PERIODIC TRENDS
Periodicity
 Elements in the PT are arranged in order of
increasing atomic number.
 Elements in the same group - same chemical
and physical properties.
 Across the period - repeating pattern of
physical and chemical properties known as
periodicity.
Periodic Trends
Properties such as
 Melting points
 Electronegativity
 First ionisation energy
show periodicity
Ionisation Energy
The first ionisation energy is the energy required to
remove the valence electron from 1 mole of
gaseous atom to form a positive ion.
The equation for 1st IE of element M
M(g)  M+(g) + e-
Example
H(g)  H+(g) + e- ΔH = +1310kJmol-1
The amout of energy required to carry out this
process for 1 mole of hydrogen atoms is
1310kilojoules.
Electronegativity
 The electronegativity is the ability of an atom in
a covalent bond to attract shared pairs electrons
to itself.
 The greater the electronegativity of an atom, the
greater its ability to attract shared pairs of
electrons to itself.
 Electronegativity value is based on the Pauling
scale. A value of 4.0 is give to F (most
electronegative atom). The least electronegative
elements, Ce and Fr both have a value of 0.7
Atomic Radii
(Source: http://www.chem1.com/acad/webtext/atoms/atpt-6.html)
 Metallic radius is half the distance between nuclei in a metallic
crystal.
 Covalent radius is half the distance between like atoms that are
bonded together in a molecule.
 van der Waals’ radius is the effective radius of adjacent atoms
which are not chemically bonded in a solid, but are presumably in
"contact". An example would be the distance between the iodine
atoms of adjacent I2 molecules in crystalline iodine.
Atomic & ionic radii down a
group
 Atomic radii is determined by 2 opposing factors
- Shielding effect by the electrons of the inner
shell(s)
- Nuclear charge (due to protons)
 Moving down the group, both the nuclear charge
and shielding effect increase. However, the outer
electrons enter new shells. So, although the
nucleus gains protons, the electrons are not
further away, but also more effectively screened
by an addtional shell of electrons.
Atomic radius
increases down
the group
Ionic radii for ions of
the same charge also
increases down a
group for the same
reason.
Trends in first ionisation
energy
 The atomic radius increases down the group
as additional electrons are added, causing the
shielding effect to increase.
 The further the outer shell is from the
nucleus, the smaller the attractive force
exerted by the protons in the nucleus.
 More easily an outer electron can be
removed, the lower the ionisation energy.
The first IE decreases down the group
Trends in electronegativity
 There is an increasing distance between the
nucleus and electrons down the group. Hence,
the attractive force is decreased.
 Although the nuclear charge increases down the
group, this is counteracted by the increased
shielding effect due to additional electron shells.
The electronegativity decreases down the group
Trends in melting point
Group I
 Metals are held together by
metallic bonding.
 The strength of metallic bonding decreases because the
attractive forces between the delocalised electrons and the
nuclues decreases owing to the increase in the distance. The
increase in the nuclear charge is counteracted by the
increase in shielding.
Group 7
 As the molecules become large, the attractive forces
between them increases with the number of electrons in
atoms or molecules.
The melting point decreases down group I
The melting point increases down group 7
Atomic radii across a period
 Across the period from left to right, the no. of
protons and electrons increase by one.
 The electrons are added to the same shell,
hence, there is only a slight increase in the
shielding effect across the period.
 At the same time, additional protons are added
to the nucleus, increasing the nuclear charge.
 The effect of increasing nuclear charge more
than outweighs the small increase and
consequently all the electrons are pulled closer
to the nucleus.
The atomic radius decreases across the period
Ionic radii across a period
 The radii of positive ions decrease from Na+ to Al 3+
 The radii of positive ions decrease from P3- (phosphide ion) to Cl  The ionic radii increase from the Al 3+ to P3- .
Isoelectronic Species
 Isoelectronic species are atoms and ions that
have the same number of electrons
Species
Na+
Mg2+ Al3+
Species
P3-
S2-
Cl-
Nuclear charge
+11
+12
+13
Nuclear charge
+15
+16
+17
Number of electrons 10
10
10
Number of electrons 18
18
18
Ionic configuration
65
45
Ionic configuration
190
181
98
212
 The large increase in size from Al3+ to P3- is due to
the presence of additional electron shell which
causes a large increase in the shielding effect
resullting in an increase in the ionic radius.
First ionisation across a
period
 When moving across the period from left to
right, the nuclear charge increases but the
shielding effect only increases slightly (since
electrons enter the same shell).
 Hence, the electron shells are pulled
progressively closer to the nucleus.
The electronegativity increases across the
period
Comparing the eletronegativity
values
Summary of Periodic Trends
Chemical Properties
Group I alkali metals
 Li, Na and K contain 1 valence electron.
 Reactive metals, stored under liquid paraffin
to prevent them from reacting with air.
 Readily lose their valence electron -good
reducing agent
 Reactivity increases down the group
 React with water to form an alkali solution of the
metal hydroxide and hydrogen gas.
(i) 2Li(s) + 2H2O(l)  LiOH(aq) + H2(g)
Lithium floats and reacts quietly
(ii) 2Na(s) + 2H2O(l)  NaOH(aq) + H2(g)
Sodium melts into a ball which darts around on the
surface
(iii) 2Ks) + 2H2O(l)  KOH(aq) + H2(g)
Heat generated from the reaction with potassium
ignites the hydrogen.
 React readily with chlorine, bromine and iodine
to form ionic salts, e.g.
(i) 2Na(s) + Cl2(g)  2NaCl(s)
(ii) 2K(s) + Br2(l)  2KBr(s)
(iii) 2Ks) + I2(g)  2LiI(s)
Chemical Properties
 Chlorine is a stronger oxidizing agent than
bromine, so can remove the electron from
bromide ions in solution to form chloride ions
and bromine.
 Similarly, both chlorine and bromine can oxidize
iodide ions to form iodine.
(i) Cl2(aq) + 2Br-(aq)  2Cl-(aq) + Br2(aq)
(ii) Cl2(aq) + 2I-(aq)  2Cl-(aq) + I2(aq)
(iii) Br2(aq) + 2I-(aq)  2Br-(aq) + I2(aq)
Test for halide ions
 The presence of halide ions in solution can be
detected by adding silver nitrate solution.
Ag+(aq ) + X- (aq)  AgX(s)
X = Cl,Br or I
AgCl white
AgBr cream
AgI yellow
light
Ag(s) + ½ X2