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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Lecture 6 Chemical Properties and the Periodic Table
________________________________________________________________________________
Part 1
Part 2
Part 3
Part 4
Part 5
Periodic Properties of the Elements
Groups 1a & 2a
Metals and Non-metals
Groups 3a to 8a; the Non-metals
Groups 3b to 12b; the Transition Metals
________________________________________________________________________________
Part 1
Periodic Properties of the Elements
We have already seen how electronegativity increases on moving across the
periodic table. Can you remember why? As a result electronegativity is a
periodic property, so for example, the halogens (Group 7a elements) are all
highly electronegative.
Many other atomic properties are periodic because of the periodic nature of
electronic configurations of the atoms.
Note:
As you go down a group, the 1s orbital is moves closer to the nucleus,
because of greater attraction due to the increase in the nuclear charge, Z.
For many purposes we can consider the nucleus and inner (non-valence)
electrons as being an effective field in which the valence electrons move.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
In other words the effective nuclear charge, Z*, experienced by any given
electron is given by:
Z* = Z - σ
Where
Z is the actual nuclear charge
σ is the screening constant
Slater’s Rules are used to estimate σ
1. Write down electron configuration :
eg
(1s) (2s, 2p) (3s, 3p, 3d) (4s, 4p, 4d)
2. Electrons to the right of the electron of interest contribute zero to σ
3. Electrons in the same (s,p) group contribute 0.35 to σ,
for 1s the value used is 0.30
4. All electrons in the (n-1) group contribute 0.85 to σ
5. All electrons in the (n-2) and lower contribute 1.00 to σ
Consider He, Ne, Ar:
He
1s2
1s electrons
Z*
Ne
=
2 - 0.30
=
1.7
=
9.7
1s2 2s2 2p6
Z*
=
10 - σ
For a 2s or 2p electron in Ne
Z*
Ar
=
1s2 2s2
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10 -( 7 x 0.35) - (2 x 0.85) = 5.85
2p6 3s2
3p6
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
1s electron
Z*
=
18 - 0.30
=
17.70
For a 2s or 2p electron in Ar
Z*
=
18 - 4.15
=
13.85
Z*
=
18 - ( 7 x 0.35) - (2 x 0.85) = 13.85
For a 3s or 3p electron in Ar
Z*
=
18 - ( 7 x 0.35) - (8 x 0.85) - 2 = 6.75
Consequences:
1. Quantum mechanics suggests that electrons in an atom can be described in
terms of their most probable distance from the nucleus
2. For a given shell, the electron distribution function moves closer to the
nucleus as Z increases
3. As Z increases, the inner shells are more tightly held and only the outer
more diffuse electrons can be disturbed
These findings affect many atomic properties, so there are several periodic
properties:
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(a) Atomic Radii:
The atomic radius can be calculated by measuring the internuclear distance
in compounds of like elements e.g. H2, N2, Cl2
Where the atomic radius is taken to be half the internuclear distance between
like atoms.
Trends in the atomic radius are observed:
1. Across a period e.g. Li → F, size of the atom decreases.
This is because Z increases, but each electron added does not fully
shield the outer electrons from the nuclear charge.
⇒ Z* increases and the size reduces.
2. Going down a group the size of the atom increases
This is because we are putting electrons into orbitals of higher
principal quantum number, which are further from the nucleus.
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(b) Ionisation Energy:
Is defined as the energy required to remove an electron from an atom or ion M(g) → M+ (g) + e
M+ (g) → M2+ (g) + e
- 1st IE
- 2nd IE etc.
For any element 1st IE < 2nd IE < 3rd IE …etc...
The repulsive forces between electrons help in the removal of electrons
On removing an electron, the screening of one electron by the others is
reduced, therefore the effective nuclear charge (Z*) increases and further
removal of electrons becomes more difficult.
