spectrometry indicated that main
lines of hydrogen spectra composed of more than one line – line splitting
(Michelson, 1891)
 Sommerfeld (1915) used elliptical orbits to extend the knowledge of the
time by explaining Michelson’s work
 He introduced the concept of there being additional electron subshells (or
sublevels) that formed part of the energy levels and used the concept of
the secondary quantum number to describe this concept
 I.e. each energy “step” was a group of several little steps
 “l” relates to the shape of the electron orbit and the number of values for
l equals the volume of n
 i.e. if n=3, then l=0,1,2
as letters: l = s, p, d, f, g
Magnetic Quantum Number, ml
 it was observed that if a gas discharge tube was placed near a strong
magnet, some single lines split into new lines
 called normal Zeeman effect after Zeeman who 1st observed this (1897)
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 this was explained by Sommerfeld & Debye (1916) who thought that
orbits may exist at varying angles and that the energies may be different
when near strong magnets
 for each value of l, ml can vary from –l to +l (each value represents a
different orientation)
 i.e. if l=1 then ml can be –1,0 or +1
if l=2 then ml can be –2,-1,0,+1, or +2
Summary: The magnetic quantum number ml , relates to the direction of the
orbit of the electrons. The number of values of ml represents the number of
orientations of the orbits that we can have.
Spin Quantum Number, ms
 needed to explain additional spectral line-splitting & different kinds of
magnetism
 ferromagnetism-associated with substances containing Fe, Co & Ni
 paramagnetism-weak attraction to strong magnets (individual atoms vs.
collection of atoms)
 paramagnetism couldn’t be explained until Wolfgang Pauli suggested that
electrons spin on their axis (1925)
 could spin only 2 ways (clockwise vs. counterclockwise) and he used only
2 numbers to describe this:
ms = +1/2 (clockwise) or –1/2 (counterclockwise)
 opposite pairs of electron spins represent a stable arrangement
 when electrons are paired – spin in opposite directions – the magnetic
field is neutralized, while an individual electron spin can be affected by a
magnet
ATOMIC STRUCTURE & THE PERIODIC TABLE
 using the quantum numbers gives an ordered description of the electrons
in a particular atom
 the secondary quantum number, l, is presented as s, p, d, and f
to represent the shape of the orbitals
Value of l
Letter designation
Name designation
0
s
sharp
1
p
principal
2
d
diffuse
3
f
fundamental
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Electron orbitals
 4 quantum numbers apply equally for the electron orbits (paths), as well
as, electron orbitals (clouds) [see Fig. 1 & Table 2 on p. 185]
 1st 2 quantum numbers (n & l) describe electrons that have different
energies under normal circumstances in multi-electron atoms
 the last 2 quantum numbers describe electrons that have different
energies under special conditions (e.g. strong magnetic fields
 we use a number for the main energy level and a letter for the energy
sublevel
 simpler to use 1s than n=1, l=0 and 2p is n=2, l=1…etc
 if we needed the third quantum number ml we would need another
designation, 2px, 2py, 2pz
Value of l
0
1
2
3
Sublevel symbol
s
p
d
f
Number of orbitals
1
3
5
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Creating Energy-Level Diagrams
 these diagrams indicate which orbital energy levels are occupied by
electrons for an atom or ion
 In fig.2 on p. 187, as atoms become larger & the main energy levels come
closer, some sublevels may overlap
 Generally the sublevels for a particular value of n, increase in energy in
the order of s<p<d<f.
There are also some rules that apply to creating an energy-level diagram
 Pauli Exclusion Principle - states that no 2 electrons in an atom have
the same 4 quantum numbers (usually you place 1 electron with  and
the other electron )
 Aufbau Principle – an energy sublevel must be filled before moving into
the next higher energy sublevel
 Hund’s Rule – one electron is placed into each of the suborbitals before
doubling up any pair of electrons
 These are the “rules” that are followed until all the electrons have been
placed in their proper place
(see Fig. 6 p. 188 for order diagram.)
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1s, 2s, 2p,3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 6f, 7d, 7f.
Order of filling sub-shells:
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
7d
6d
5d
4d
3d
7f
6f
5f
4f
7g
6g
5g
The filling of sub-shells can also be shown on an energy level diagram as
follows:
6s
5p 5p
5p
5s
4p 4p
4d 4d
4d
4d
3d
3d 3d
3d
3d
4p
4s
3s
4d
3p 3p
3p
2p 2p
2p
2s
1s
This method allows one to see the different energy levels in diagram form –
energy is on the vertical axis. Each underline (orbital) can take up to 2
electrons – one electron  (clockwise spin) and the other electron 
(counterclockwise spin)
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Sample problem 4:
O = 8 electrons
6s
5p 5p
5s
4s
3s

