Download PPT chapter 3

Chapter 3:
Electrons in atoms
Learning outcomes:
 Energy
levels and shapes of orbitals
 Electronic
configurations
 Ionisation
energy, trends across a period
The quantum mechanical model
How are electrons arranged?
Electrons are not evenly spread but exist in layers called
shells. The arrangement of electrons in the shells is called
the electron structure or electronic configuration.
1st shell
n=1
Max:2
n2
2nd shell
n=2
3rd shell
n=3
Each shell has a maximum number of electrons that it can
hold. Electrons fill the shells nearest the nucleus first.
1st shell holds
a maximum of
2 electrons
2nd shell holds
a maximum of
8 electrons
3rd shell holds
a maximum of
18 electrons
Time for a break and practice
Please make
check-up 1 on
page
Simplified electron configuration
7N
2, 5
17Cl
2, 8, 7
+
11Na
2, 8
8O
2-
2, 8
2nd shell holds
a maximum of
8 electrons
1st shell holds
a maximum of
2 electrons
For “ion” the number of
proton is NOT the same
as electrons
Electronic configuration


This model assumes the electrons
have the same location and energy
Until it was discovered that electrons
have different locations and energy
Li
Li+
+ e-
(1st ionisation energy: E1)
Li+
Li2+ + e-
(2nd ionisation energy: E2)
Li2+
Li3+ + e-
(3rd ionisation energy: E3)
E1 ≠ E2 ≠ E3
Ionisation energy
Li
Li+
+ e-
(1st ionisation energy: E1)
Li+
Li2+ + e-
(2nd ionisation energy: E2)
Li2+
Li3+ + e-
(3rd ionisation energy: E3)
Li 2,1
Which electron will be easiest
to remove?
E1 < E2 < E3
ΔHi1 < ΔHi2 < ΔHi3
Table 3.2 in the book on p. 35




For every element, the successive ionisation energy increases;
for every next electron it is more difficult to remove
We can in theory continue removing electrons until only the
nucleus is left
We call this sequence the “successive ionisation energy”
Sometimes we find a big gap/jump in ionisation energy
Example: sodium

The first ionisation energy is quite low,
it is likely quite far from the nucleus

The 2nd to the 9th ionisation energy are
in a gradual successive increase
indicating these electrons are in the
same shell

The 10th and 11th electrons have high
ionisation energies compared to the
rest, they must be near the nucleus.

The jump between the 9th and 10th
suggests a change in shell
Factors affecting the first
ionization energy
Nuclear charge (number of protons)
the bigger nuclear charge, the higher 1st ionization energy.
 Atomic radius (distance effect)
the bigger atomic radius, the lower 1st ionization energy.
 Shielding effect (number of shells)
the bigger Shielding effect, the lower 1st ionization energy

The first ionization energies of the first 20
elements in the periodic table is shown below:
first ionisation energy (kJ
per mole)
Variation of first ionisation energy with atomic
number for the first twenty elements
2500
2000
1500
1000
500
0
0
5
10
atomic number
15
20
Worked example
The model of the atom
A model is what fits logic, experimental observations and
mathematical calculations
Symbol
Simple electronic configuration
(last number is Group)
17Cl
1st shell, with
a maximum of
2 electrons
2, 8, 7
2nd shell, with
a maximum of
8 electrons
Number of shells
(=period)
3
3nd shell, with
a maximum of
18 electrons
Symbol
Simple electronic configuration
(last number is Group)
Number of shells
(=period)
6C
2, 4
2
10Ne
2, 8
2
11Na
2, 8, 1
3
19K
4
Where in the atom is the electron?
According to quantum mechanics it is
most likely to find the electron for the
of the H-atom at 0.0000000000529
meter (52.9 pm) from the nucleus
Shells
n=2