For example silicon, Si
[Ne] 3s2 3p2
Ionisation energies (kJ/mol)
1st
780
2nd
1,575
3rd
3,220
4th
4,350
5th
16,100
IE (1) → IE (4) gradual increase as 3s and 3p electrons removed
Then, a large increase in IE (5) as a 2p electron is removed. This electron is
closer to the nucleus (lower n) and much more tightly held.
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Across a Period e.g. Li → Ne
Alkali elements have lowest IE,
while inert gas elements have highest IE.
Note
There are a number of discrepancies in going across a period;
i.e. the trends are not perfect.
1. For B, C, N the IE increases, but the increase is lower than might be
expected. Because we are now removing p electrons, these don’t
penetrate as close to the nucleus as do s electrons; therefore they are more
easily removed.
2. For O and F the IE are again lower than predicted, because we are now
removing electrons that are spin paired. It is easier to remove these
because the electrons are repelling each other, which aids the removal of
electrons.
The trend going down any Group:
• The 1st IE decreases significantly.
• This is because of the increasing number of inner filled shells screening
the valence electrons, which become easier to remove.
• So the following process becomes more favourable energetically:
M
→
M+
+
1e-
Later we will see that this is associated with the generally observed increase
in metallic character going down a group.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
(c) Electron Affinity:
Is the energy change that takes place when an electron is added to a gaseous
atom or ion:
M(g) + e- → M-(g)
M(g)+ + e- → M(g) reverse of I.E.
A negative value for EA indicates a release of energy when the anion is
formed, i.e. that the anion is more stable than the neutral atom.
Consider the halogens:
Hal (ns2 np5) + e → hal- (ns2 np6)
This should be favoured since X- has a full octet.
E.A (kJ/mol)
→
F-
-330
Cl
Cl-
-350
Br
Br-
-325
I
I-
-295
F
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The electron goes into an orbital further from the nucleus as we go down the
group (F→I), you would predict that the EA should decrease.
So you would expect F to have highest EA
It turns out that electron-electron repulsion terms are more important in this
case and they counterbalance the above stabilisation.
⇒ the EA of all the halogens are similar.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Part 2
Groups 1a & 2a
The chemistry of the elements in Groups 1a and 2a are determined by their
low electronegativities & ionisation energies and also by their relatively
large atomic radii.
Group 1a; hydrogen
Having said this, the chemistry of hydrogen is completely different from the
other elements in Group 1a:
Unlike the other elements it has a relatively high IE and electronegativity, it
commonly forms covalent bonds, particularly to carbon. These properties
arise because of its very small atomic radius.
Exercise:
Explain (i) why hydrogen is so small, (ii) why its small size results in
high IE and electronegativity.
Note:
When we refer to the chemistry of hydrogen, we are actually talking about;
H
1s1
the hydrogen atom;
1 proton and 1 electron
1s2
the hydride anion;
1 proton and 2 electrons
or
H-
The latter can and does occur, as the hydrogen atom has moderate electron
affinity.
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On the other hand, the hydrogen cation:
H+
1s0
consists of just a proton
This can occur (though always associated with other ions).
It is the entity, which is transferred in acid-base chemistry.
So an acid is a proton (or H+) donor and a base is a proton acceptor.
This type of behaviour will be dealt with later in terms of the molecules,
which have this ability.
The chemistry of hydrogen is dominated by covalent bond formation in
which hydrogen shares electrons with other elements
H•
+
•X →
H:X
If X not equal to hydrogen, the covalent bond will be polar
If X has high electronegativity e.g. fluorine, F
HF
δ+
δ-
H
F
-the δ refers to the molecular centre of charge
hydrofluoric acid
If X is less electronegative than hydrogen (i.e. is electropositive), e.g. Na
δ-
δ+
H
Na
HNa sodium hydride.
Metal hydrides generally are salt-like materials, i.e. like the ionic NaCl.
