Transcript Slide 1

Trends in the Periodic
Table
(Chpt. 7)
The three properties of elements whose changes
across the periodic table are to be investigated are:
1. Atomic radius (size)
2. Ionization energy
3. Electronegativity
1. Trends in Atomic Radii
The atomic radius (covalent radius) of an atom is
defined as half the distance between the nuclei of
two atoms of the same element that are joined
together by a single covalent bond
e.g. in a molecule of hydrogen it is found that the
distance between the two nuclei is 0.074 nm.
Therefore the covalent radius of a hydrogen atom
is 0.074/2 = 0.037nm
*Note: Noble gases do not form covalent bonds with one
another so they have NO atomic radius
*Note – understanding the trends in atomic radii values:
The size of an atom depends on the attraction between
the positively charged protons and negatively charged
electrons in the atom:
- Large attraction – positive protons will pull
outer electrons closer to the nucleus - leads to
smaller atomic radius
- Small attraction – electrons will be found
further from the nucleus – leads to a larger
atomic radius
Screening (shielding) Effect:
means that the inner shell or
shells of electrons help to
shield the outer electrons from
the positive charge in the
nucleus
(Please leave space for diagram)
Increase
Atomic radius
increases down
the group:
- new shell
- screening
effect
Atomic radius decreases across the period:
- increasing nuclear charge
- no increase in screening effect
Decrease
1. Trends in Atomic radii:
A) Across a period
1 2
3
4
radius decreases
5 6 7 8
Increasing effective nuclear charge: - number of
protons increases from left to right across a period
therefore greater attraction between nucleus and
outer electrons – shells drawn closer to nucleus
No increase in screening effect: - same number of
shells therefore no increase in screening
B) Down a group
n=1
n=2
n=3
n=4
radius increases
Although there is an increase in the
number of protons:
New Shell – additional electrons are
going into a new shell which is
further from the nucleus – radius
increases
Screening Effect – inner electrons
screen outer electrons from the
nucleus
Which is Bigger???
• Na or K ?
• Na or Mg ?
• Al or I ?
2. Trends in Ionisation Energy
•
Ionisation energy is a term used to describe
the tendency of an atom to lose an electron
The First Ionisation Energy of an atom is the minimum
amount of energy required to completely remove the
most loosely bound electron from a neutral gaseous
atom in its ground state
•
1st ionisation energy equation for hydrogen
and sodium
H(g) - eH+(g)
Na(g) - eNa+(g)
• The second ionisation energy of an element refers
to the removal of a second electron from the
positive ion formed when the first electron is
removed e.g. second ionisation energy of sodium :
Na+(g) - e-
Na2+(g)
• Ionisation energy unit – ‘kilojoules per mole’
• Table of first ionisation energies given on pg. 80 in
the log tables
Ionisation Energy increases across the period:
- increasing atomic charge
- decreasing atomic radius
Increase
Ionisation Energy
Decrease
decreases down
the group:
- increasing
atomic radius
- screening
effect
A) Down a group
ionisation energy decreases
Although there is an increase in the number of
protons (nuclear charge):
Increasing atomic radius – radius increases
therefore number of shells of electrons increases
– outermost electrons are moving further away
from attractive force of nucleus and it becomes
easier to remove an electron from the outer shell.
Screening Effect – inner electrons screen outer
electrons from the positively charged nucleus –
becomes easier to remove outermost electrons
and ionisation energy values decrease.
B) Across a period
ionisation energy increases
Increasing effective nuclear charge – number of
protons in nucleus is increasing as move from left
to right across a period. As a result, the attraction
between the nucleus and the outer electrons is
increasing. More energy is now required to remove
an electron from the outermost shell
Decreasing atomic radius – atomic radius decreases
from left to right so outer electrons drawn closer to
nucleus. Due to increased attraction between
electron in outermost shell and nucleus the
ionisation energy values increase.
Which has a higher 1st ionization energy?
Mg or Ca ?
Al or S ?
Cs or Ba ?
Exceptions to the General Trend Across a Period
If plot a graph of ionisation energy Vs atomic
number for first 20 elements it is clear that, in
any one period, some elements do not follow the
smooth increase
First Ionisation Energy
Ionisation Energies for n = 2 period
Ne
F
N
C
Be
Li
O
B
Atomic Number
Ionisation Energies for n = 3 period
First Ionisation Energy
Ar
Cl
P
Mg
Na
Si
S
Al
Atomic Number
• In n = 2 period beryllium and nitrogen have higher
values than expected
• In the n = 3 period magnesium and phosphorous
have higher values than expected
• This irregularity can be explained by the fact that
any sublevel that is completely filled (Be, Mg) or
exactly half filled (N, P) has extra stability
• Because of this extra stability their ionisation
energy values are higher
Be: 1s2 , 2s2
Mg: 1s2 , 2s2 , 2p6 , 3s2
N: 1s2 , 2s2 , 2p3
P: 1s2 , 2s2 , 2p6 , 3s2 , 3p3
*Note: If asked to account fully for
trend across 2nd /3rd period write out
electron configurations of Be and N/
Mg and P
Second and Subsequent Ionisation Energies
Evidence for the fact that electrons are arranged in
shells of different energies is also provided by studying
the values of a number of ionisation energies of any
one particular element.
