Transcript fhs as 1.5 lesson5 cro trends in the periodic table

5: Trends in the periodic table
j.
represent data, in a graphical form, for elements 1 to 36 and use this to explain
the meaning of the term ‘periodic property’
k. explain trends in the following properties of the element from periods 2 and 3 of
the periodic table:
i. melting temperature of the elements based on given data using the structure
and the bonding between the atoms or molecules of the element
ii. ionization energy based on given data or recall of the shape of the plots of
ionization energy versus atomic number using ideas of electronic structure
and the way that electron energy levels vary across the period.
Connector – Explain what is meant by
• nuclear charge
• atomic or ionic radius
• screening effects
• ionisation energy
• electron affinity
• nuclear charge
• atomic or ionic radius
• screening effects
The number of protons present in the
nucleus of an atom or ion
Half the distance between the nuclei of two
touching atoms.
This is the reduced attraction between the
nucleus and the outer energy level electrons
when there are electrons in energy levels
between them.
• ionisation energy
is the amount of energy required to remove a
mole of electrons from a mole of gaseous
atoms or ions.
e.g.
• electron affinity
Na (g) → Na+ (g) + e -
is the amount of energy released when a mole
of gaseous atoms each gain an electron.
e.g.
General Periodic Trends
We need to consider and explain how the following factors change,
either across a period or down a group.
• Atomic radius
• Ionic radius
• Ionisation energy
• Electron affinity
General Periodic Trends – Across a period
1. How does the nuclear charge change across a period?
2. What effect does this have on the electrons in the outer energy
level?
3. What is the overall effect on atomic radius across the period?
Atomic radius – Across a Period
• As you move along a period, the nuclear charge
becomes increasingly positive as the number of
protons in the nucleus increase.
• Although the number of electrons also increases, the
outer electrons are all in the same energy level.
• This means that electrons are attracted more
strongly to the nucleus, thus reducing the atomic
radius across a period.
General Periodic Trends – Down a group
1. How does the arrangement of the electrons in an atom
change going down a group?
2. What happens to the nuclear charge?
3. What are the effects of 1 & 2 on the electrons in the outer
energy level?
4. What is the overall effect on atomic radius down a group?
Atomic radius – Down a Group
• As you go down a group, the outer electron(s) enter
into a new energy level.
• Although, the nucleus also gains positive
protons(charge), the electrons are both further away
and screened by more electron energy level.
• As a result, electrons are not held so tightly, and the
atomic radii increases.
Atomic Radius: the radius of an atom in picometers
1
2
13
14
15
16
17
18
Atomic radius vs. atomic number
Atomic Radius (pm)
250
K
200
Na
Li
150
Mg
Al Si
Be
100
Ca
P S Cl
B C N
O F
Ar
Ne
50
H
0
0
He
2
4
6
8
10
12
Element
14
16
18
20
Which is Bigger?
• Na or K ?
• Na or Mg ?
• Al or S ?
Ionic radius
Li,152 pm
3e and 3p
Does the size go
+
up or down when
Li + , 60 pm
losing
an
electron
to
2e and 3 p
form a cation?
Ionic radius
+
Li,152 pm
3e and 3p
Li + , 78 pm
2e and 3 p
Forming a cation.
• CATIONS are SMALLER than the atoms
from which they come.
• The electron/proton attraction has gone
UP and so size DECREASES.
Ionic radius
Does the size go up or
down when gaining
an electron to form an
anion?
Ionic radius
F, 71 pm
9e and 9p
F- , 133 pm
10 e and 9 p
Forming an
anion.
• ANIONS are LARGER than the atoms
from which they come.
• The electron/proton attraction has gone
DOWN and so size INCREASES.
• Trends in ion sizes are the same as
atom sizes.
Trends in Ion Sizes
Which is Bigger?
•
•
•
•
Cl or Cl- ?
K+ or K ?
K+ or Ca+2 ?
I- or Br- ?
Periodicity
• The repeating patterns or trends are
known as the periodicity of the elements
and the properties are known as periodic
properties.
Ionization Energy
It is the amount of
energy required to
remove a mole of
electrons from a mole of
gaseous atoms or ions.
Ionization Energy
• Ionisation energy is affected by three main
factors:
• The attraction between the nucleus and
the outermost electron
• The size of the positive nuclear charge
• Screening effect
Ionization energy vs. atomic number
He
Ionization energy (kJ/mol)
2500
Ne
2000
Ar
F
1500
N
H
Cl
C
Be
1000
O
P S
B
500
Mg Si
Al
Li
Ca
Na
K
0
0
2
4
6
8
10
12
Element
14
16
18
20
st
1
Which has a higher
ionization energy?
• Mg or Ca ?
• Al or S ?
• Cs or Ba ?
On the graph, the end of
each period is marked by a
peak of a high I.E. of a
noble gas
This distinctively high I.E.,
as we have seen, is due to
the fact that noble gases
have a complete outer shell
(the most stable electronic
configuration
This graph shows that the first
I.E. doesn’t increase smoothly.
This is due to the presence of
subshells within shells.
In the period Li to Ne, Be has a
higher I.E. than B.
Same when comparing Mg to
Al.
This is because, removing an
electron from B or Al, would
mean taking it from a 2p
subshell.
However, taking one from Be or
Mg, means removing from a full
s subshell. These are stable, so
more energy required.
Also, N and P also have peaks
in the graph.
This is because, they have half
filled p subshells.
Again underlying how the most
stable configuration is when
you have a full or half-full shell
or subshell.
Its graphs like this that prove
the existence of subshells.
ELECTRON AFFINITY
A few elements GAIN electrons
to form anions.
1) Electron affinity is the energy
change which occurs when an
electron is accepted by an atom
in the gaseous state.
A(g) + e A-(g)
a) Ionization Energy is always endothermic. It
always takes energy to remove an electron.
Change in Enthalpy (∆ H) is always +ve.
b) Electron Affinity can be either endothermic or
exothermic depending on the element.
c) An exothermic (Change in Enthalpy is -ve)
value for the electron affinity indicates that
energy is released upon the addition of an
electron to a gaseous atom.
d) The greater the negative value of the electron
affinity, the greater the tendency of an atom to
accept an electron.
e) A +∆H indicates that energy must be absorbed for an
atom to gain an electron.
f) As we go from left to right on the periodic table, the
elements have, in general, an increasing tendency to
form negative ions.
Electron Affinity of Oxygen
O atom [He] 
 

