chapter 3: electrons in atoms. learning outcomes: energy levels and shapes of orbitals electronic...
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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..
3rd shell n=3
2nd shell n=2
1st shell n=1 Max:2
n2
Each shell has a maximum number of electrons that it can hold. Electrons fill the shells nearest the nucleus first.
3rd shell holdsa maximum of18 electrons
2nd shell holdsa maximum of
8 electrons
1st shell holdsa maximum of
2 electrons
Time for a break and practice
Please make check-up 1 on page
Simplified electron configuration
2, 5
2, 8
2, 8, 7
2, 8
7N2nd shell holdsa maximum of
8 electrons
1st shell holdsa maximum of
2 electrons
17Cl
11Na+
8O2-
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:
Variation of first ionisation energy with atomic number for the first twenty elements
0
500
1000
1500
2000
2500
0 5 10 15 20
atomic number
firs
t io
nis
atio
n e
ner
gy
(kJ
per
mo
le)
Worked example
The model of the atom
A model is what fits logic, experimental observations and mathematical calculations
17Cl 2, 8, 7 3
2nd shell, witha maximum of
8 electrons
Symbol Simple electronic configuration Number of shells (last number is Group)
(=period)
1st shell, witha maximum of
2 electrons
3nd shell, witha maximum of18 electrons
6C
10Ne
11Na
Symbol Simple electronic configuration Number of shells (last number is Group)
(=period)
2, 4 2
2, 8 2
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 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
n = 1
n = 2
Quantum shell 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 2s2p
26
n = 3 3 3s3p3d
2610
Subshells and their shapes
Atomic orbital is a space around the nucleus holding 1 or 2 electrons
n = 1 1s
Where in the atom are the electrons?
energy
2He 2
2
Simple electronic configuration Complicated electronic configuration
1s2
Principle quantum shell
Sub-shellNumber of electrons
n = 1
n = 2
1s (e<2)
2s (e<2)
2p (e<6)
Where in the atom are the electrons?
energy
8O 2, 6
2
24
Simple electronic configuration Complicated electronic configuration
1s22s22p4
n = 1
n = 2
1s (e<2)
2s (e<2)
2p (e<6)
Where in the atom are the electrons?
energy
11Na 2, 8, 1
226
Simple electronic configuration Complicated electronic configuration
1
n = 33d (e<10)
3p (e<6)3s (e<2)
n = 1
n = 2
1s (e<2)
2s (e<2)
2p (e<6)
Where in the atom are the electrons?
energyn = 3
3d (e<10)
3p (e<6)
3s (e<2)
n = 4
4s (e<2)
4d (e<10)
4p (e<6)
4f
Subshells and atomic orbitals
From simple to complicated electron configuration to noble gas electronic configuration notation
2, 6 1s22s22p4 [He] 2s2sp4
2, 8, 7 1s22s22p63s23p5 [Ne] 3s23p5
8O
17Cl
19K
35Br
Element: Simple: Complicated: Noble gas:
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 group block
1s22s22p5
1s22s22p63s23p64s23d104p2
1s22s1
1s22s22p63s2
1s22s22p63s23p64s23d5
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
2, 6 becomes 1s22s22p6
2, 8, 7 becomes 1s22s22p63s23p5
2, 8, 8 becomes 1s22s22p63s23p6
7N
8O2-
17Cl
19K+
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:
Variation of first ionisation energy with atomic number for the first twenty elements
0
500
1000
1500
2000
2500
0 5 10 15 20
atomic number
firs
t io
nis
atio
n e
ner
gy
(kJ
per
mo
le)
Ionisation: Trend down a group
General trend decrease further away from the nucleus increased shielding despite increased nuclear charge
Li = 519 kJ/molNa = 494 kJ/molK = 418 kJ/molRb = 403 kJ/mol
Worked example