Transcript Periodicity Physical properties

Periodicity
Physical Properties
First Ionisation enegy/ kJmol-1
Ionisation energies
He
2500
Ne
2000
Ar
Kr
1500
1000
500
Li
Na
K
Rb
0
0
10
20
30
Atomic number
40
50
The elements in group 1 have the lowest value in each period.
As we descend group 1 from Li  Cs the values decrease as
the outer electron is further from the nucleus and already in a
higher energy level so less energy is needed to remove it.
Across a period as each energy level is being filled, an extra
proton is also being added to the nucleus. As each electron is
added it is more strongly attracted to the increased charge on
the nucleus so becomes lower in energy and is harder to
remove.
The exceptions in this trend are due to the existence of sublevels within the main energy level.
The noble gases have the highest first ionisation energy in
each period.
Electronegativity
H
2.1
He
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Ne
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
Ar
K
0.8
Ca
1.0
Sc
1.3
Ti
1.5
V
1.6
Cr
1.6
Mn
1.5
Fe
1.8
Co
1.8
Ni
1.8
Cu
1.9
Zn
1.6
Ga
1.6
Ge
1.8
As
2.0
Se
2.4
Br
2.8
Kr
Rb
0.8
Sr
1.0
Y
1.2
Zr
1.4
Nb
1.6
Mo
1.8
Tc
1.9
Ru
2.2
Rh
2.2
Pd
2.2
Ag
1.9
Cd
1.7
In
1.7
Sn
1.8
Sb
1.9
Te
2.1
I
2.5
Xe
Cs
0.7
Ba
0.9
La
1.1
Hf
1.3
Ta
1.5
W
1.7
7
Re
1.9
Os
2.2
Ir
2.2
Pt
2.2
Au
2.4
Hg
1.9
Tl
1.8
Pb
1.8
Bi
1.9
Po
2.0
At
2.2
Rn
Do not learn these values!
Group 0 do not have values as they do not form compounds.
The value of the electronegativity is related to the size of the
atom.
Generally, as the atoms become smaller, the nucleus will tend
to attract an electron pair more strongly.
Atomic radius
It is actually impossible to measure atomic radius. Why?
What is used instead of an actual the radius of an individual
atom?
Going down a group the atomic radius increases.
This is because the outer electron is in an energy level that is
progressively further from the nucleus.
e.g Li: 2, 1
Na: 2, 8, 1
K: 2, 8, 8, 1 etc.
The atomic radius decreases across a period.
As the number of outer electrons increases so does the number
of protons in the nucleus. This increase in charge on the
nucleus increases the attraction to the outer energy level. This
results in the outer energy level becoming closer to the
nucleus.
Ionic Radius
When positive ions are formed,
the radius becomes smaller.
There is one fewer electrons
than protons so the nucleus
attracts the remaining electrons
more strongly and there is one
fewer energy level.
When negative ions are formed
there are more electrons in the
outer shell so more electron –
electron repulsion. As the
number of protons remains the
same, each electron is less
strongly attracted.
For both negative and positive ions the size of the ion
increases down the as the outer energy level is progressively
further from the nucleus.
Across a period: For both positive and negative ions we need
to compare isoelectronic ions (same number of electrons)
All have 10
electrons
Number of protons increases from Na+ (10 p) to Al3+ (13 p).
So as we move across the period the ions become smaller as
there is a stronger attraction for outer energy level by the
increasing number of protons.
All of these ions have 18 electrons but the number of
protons increases from P3- (15 p) to Cl- (17 p). So the ions
decrease in size as there is an increase in the attraction
between the outer energy level and the increasing nuclear
charge.
Melting Points
Down a group:
Element Melting point / oC
Element Melting point / oC
Li
181
F2
-220
Na
98
Cl2
-101
K
64
Br2
-72
Rb
39
I2
114
Cs
29
There is a decrease in melting
point because as the atoms get
larger, the forces of attraction
between them decrease. Metallic
bonding increase with valence
There is an increase in melting
point as the molecules get
bigger down the group due to
increased van der Waals forces
(see later).
Across a period:
3000
2500
mp / bp
2000
1500
1000
500
0
11
12
13
14
15
16
17
atomic number
mp/K
bp/K
18
Linked to the bond strength and structure of the elements.
Na  Al
These elements are metals. The melting and boiling
points increase from sodium to aluminium.
This is because the atoms are smaller and have an
increasing nuclear charge so there is a stronger attraction
for the delocalised electrons. Therefore, the strength of
the metal-metal bonds increase.
Alternatively: Increase in charge of metal ion  increase
in attraction for delocalised electrons.
Silicon has the highest m.p. and b.p. as it is a macromolecule
with a diamond-like structure. Strong covalent bonds link all
of the atoms in 3 dimensions. A large amount of energy is
required to break these bonds.
Phosphorus, sulphur and chlorine are all molecular substances.
The m.p.of each is determined by the strengths of van der
Waals forces (which in turn depends upon size of the
molecule). Each has a low m.p. as the van der Waals forces
are weak.
n.b P4 but S8 and then Cl2!
Argon has the lowest m.p and b.p as it exists as single atoms.
It has few electrons so and van der Waals forces are very
weak.