Covalent hydrides:
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Hydrogen forms strong stable single covalent bonds with elements
including: B, C, N;
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eg. BH3, CH4, NH3
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Group 1a continued, the alkali metals
Li
[He] 2s1
Na
[Ne] 3s1
K
[Ar] 4s1
Rb
[Kr] 5s1
Cs
[Xe] 6s1
Fr
[Rn] 7s1
• Na and K are the most abundant
• Due to low electronegativity and ionisation energy the compounds
formed are mainly ionic although some covalent bonding does occur.
i.e. the dominant process in their chemistry is:
M
→
M+
+
1e-
The loss of an electron is termed oxidation;
so alkali metals are easily oxidised.
Reactions of Alkali Metals
• Reaction with Halogens (Group 7a elements)
Due to the large electronegativity difference, all known halides are ionic
(CsF). So these reactions invariably result in the formation of salts.
2M
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+
X2
→
2MX
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
• Oxygen compounds
They react rapidly with oxygen; different types of oxides can exist
Li
+
O2
→
Li2O
Na
+
O2
→
Na2O2
O2
→
KO2
K
+
Exercise: Draw Lewis structures for the three oxides.
Hint: The lithium compound contains the oxide ion [O]2-, the sodium
compound contains the peroxide ion [O2]2- and the potassium compound
contains the super-oxide ion [O2]-.
Hydroxides
Strongly basic hydroxides are formed when the metals react with water:
Li
+
H2O
→
LiOH
+
H2
Na
+
H2O
→
NaOH
+
H2
K
+
H2O
→
KOH
+
H2
Exercise:
On the basis of the known trends in electronegativity and ionisation energy
going down the group; which of these reactions leading to the formation of
Group 1a hydroxides is the most vigorous? -In fact it is explosive!
Reaction with ammonia
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The alkali metals react with ammonia, NH3, to yield H2 gas and a metal
amide
M
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+
NH3 →
MNH2
+
H2
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Group 2a, the alkaline earth metals
Be
[He] 2s2
Mg
[Ne] 3s2
Ca
[Ar] 4s2
Sr
[Kr] 5s2
Ba
[Xe] 6s2
Ra
[Rn] 7s2
• Like the alkali metals they tend to oxidise, in this case they usually lose
two electrons to yield dipositive ions:
→
M
M2+ +
2e-
• They tend to be less reactive than the group 1a elements
-this is because the first ionization energies are larger
• They react with halogens to form ionic halides, which are usually salts.
Ca
+
Cl2
→
CaCl2
Beryllium is exceptional in the Group; BeF2 and BeCl2 are molecular
substances. This is because of Be’s higher electronegativity.
• They react with oxygen to form oxides
2 Mg +
O2
→
2 MgO
• Again, with the exception of Be the alkaline earth metals react with water
to yield metal hydroxides M(OH)2, which are weakly basic:
Mg
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+
2H2O →
MgOH
+
H2
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Part 3
1a
1
3
11
Na
19
37
2a
H
Li
4
4b
5b
6b
7b
8b
9b
12
Rb
38
Ca
56
Sr
Ba
B
13
Mg
Sc
39
57
22
Y
40
La
72
Ti
23
V
24
Nb
42
Zr
41
Hf
73
Ta
Cr
25
Mo
43
74
W
Mn
75
26
Fe
27
Co
28
45
Rh
46
Tc
44
Ru
Re
76
Os
77
5a
6a
7a
Ir
Ni
Pd
47
Ag
48
Cd
49
79
Au
80
Hg
81
78
Pt
Zn
31
Al
29
Cu
30
8a
2
5
21
4a
10b 11b 12b
Be
K
Cs
3a
3b
20
55
Metals and Non-metals
Ga
6
C
14
32
Si
Ge
7
N
15
33
P
8
16
17
Se
35
As
Sb
52
50
Sn
51
Th
82
Pb
83
Bi
84
F
S
34
In
O
9
Te
Po
He
10
Ne
Cl
18
Ar
Br
36
Kr
53
85
I
54
At
86
Xer
Rn
metals
non-metals
There is an approximate division between metals and non-metals.
This is shown in the PTE above.
