Transcript Lesson 3 & 4 Covalent radii and ionisation

(B) Periodicity
(B) Periodicity
Covalent radius and Ionisation
After completing this topic you should be able to :
• The covalent radius is a measure of the size of an atom.
• The trends in covalent radius across periods and down
groups can be explained in terms of the number of occupied
shells, and the nuclear charge.
• The trends in ionisation energies across periods and down
groups can be explained in terms of the atomic size, nuclear
charge and the screening effect due to inner shell electrons.
Covalent radius
Atomic Size
There is no definite
edge to an atom.
However, bond lengths
can be worked out.
Atomic radius,
½ the distance between nuclei.
To find the bond length, add 2 covalent radii together.
pm = picometre X 10
– 12
m
Variation of covalent radius with atomic number
Adapted from New Higher Chemistry E Allan J Harris
The covalent radii of the elements in
any period decrease with increasing atomic number.
Variation of covalent radius with atomic number
Adapted from New Higher Chemistry E Allan J Harris
The covalent radii of the elements in
any group increase with increasing atomic number.
Variation of covalent radius with atomic number
Adapted from New Higher Chemistry E Allan J Harris
No values are given for the Nobel gases
Why?
Unreactive so do not form
bonds
Covalent radius
COVALENT RADIUS
Explain the change in covalent radius as you go along a period.
The covalent radius of an element is half the distance
between the nuclei of two of its covalently bonded
atoms.
The covalent radius decreases as you go along a period
As you go along a period there is a greater positive
charge on the nucleus The shells or energy levels of
electrons are more strongly attracted to the nucleus
and therefore the size of the atoms decreases.
COVALENT RADIUS
Explain the change in covalent radius as you go down a group.
The covalent radius of an element is half the
distance between the nuclei of two of its
covalently bonded atoms.
The covalent radius increases as you go down a group.
As you go down a group there are more energy levels of
electrons. The outer electrons are further away from
the nucleus so the atoms are larger.
Ionisation energy
Ionisation energies
This is defined as "the amount of energy
required to remove one mole of electrons from
one mole of atoms in the gaseous state”
Energy
e e
M (g)  M+(g) + e 1st ionisation
+
M +(g)
The outermost electron will be
the most weakly held and is
removed first
Ionisation energies
This is defined as "the amount of energy required to
remove one mole of electrons from one mole of atoms in
the gaseous state”
Energy
e
M (g)  M+(g) + e 1st ionisation
e
+
M
2+
(g)
M(g)+  M(g)2+ + e 2nd ionisation
ionization energy
The ionization energy of a chemical element, expressed in joules (or
electron volts), is usually measured in an electric discharge tube in
which a fast-moving electron generated by an electric current collides
with a gaseous atom of the element, causing it to eject one of its
electrons.
The ionisation energy is an enthalpy change and
therefore is measured per mole.
Units kJmol-1 (kilojoules per mole).
Ionisation energies kJ mol-1
Overall increase along period
Decrease down group
Li
526
Be
905
B
807
C
N
O
F
Ne
1090 1410 1320 1690 2090
Na
502
K
425
Rb
409
Mg
744
Ca
596
Sr
556
Al
584
Ga
577
In
556
Si
792
Ge
762
Sn
715
P
1020
As
953
Sb
816
S
1010
Se
941
Te
870
Cl
1260
Br
1150
I
1020
Ar
1530
Kr
1350
Xe
1170
First and Second ionisation energies of the first 20 elements
Adapted from New Higher Chemistry E Allan J Harris
In each period there is an overall
increase peaking at the noble gas
FIRST IONISATION ENERGY
Explain the change in first ionisation energy as you go
along a period.
The first ionisation energy is the energy required
to remove 1 mole of electrons from 1 mole of
atoms in the gaseous state.
The first ionisation energy increases as you go
along a period
As you go along a period there is a greater
positive charge on the nucleus. There is a
greater attraction between the outer electron
and the nucleus. More energy needs to be
supplied to remove the electron.
First and Second ionisation energies of the first 20 elements
Adapted from New Higher Chemistry E Allan J Harris
Down a group first ionisation energy
decreases
FIRST IONISATION ENERGY
Explain the change in first ionisation energy as you go
down a group.
The first ionisation energy is the energy required to
remove 1 mole of electrons from 1 mole of atoms in
the gaseous state.
The first ionisation energy decreases as you go down
a group.
As you go down a group there are more energy levels
of electrons. The outer electron is further from the
nucleus. The inner electrons shield the outer
electron from the effect of the nucleus. Less
energy is needed to remove the outer electron.
First and Second ionisation energies of the first 20 elements
Adapted from New Higher Chemistry E Allan J Harris
For each element the second ionisation energy is
higher than the first ionisation energy.