+ electron
O - ion [He] 
 
EA = - 141 kJ

∆E is EXOthermic
because O has
an affinity for an
e-.
Electron Affinity of Nitrogen
N atom [He] 
 

+ electron
N- ion
[He] 

EA = 0 kJ


∆E is zero for Ndue to electronelectron
repulsions.
Spot the trend of Electron affinity in a group
and in a period
Electron Affinity Trends
(Same as for Ionization Energy)
• Group Trends: Electron affinity decreases
from top to bottom on the periodic table
• Outermost electrons are further away from
the nucleus and therefore easier to remove
• Shielding effect increases down the group
• Period Trends: increases from left to right
• Nuclear charge is increasing with no
increase in shielding effect
• Outermost electrons are closer to the
nucleus
Trends in Electron Affinity
Electron affinity
MORE STABLE
MORE STABLE
LESS STABLE
Electron affinity
Melting Temperature
Melting Temp. (for a substance) =
tamp at which pure solid is in
equilibrium with pure liquid at
atmospheric pressure.
This is determined by the packing
and bonding of atoms in a
substance.
Melting Temperature
Metals like Li, Mg, and Al have
peaks due to strong metallic
bonding (ions packed in sea of
electrons)
Si and C are giant covalent
structures with extremely strong
covalent bonding. So at peaks.
After peaks, the elements tend to
exist as simple covalent molecules
(diatomic etc). Therefore, despite
strong intramolecular forces
(covalent bond) have weak
intermolecular attraction. So easy
to pull individual molecules apart.
Electronegativity
Electronegativity is a measure of the tendency of an atom to attract
a bonding pair of electrons.
The trend across Period 3 looks like this:
Explaining the trend
As you go across the period, the bonding electrons are always in the
same energy level. They are always being screened by the same inner
electrons. All that differs is the number of protons in the nucleus. As you
go from sodium to chlorine, the number of protons steadily increases and
so attracts the bonding pair more closely.
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