However, around the division the metallic elements have some non-metallic
properties, and visa versa.
Metals
• In a metal there are no molecules.
• Instead the atoms are arranged in ordered arrays (or grids), similar to the
arrays of atoms in the crystals of NaCl.
• However, unlike in silicon, in metals the electrons in the bonds are not
fixed in positions.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
each sphere represents
an atom in the metal
each atom is bonded
to the atoms around it
the electrons in each bond
(three are drawn in here)
are continuously swapping
positions, randomly
The electrons in metals are moving around continuously.
In fact although the atoms in a metal have fixed positions, the electrons
behave as if they were in a “gas”, i.e. they move about (diffuse) at random.
the electrons are better
thought of as being a “gas”,
shown here as a blue haze
The electrical conductivity of metals arises from the fact that these electrons
are mobile and so can carry electric charge through the metal;
i.e. there are “mobile charge carriers” available.
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Orbital view of metals:
• In a diatomic molecule, eg H2, the two AOs give rise to two MOs; a
bonding and antibonding pair.
• So initially there are two isolated AOs of the same energy, but on
bonding two states of different energy are formed.
• It is always the case that the more orbitals overlap the more states are
formed.
• In a metal there are so many orbitals overlapping that effectively an
infinite number of states with different energies arise; a band.
2nd row diatomic
molecule
H2
metal
empty
anti-bonding
band
σ∗2pz
σ∗1s
*
π*2px, π2p
y
σ1s
more
orbitals
π2px, π2py
σ2pz
band gap
more
orbitals
σ∗2s
partially
filled
conducting
band
σ2s
• This band is a state (orbital) extending over the metallic solid.
• An occupied band is called the valence band. If it is full, or the gap to the
empty upper band (antibonding states) is significant, the electrons are not
free to move.
• In metals one of these criteria is met, so metals conduct electricity.
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This simple model can be used to rationalise the electrical conductivities of a
range of materials:
metal
metal
semi-conductor
insulator
∆E
∆E
valence band is
not full, so it
serves as the
conducting band
conducting band
full valence band
as in semiconductors
overlaps the
doesn’t overlap with
but the energy
valence band the empty conducting
gap is large
band but the energy
gap is small
• Metals conduct electricity due to the presence of a partially filled
conducting band.
• Semiconductors conduct electricity if the energy gap to the conducting
band can be jumped, so if the temperature is raised conductivity
increases.
More commonly a low concentration of an impurity is added which
introduces some extra electrons into the system (an n-type implant).
• Insulators do not conduct electricity, as very few electrons can jump the
energy gap.
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The Physical Properties of Metals:
•
Under normal conditions they are usually solids, with high melting and
boiling points. For instance copper, Cu, melts at 3000°C.
There is one exception; mercury, Hg, is a liquid at room temperature.
•
Because of the close packing of the atoms, they are dense materials. So
a given volume of a metal will be heavier than other materials.
•
They usually are hard materials, but they are also usually malleable,
i.e. they can be shaped. There are some exceptions, in particular the
Group 1a elements (alkali metals), which are relatively soft solids.
•
They are good conductors of heat and electricity.
The Chemical Properties of Metals:
These arise because they tend oxidise (lose electrons), for example:
Fe →
•
Metals react with oxygen to form metal oxides, which are bases.
4Fe(s) + 3O2(g) →
•
Fe3+ + 3e-
2Fe2O3(s)
RUST
Metals react with water to form Metal hydroxides, which are bases, and
hydrogen.
Na(s) + 2H2O (l) → 2NaOH (l)
•
+ H2(g)
Metals react with acids to form metal salts and hydrogen.
Zn(s) + H2SO4 (l) → ZnSO4 (l)
W(s) + H2SO4 (l) →
+ H2(g)
no reaction
so tungsten is a less active metal
•
Metals form chlorides, which are ionic materials.
2Na(s) + 2HCl(g) → 2NaCl (l)
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+ H2(g)
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Non-metals
As we have seen previously, non-metallic elements tend to form molecular
substances. These molecules do not pack as closely as the atoms in a metal.