Explain why the second ionisation energy of an element
is always greater than the first ionisation energy:
First ionisation energy –
first mole of electrons
removed
Second ionisation energy –
second mole of electrons
removed
M(g)  M+(g) + e
M+(g)  M2+(g) + e
In the second ionisation energy negative electrons
are being removed from positive ions rather than
neutral atoms.
In the positive ion there is a greater attraction for
the electrons so more energy is needed to remove
the second mole of electrons.
Explain why the second ionisation energy of K is much
greater than the second ionisation energy of Mg:
K
(g)
2,8,8,1
 K+ (g) + e
2,8,8
K+ (g)  K2+ (g) + e
2,8,8
2,8,7
Mg (g)  Mg+ (g) + e
2,8,2
2,8,1
Mg+ (g)  Mg2+ (g) + e
2,8,1
2,8
The second ionisation of K
involves removing an
electron from a stable
electron arrangement.
The second ionisation of
Mg involves removing an
electron to form a stable
electron arrangement.
This requires a lot of energy
This requires less energy
Successive ionisation Energies
first ionisation energy
E(g)
second ionisation energy E+(g)
third ionisation energy E 2+(g)
fourth ionisation energy E 3+(g)
E+(g)
+ eE 2+ (g) + eE 3+ (g) + eE 4+ (g) + e-
- ionisation energies increase as successive electrons
are removed
- removing an electron from a filled inner shell
requires a large increase in energy
The first four ionisation energies of aluminium, for
example, are given by
Al(g)
Al+(g)
+ e1st I.E. = 577 kJ mol-1
Al+(g)
Al2+ (g) + e2nd I.E. = 1820 kJ mol-1
Al2+(g)
Al3+ (g) + e3rd I.E. = 2740 kJ mol-1
Al3+(g)
Al4+ (g) + e4th I.E. = 11600 kJ mol-1
In order to form an Al3+(g) ion from Al(g) you would have
to supply:
577 + 1820 + 2740 = 5137 kJ mol-1
First and Second ionisation energies of the first 20 elements
Adapted from New Higher Chemistry E Allan J Harris
It is worth noting the Nobel gases have the highest value
for each period. This goes some way to explaining the
great stability of filled orbital's and the resistance of
the Nobel gases to form compounds.
Nobel gas compounds
If some other change can compensate for the energy
required then ionic compounds of Nobel gases can be
made.
Can you suggest why the first Nobel gas compound
prepared contained Xe?
Questions for you to try:
1.Explain why the third ionisation energy of
Magnesium is so much greater than its second.
Removing the third mole of electrons
involves breaking into a stable complete
energy level of electrons.
2. Calculate the energy change for the following
changes.
a)Ca (g) → Ca2+ (g) + 2e596 + 1160 = 1756 kJmol-1
b)B2+(g) → B4+ (g) + 2e3660 + 25000 = 28660 kJmol-1
3. Which of the following equations represents
the first ionisation energy of fluorine?
A F–(g) → F(g) + e–
B F–(g) → F2(g) + e–
C F(g) → F+(g) + e–
D F2(g) → F+(g) + e–
C
4. Which line in the table is likely to be
correct for the element francium?
State at 30 °C
First
ionisation
energy/kJ
mol–1
A
solid
less than 382
B
liquid
less than 382
C
solid
greater than
382
D
liquid
greater than
382
B
5. As the atomic number of the alkali metals increases
A the first ionisation energy decreases
B the atomic size decreases
C the density decreases
D the melting point increases.
A
6. Explain why
a) A potassium atom is larger than a sodium atom
Potassium has an extra energy level (shell) of
electrons.
b) The Chlorine atom is smaller than a sodium atom
Both atoms have the same number of energy levels,
but the chlorine has a greater nuclear charge than
sodium. This attracts the outer electrons more
strongly and causes the decrease in atomic radius.
7. Atoms of different elements are different sizes.
What is the trend in atomic size across the period
from sodium to argon?
Decreases or gets smaller
8. Which of the following reactions refers to the
third ionisation energy of aluminium?
A Al(s) → Al3+(g) + 3e–
B Al(g) → Al3+(g) + 3e–
C Al2+(g) → Al3+(g) + e–
D Al3+(g) → Al4+(g) + e–
C
9. Atoms of different elements have different
ionisation energies.
Explain clearly why the first ionisation energy of
potassium is less than the first ionisation energy of
sodium.
Potassium has more electron shells (or outer electron is
further from the nucleus)
The inner electrons (electron shells) shield (screen) the
outer electron from the attraction of the nucleus.
Therefore the outer electron is held less tightly in
potassium
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