2s

1s
4p 4p
5p
4d
4d 4d
4d
4d
3d
3d 3d
3d
3d
4d
4d 4d
4d
4d
3d
3d 3d
3d
3d
4p
3p 3p
 
2p 2p
3p

2p
5p 5p
5p
Fe = 26 electrons
6s
5s
4s
4p 4p
4p
3p 3p
3p
2p 2p
2p
3s
2s
1s
Creating Energy-Level Diagrams for Anions
 Are done the same way except add the extra electrons that corresponds
to the ion charge to the total number of electrons before showing the
distribution of electrons
e.g. S has 16 electrons but S-2
has 18 electrons
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Creating Energy-Level Diagrams for Cations
 Draw the energy-level diagram for the neutral atom first, then remove the
number of electrons needed (for the correct ion charge) from the orbitals
with the highest principal quantum number
Homework:
Nelson 12 p. 191 #’s 1-4
Electron Configuration
 Provide the same information in a more concise format
Examples:
 O: 1s2 2s2 2p4
 S-2: 1s2 2s2 2p6 3s2 3p6
 Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d6
Sample Problems p.192 & 193
Shorthand Form of Electron Configurations
 There is an internationally accepted form of shorthand that can be used
where the core electrons are expressed by the preceding noble gas and
just adding the electrons beyond the noble gas
 E.g. Cl: 1s2 2s2 2p6 3s2 3p5 becomes [Ne] 3s2 3p5
Homework:
Nelson 12 p. 194 #’s 6-11
Explaining the Periodic Table
Period
Period
Period
Period
Period
Period
1
2
3
4-5
6-7
# of elements
2
8
18
18
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Electron Distribution:
Groups: 1-2 13-18 3-12
Orbitals: s
p
d
2
2
6
2
6
10
2
6
10
2
6
10
f
14
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 Can use the filling of the subshells to explain the Periodic Table
 Noble gases all have ns2 np6 in the outer shell of electrons
 These types of patterns of configuration are seen in the representative
elements
 Transition elements can be explained as elements filling the d elements
therefore the transition elements are sometimes referred to as the d block
of elements
 Lanthanides and actinides are explained by the filling of the f orbitals
Explaining Ion Charges
 Zn: [Ar] 4s2 3d10 has 12 outer electrons but Zn+2: [Ar] 3d10 and it is
unlikely that Zn would give up 10 electrons to go to 4s sublevel
 Pb: [Xe] 6s2 4f14 5d10 6p2 lead can either lose 2 6p electrons to become
Pb2+ or lose 4 electrons from the 6s and 6p orbitals to form Pb4+
Explaining Magnetism
 To explain magnetism we can draw the electron configuration of a
ferromagnetic element, e.g. Fe [Ar] 4s2 3d6 = 1 pair + 4 unpaired=
unpaired electrons cause the magnetism?
 Ruthenium which is right under iron is only paramagnetic (weakly
magnetic)
 Result occurs probably because the atoms form groups called domains
that cause this type of magnetism
 “Ferromagnetism is based on the properties of a collection of atoms,
rather than just one atom”
Anomalous Electron Configurations
 List on p.196 shows some anomalies in predictions of electron
configurations
 evidence suggests that half-filled and filled subshells are more stable
(lower energy) than unfilled subshells
 i.e. Cr: [Ar] 4s2 3d4 expected but in fact it is [Ar] 4s1 3d5
 appears more important for d subshells and in the case of chromium an s
electron is promoted to the d subshell to create two half-filled subshells
 the justification is that the overall energy state is lower after the
promotion of the electron
Homework: Nelson 12 p. 197 #’s 1-14
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PROBABILITY MODEL
(WAVE MECHANICAL MODEL or
QUANTUM MECHANICAL MODEL)
DeBroglie (1924)
 Proposed that moving electrons could be considered to behave like waves
 experimentation showed that electrons can exhibit diffraction patterns like
light – a key part of the wave model of light
 derived expression relating wave properties to particle properties
vm/h = ‫ג‬
where: fo htgnelevaw = ‫ג‬
particle wave
h = Planck’s Constant
m = mass of particle
v = speed of particle
for a baseball of m= 0.145 g with v=27 m/s
01 = ‫ג‬-34 m
for an electron of m = with v=3.00x108 m/s
01 x 001 = ‫ג‬-12 m (100 pm)
visible light has a wavelength range of 400 to 750 nm
Schrodinger’s Theory (1926)
 extended wave properties of electrons to their behaviour in atoms and
molecules
 consider Bohr’s Model – electron orbits nucleus like the earth around the
sun in a continuous path
 in quantum mechanics – electron does NOT have a precise orbit in an
atom
 using principle of standing waves (wave theory) and the quantum numbers,
Schrodinger was able to devise an expression, called the wave equation, that
could identify the space an electron will occupy
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Heisenberg’s Uncertainty Principle (1927)
 Heisenberg claimed that it was impossible to know at the same time
(absolutely), both the position and momentum of an electron in an atom
(Δx)(Δpx) ≥ h/4
where: Δx = uncertainty in position
Δpx = uncertainty in momentum
 another way of saying this is:
“if you know where a particle is, you cannot know where it is going”
 because of the interaction of photons of light with electrons, once one
identifies where the electron is, then the light will excite the electron and
send to a place unknown
 nonetheless, one can make a statistical statement about where an
electron is in an atom – one can indicate the probability of finding an
electron at a certain point in an atom
 to do this we can employ the wave equation
 suggests that the atom does not have a definite boundary – unlike the
Bohr model
probability
of finding
particle at
position r
for 1s
orbital
50
100
150
r (pm) – distance from nucleus
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