H
Si
H
weak force of
attraction
H
Si
H
H
H
H
H
Nearby molecules do not approach each other very closely.
The only attractive forces are quite weak, these weak intermolecular forces
are called “Van der Waals interactions”, they are only important for very
large molecules.
This is in contrast to the strong forces which hold metals and ionic solids
together, so metals and ionic solids are dense materials.
The Physical Properties of Non-Metals:
•
•
•
•
They have low melting and boiling points.
They have low density compared for instance to metals.
They form soft solids liquids or gases (exceptions include diamond).
They are poor conductors of heat and electricity.
e a poor conductor, so in Intel it is used as an insulator between
metallic layers (exceptions include graphite, a form of carbon, which
conducts heat and electricity).
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
The Chemical Properties of Non-Metals:
•
They usually gain electrons to form negative ions (anions).
F + 1e- →F-
•
When reacted with oxygen they form oxides, which are acidic
C(s) + O2(g)
CO2 (g) + H2O(g)
•
•
→
CO2 (g)
H2CO3 (g)
They do not react with dilute acids.
They form chlorides, which are liquids or gases. So CCl4, SiF4 and
SiH2Cl2 are volatile liquids or gases at standard temperature and
pressure (STP).
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Part 4
Groups 3a to 8a; the Non-metals
The chemistry of the elements from Group 3a to 8a are determined by their
moderate to high electronegativities & ionisation energies and to a lesser
extent by their larger atomic radii & moderate electron affinities. So they
tend to form covalent bonds which result in the formation of molecular
substances.
Group 3a, boron and the electron deficient elements
5
[He] 2s22p1
B
13
Al
[Ne] 3s23p1
31
Ga
[Ar] 4s24p1
49
In
[Kr] 5s25p1
• On going down the group, the metallic character of the elements
increases.
• Boron is unique in the group in that it is clearly a non-metal, we will
concentrate on its properties, as it is very interesting.
• The molecules boron forms are unique in that they do not conform fully
to Lewis theory, for instance BH3 is a stable molecule, but there is no
octet of electrons on boron. In fact there are only 6 electrons formally.
• For this reason we say that boron forms electron-deficient molecules.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Exercise:
Firstly draw the Lewis structure of ethane, C2H6, and then of diborane, B2H6.
Do you agree that the only possible structure for B2H6 is the following?
H
H
H B B H
H
H
2-
However, both ethane and diborane exist; they are neutral (uncharged) and
are covalent.
If you draw B2H6 as a neutral molecule with a boron-boron bond there are
not enough electrons to form all the bonds
But B2H6 exists and like C2H6 it is neutral (uncharged) and covalent.
The actual structure is now accepted as:
H
H
B
H
H
B
H
H
• Here there are only really six bonds (12 electrons); the two bridging B-HB groups are held together by a pair of three-centre two-electron bonds
and there is no direct boron-boron bond.
• This is a classical case of the failure of the VB/Lewis approach.
• However MO theory can easily accommodate this type of bonding;
an MO containing two electrons arising from the overlap of AOs on three
atoms is perfectly acceptable in MO theory.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Group 4a, the carbon group
6
[He] 2s22p2
C
14
Si
[Ne] 3s23p2
32
Ge
[Ar] 4s24p2
50
Sn
[Kr] 5s25p2
• Carbon is unique in the group in that it is clearly a non-metal.
• It is also the element, from which by far the greatest number of molecules
are formed; the chemistry of carbon is called organic chemistry.
• This great diversity exists because of the carbon-carbon and carbonhydrogen bonds are stable covalent bonds.
• Organic chemistry does not form part of this course.
• The chemistry of silicon is less rich than that of carbon, partially because
Si-Si and Si-H bonds are weaker.
• On going down the group, the metallic character of the elements
increases, so tin, Sn, is clearly a metal.
