A topic which is part of atomic structure

1. The Period Table and
Periodic Trends in Atomic
Properties

2. The Periodic Table
Main features
• The layout of the periodic table gives valuable information
about the elements: chemical and physical etc
• The vertical columns are known as groups. Elements in the
same group have the same configuration in the valence
shell (outer most shell): Group I [X]ns1, Group II [X]ns2
• The horizontal rows are called periods. There are seven
periods each representing quantum numbers n = 1 → n= 7
• Each period is filled more or less sequentially. The period an
element is in tells us the highest shell occupied
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3. Effective nuclear charge
• Properties of elements are affected by the
amount of positive charge experienced by the
outer electrons
• The positive charge is always less than the
full nuclear charge (except for hydrogen)
• The negative charge of the inner electrons
partially offset (neutralise) the positive charge
of the nucleus
• The effective nuclear charge is determined by
the difference between the charge of the
nucleus and the charge on the core
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Example
• Li 1s22s1
• (1s2) core lies below the valence electron
• Are tightly packed around the nucleus, for
most time lie between and the outer
electron
• The core has a charge of 2− and nucleus
charge 3+
• Effective charge ‘felt’ by the outer
electron is 1+

4. Atomic Size
• Two important
factors to consider:
• increasing
principal number
• effective nuclear
charge
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Down the group
 Effective nuclear charge remains nearly constant
whilst the principal quantum number of the
valence shell increases
E.g. Group IA elements:
Li 2s1, Na 3s1, K 4s1, …
 For each of these elements the effective nuclear
charge is  1+
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5. Atomic Size
• n increases as one descends the group
• The orbital containing the valence electrons become
bigger hence atom larger
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• E.g. Li to F nuclear charge
changes 3+ to 9+ whilst core (1s2)
remains the same
• Consequently. the outer electrons
feel a greater positive charge that
causes the electrons to drawn
inwards – hence size of atoms
reduces
Across period
• Left to right the nuclear charge
increases
• Outer shells become more populated
but the inner core remains the same
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6. Atomic radius vs. atomic number
Ca
K
H He
Li
B
Be
C N O F
Ne
Na
Mg
Al Si P S Cl
Ar
0
50
100
150
200
250
0 2 4 6 8 10 12 14 16 18 20
Element
Atomic
Radius
(pm)
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7. Ionisation energy
• The energy required to remove an electron
from an isolated, gaseous atom or ion in
its ground state.
X (g) → X+(g) + e−
• Measures amount of work needed to pull
out an electron – gives an idea of how
tightly bound the electron is to the nucleus.
• Requires energy input – the tightly bound,
the more energy required to remove the
electron.
• Successive ionisation energies become
increasingly larger
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Decreasing
Increasing
Li (g) → Li+(g) + e− IE1 = 520 kJ
Li+( (g) → Li2+(g) + e− IE2 = 7297 kJ
___________________________________
Li (g) → Li2+(g) + 2e− IEtotal = 7817 kJ
IE1  IE2  IE3 …

8. Ionization energy vs. atomic number
Ca
K
H
He
Li
B
Be C
N
O
F
Ne
Na
Mg
Al
Si
P S
Cl
Ar
0
500
1000
1500
2000
2500
0 2 4 6 8 10 12 14 16 18 20
Element
Ionization
energy
(kJ/mol)
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9. Ionic sizes
• When atom gains an electron (anion), the size increases i.e. sizes of ion is
greater than neutral atom. Generally: X3−  X2−  X−  X
• When electrons are added mutual repulsion between electrons increases.
Electrons push each other apart
• When electrons are removed (cation) the size decreases i.e. cation size is
smaller than the neutral atom. Generally: A3+  A2+  A+  A
• Electron-electron repulsions reduce when electrons are removed allowing
remaining electrons to be drawn closer around the nucleus.
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10. Ionisation energy
• The energy required to remove an electron from an isolated, gaseous atom or
ion in its ground state.
X (g) → X+(g) + e−
• Measures amount of work needed to pull out an electron – gives an idea of how
tightly bound the electron is to the nucleus.
• Requires energy input – the tightly bound, the more energy required to remove
the electron.
• Successive ionisation energies become increasingly larger
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11. Electron Affinity
• Electron affinity reflects the ability of an
atom to accept an electron.
• It is the energy change that occurs
when an electron is added to a
gaseous atom.
Example
F(g) + e− → F−(g) EA = − 328 kJ mol
−1
Note that the process is exothermic when
F gains an electron.
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• For a free electron approaching an atom
there is opposing attractive effects of the
nucleus and repulsive forces in the outer
electrons
• If attractive forces exceed the repulsive, an
electron is gained and energy is given off
• By gaining an electron, F−, a fluorine atom
acquires the very stable a noble gas
configuration of Ne
F (1s22s22p5) + e− → F− (1s22s22p6)
Why gain an electron?

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Electron Affinity

13. Electronegativity
• Electronegativity is a measure of the attraction of an atom for the electrons in
a chemical bond.
• The higher the electronegativity of an atom, the greater its attraction for
bonding electrons.
• It is related to ionization energy. Electrons with low ionization energies have
low electronegativities because their nuclei do not exert a strong attractive
force an electrons.
• Elements with have high ionisation energies have high electronegativities due
to the strong pull exerted on electrons by the nucleus (effective nuclear
charge)
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14. Electronegativity and Bond Polarity
• The difference in electronegativity between two atoms gives
a measure of the polarity between two atoms
• Suppose a positive and negative charge are separated from
one another
Direction of dipole
Dipole moment = distance x charge
 = q x d
Note: q is Coulomb and d in m
q + q−
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