n=1

Quantum shell
Principal quantum shells (n=1, n=2
etc.)
Remember for each the max number
of electrons is 2n2 (so for n=2, max 8
electrons)
We know from experiments and
calculations these 8 electrons have
different energies…. so we need a
new model of the atom where we can
distinguish between electron energy
Subshells
The quantum mechanical model
Simplified model
Realistic model
Principal
quantum
shell
Number of
Sub-shells
Name of the
Sub-shell
Max. number
of electrons
n=1
1
1s
2
n=2
2
2s
2p
2
6
n=3
3
3s
3p
3d
2
6
10
Subshells and their shapes
Atomic orbital is a
space around the
nucleus holding 1 or 2
electrons
Where in the atom are the
electrons?
Simple electronic configuration
Complicated electronic configuration
2He
2
1s2
Principle
Sub-shell
quantum shell
energy
n=1
2
1s
Number of
electrons
Where in the atom are the
Simple electronic configurationelectrons?
Complicated electronic configuration
8O
1s22s22p4
2, 6
energy
n=2
n=1
4
2
2p
(e<6)
2s
(e<2)
2
1s
(e<2)
Where in the atom are the
Simple electronic configurationelectrons?
Complicated electronic configuration
11Na
2, 8, 1
n=3
energy
n=2
n=1
1
6
2
2
3d
(e<10)
3p
3s
(e<6)
(e<2)
2p
(e<6)
2s
(e<2)
1s
(e<2)
Where in the atom are the
electrons?
4f
n=4
energy
n=3
n=2
n=1
4d
(e<10)
4p
(e<6)
3d
(e<10)
4s
(e<2)
3p
(e<6)
3s
(e<2)
2p
(e<6)
2s
(e<2)
1s
(e<2)
Subshells and atomic orbitals
From simple to complicated electron configuration to noble
gas electronic configuration notation
Element:
8O
17Cl
19K
35Br
Simple:
2, 6
2, 8, 7
Complicated:
Noble gas:
1s22s22p4
[He] 2s2sp4
1s22s22p63s23p5 [Ne] 3s23p5
Note the following:





Potassium: 1s22s22p63s23p64s1
The 3d subshell: 3d<4p
Chromium and copper are exceptions:
Cr: [Ar] 4s13d5 rather than [Ar] 4s23d4 and
Cu: [Ar] 4s13d10 rather than [Ar] 4s23d9
The blocks of the periodic table



Elements in Group 1 and Group 2 are in the s-block and have their
outer electrons in an s subshell.
Elements in Group 3 to 18 have outer electrons in a p subshell.
Elements that add electrons to the d subshells are called the d-block
elements.
Use the electronic configuration to find the group
…s2
is in group:
…p1
is in group:
…p3
is in group:
…p6
is in group:
…d3
is in group:
…d7
is in group:
In which period, group and block of
the following electron configuration?
period
1s22s22p5
1s22s22p63s23p64s23d104p2
1s22s1
1s22s22p63s2
1s22s22p63s23p64s23d5
group
block
RULES FOR FILLING ENERGY LEVELS
Aufbau Principle “Electrons enter the lowest
energy orbital first”
 Pauli’s Exclusion “Sub-Orbitals can hold a max. of
2 electrons provided they have opposite spin”
 Hund’s Rule “Orbitals of the same energy remain
singly occupied before pairing up.

Examples
N = 1s22s22p3
O = 1s22s22p4
From simple to complicated electron
configuration
2, 5 becomes 1s22s22p3
7N
8O
2-
17Cl
+
19K
2, 6 becomes 1s22s22p6
2, 8, 7 becomes 1s22s22p63s23p5
2, 8, 8 becomes 1s22s22p63s23p6
Ionisation: trend across a period




General increase across period
Rapid decrease between last
element of a period and 1st of a new
period
Be and B because 2s and 2p
N and O
The first ionization energies of the first 20
elements in the periodic table is shown below:
first ionisation energy (kJ
per mole)
Variation of first ionisation energy with atomic
number for the first twenty elements
2500
2000
1500
1000
500
0
0
5
10
atomic number
15
20
Ionisation: Trend down a group
 General
trend decrease
 further away from the nucleus
 increased shielding
 despite increased nuclear charge
Li = 519 kJ/mol
Na = 494 kJ/mol
K = 418 kJ/mol
Rb = 403 kJ/mol
Worked example