Some organic molecules:
NH2
N
H
N
OH
O
HO
N
H
O
H
H2N
C
OH
OH
H
CH3
glucose - a sugar
is a six-membered ring involving
five carbons and an oxygen
H
H
OH
H
H
HO
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N
HO
O
H
HO
CH
alanine - an amino acid
is one of the building blocks of proteins
it has a carbon-nitrogen backbone
H
the adenine nucleoside
- is one of the building blocks of DNA
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
The chemistry of silicon, Si:
• Unlike carbon, silicon does not tend to form large molecules as S-S and
S-H bonds are weaker than the carbon counterparts.
• The major applications of silicon are in the semiconductor industry.
Silicon oxidation:
Si(s) + O2 (g)
→
SiO2 (s)
Silicon oxide is an electrical insulator, so it is used to electrically
isolates individual transistors and layers on an IC.
Furthermore it acts as a contaminant sink, i.e. contaminants from
various parts of the chip tend to gather in the oxide, where they
are subsequently removed (sacrificial oxide).
Formation of silicon nitride:
3SiH2Cl2 (g) + 4NH3(g) → Si3N4 (s) + 6HCl(g) + 6H2(g)
Silicon nitride (Si3N4) is a white, chemically inert (i.e.
unreactive), amorphous powder (i.e. unlike silicon, it has no
defined crystal structure). It has the ability to form thin films.
• These reactions take place using chemical vapour deposition (CVD), the
reactants are gases which coat the surface of the solid silicon wafer with a
silicon derivative.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Group 5a, the nitrogen group
7
[He] 2s22p3
N
15
P
[Ne] 3s23p3
33
As
[Ar] 4s24p3
51
Sb
[Kr] 5s25p3
• Again, on going down the group, the metallic character of the elements
increases.
• Nitrogen and phosphorus exhibit valencies of 3 or 5; thus PCl3 and PCl5
both exist, and in both cases the phosphous atom is neutral.
Chemistry of Nitrogen
• Dinitrogen, N2, is an extremely stable molecule. This is because of the
presence of a covalent triple bond (draw the Lewis structure to convince
yourself). Air in 80% N2.
• Ammonia, NH3, is a very important molecule, it is produced
commercially on a huge scale in the Haber process:
N2 (g)
+
3H2(g)
2NH3(g)
- historically this was one of the first industrial scale chemical reactions.
The ammonia produced is used in the fertiliser, explosive and polymer
industries.
• There are at least six stable oxides of nitrogen; NO2, NO, N2O, N2O3,
N2O4 and N2O5, in which the nitrogen atoms exhibit a range of oxidation
states (formal charges) from +1 to +5.
As an exercise you could try drawing some of these.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Chemistry of Phosphorus
• The most important phosphorus molecules are phosphates or
phosphorous oxides, which (like nitrates) are important fertilisers.
• The most common form is the (PO4)3- anion, which is relatively stable
due to delocalisation of the negative charge. Can you draw its Lewis
structure?
• The true structure (the average of the four possible equivalent Lewis
structures) is a tetrahedron with the “3-” charge is distributed evenly over
the four oxygen atoms, which are positioned at the apices of the
tetrahedron (the phosphorus atom is at the centre).
• On heating, phosphate groups have a tendency to lose water and form
polyphosphates.
2H3PO4
→
H4P2O7
+
H2O
Pyrophosphoric acid, H4P2O7, is used for etching in the semiconductor
industry.
• Long-chain polyphosphates can also form; their structure is a series of
tetrahedrons, linked at the apices.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Group 6a, the oxygen group
8
[He] 2s22p4
O
16
S
[Ne] 3s23p4
34
Se
[Ar] 4s24p4
52
Te
[Kr] 5s25p4
• Again, on going down the group, the metallic character of the elements
increases.
• Oxygen gas, O2 is a highly oxidising substance:
i.e. atoms to which oxygen adds are usually oxidised (they lose electrons)
due to the high electronegativity of oxygen.
• The most important compound on the planet is probably water, H2O. As
we saw earlier, this is a liquid at STP (standard temperature and
pressure), even though it is a very light molecule because of the extensive
hydrogen-bonding. Life on earth evolved in water and is dependent on it;
all animal and planet cells are composed mostly of water.
• The difference between oxygen and sulphur is well illustrated by the
different properties of H2O and H2S. As we saw earlier H2S has a
different shape (because of the larger 3p orbitals) and it is a gas at STP
because of the reduced hydrogen bonding (S is less electronegative).
• The major compounds of sulphur are its two oxides; SO2 and SO3.
In water sulphur dioxide is in equilibrium with sulphurous acid
SO2 +
H2O
H2SO3
In water sulphur trioxide is in equilibrium with sulphuric acid,
SO3 +
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H2O
H2SO4
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
This is a much stronger acid than sulphurous acid, because the
[HSO4]- anion is more stable than the [HSO3]- anion.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Group 7a, the halides
9
[He] 2s22p5
F
17
Cl
[Ne] 3s23p5
35
Br
[Ar] 4s24p5
53
I
[Kr] 5s25p5
• The chemistry of this group is dominated by their very high
electronegativity, which decreases going down the group.
• They have a strong tendency to form very stable salts with Group 1a and
2a elements, e.g. NaF, NaCl etc...
• They form some of the strongest acids known, acids are proton donors.
Consider what happens to these molecules in solution:
HF + H2O
F- + H3O+
HCl + H2O
Cl- + H3O+
HF is a stronger acid than HCl, i.e. the proton concentration,[H+], is
higher (the equilibrium is even further to the right than for HCl). This is
due to the greater electronegativity of fluorine.
Chemical equilibria and acid-base chemistry will be dealt with in the
2nd chemistry module.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Group 8a, the inert (or noble) gases
2
He
1s2
10
Ne
[He] 2s22p6
18
Ar
[Ne] 3s23p6
36
Kr
[Ar] 4s24p6
54
Xe
[Kr] 5s25p6
• The chemistry of this group is dominated by their full outer shell, which
means they tend not to react as they already have an octet.
• There are some exceptions to this, particularly going down the group;
xenon which has a full 5th shell which is very large.
• Xe does form some compounds with strong oxidising agents (materials
which remove an electron) or electronegative elements such as fluorine.
• Thus
PtF6 + Xe
→
[Xe]+PtF6-
… and compounds such as XeF2 and XeF4 exist.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Part 5
Groups 3b to 12b; the Transition Metals
The chemistry of the elements in Groups 3b to 12b is influenced by their
moderately low electronegativity & ionisation energies. In this sense they
are metallic like the elements of Group 1a and 2a. However their chemistry
is dominated by the presence of partially filled d-orbitals.
Remember the order of orbital energy in atoms:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p
This means that the 4s is usually full.
So the electronic configuration of (26Fe)2+, which has 24 electrons, is:
(26Fe)2+;
1s22s22p63s23p64s23d4
• Transition metals can exist with variable oxidation state and have strong
catalytic activity.
• So materials such as FeCl2 and FeCl3 can co-exist, in which iron has a
formal charge (oxidation state) of +2 and +3 respectively.
• The chemistry and bonding of transition metal compounds (more
generally called transition metal complexes) is a huge subject. The
theories of bonding we have met in this course are severely tested by
these types of materials.
• A detailed examination of transition metal chemistry is beyond the scope
of this course.
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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6
Summary
Lecture 6 Chemical Properties and the Periodic Table
• A number of properties including electronegativity, atomic radii,
ionisation energies and electron affinities are periodic in nature.
• For this reason the chemistry of most of the elements is periodic, i.e.
elements on the same Group react in similar ways to form similar
compounds.
• The major division of the periodic table is into metals and the nonmetals. The properties of each of these two classes of materials are
readily understood using simple arguments based on the periodic
properties and MO theory.
• The chemistry of the transition metals is dominated by the presence of
partially filled d-orbitals.
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