Second ionisation higher than the first because
removing an electron from:
- an ion so there is more positive charges per
electron i.e. greater nuclear charge
- closer to nucleus as atomic radius of ion is
smaller than corresponding atom
Successive Ionisation Energies
- Bigger increase if an electron is removed
from a half filled sublevel
- Much bigger if new shell entered because
electron being removed is:
a) closer to the nucleus
b) in a full sublevel
c) has less shielding
Investigation of Successive Ionisation energies of
Aluminium
• First 3 ionisation energies increase steadily as
electrons are removed from the 3rd shell.
• 4th is a big jump as second shell is entered
• 5th to 11th get steadily bigger as successive
electrons are removed from the second shell
• 12th very big as 1st shell is entered
Note: Jumps in ionisation energies are evidence for
the existence of the energy levels (shells)
3. Trends in Electronegativity
Electronegativity is the relative attraction that an
atom in a molecule has for the shared pair of
electrons in a covalent bond
Concept proposed by
Linus Pauling
1901-1994
Decrease
Electronegativity increases across the period:
- increasing nuclear charge
- decreasing atomic radius
Increase
Electronegativity
decreases down
the group:
- increasing
atomic radius
- screening
effect
A) Down a group
electronegativity decreases
Even though nuclear charge increases down a
group:
Increasing atomic radius: - atomic radius
increases – outermost electrons are moving
further away from attractive force of nucleus.
Therefore smaller attraction between the nucleus
and the shared pair of electrons
Screening Effect: - inner electrons screen outer
electrons from the positively charged nucleus –
since it is outermost electrons involved in
bonding, the attraction of the nucleus for these
electrons decreases going down the group i.e.
electronegativity decreases.
B) Across a period
electronegativity increases
Increasing effective nuclear charge – number of
protons in nucleus is increasing as move from left
to right across a period. As a result, the attraction
between the nucleus and the outer electrons is
increasing. Therefore the electrons involved in
bonding are being more strongly attracted to the
nucleus i.e. electronegativity increases
Decreasing atomic radius – atomic radius decreases
from left to right so outer electrons drawn closer to
nucleus. Due to increased attraction between
electron in outermost shell and nucleus the
electronegativity values increase.
Trends Within Groups
The chemical properties of elements are largely
determined by the number of electrons in the
outermost shell:
All elements in group 1 have one electron in
outermost shell – all have similar chemical
properties
All elements in group 7 have seven electrons in
their outermost shell – all have similar
chemical properties
1. Trends in chemical reactivity of Alkali Metals (Group 1)
• Very reactive elements
• Low ionisation energies and electronegativity values
tend to loose electrons to form ionic compounds.
• Reactivity of alkali metals increases down the group (as
ionisation energy decreases – more easily outer electron
is lost – the more reactive the metal)
a) Reaction of Alkali Metals with oxygen:
All alkali metals react with oxygen to form oxides:
Potassium + Oxygen
2K
+ ½O2
Potassium Oxide
K2 O
b) Reaction of Alkali Metals with water:
All alkali metals react with water to form the hydroxide
of the metal and hydrogen gas is given off:
sodium + water
sodium
hydroxide
Na
NaOH
+ H2O
+ hydrogen
+
½ H2
c) Reaction of Alkali metals with dilute acids:
Extremely dangerous reaction so much hydrogen
produced that an explosion occurs:
Sodium + hydrochloric
acid
sodium + hydrogen
chloride
Na
NaCl
+
HCl
+ ½ H2
2. Trends in chemical reactivity of Halogens (Group 7):
• elements in group 7 are the most electronegative
elements in periodic table.
• electronegativity values decrease down the group –
fluorine being the most electronegative element
• due to strong attraction for electrons – very reactive
elements as they tend to remove electrons from other
substances
• The reactivity of the halogens increases up the group
(as one moves up the group the electronegativity
values increase)
Halogens as oxidising agents:
• an oxidising agent is a substance that removes
electrons from other substances
e.g. If chlorine gas is bubbled through a solution of
bromide ions (Br-) the chlorine takes the electrons
from the bromide ions and converts them into
bromine:
chlorine + bromide ion
Cl
+ Br-
½
2
chloride ion + bromine
Cl+ Br
½
2
*Note: electronegativity of Cl greater than that of Br
A similar reaction occurs between bromine and iodide
ions (I-) where the bromine (more electronegative)
would take electrons from the iodide ions( less
electronegative).
In each case the more reactive halogen
displaces the less reactive halogen from
solution
3. Trends in Physical Properties
• Noble Gases = Inert gases = very little reactivity
Even though it is not possible to consider any trends
in chemical properties of the noble gases it is possible
to study trends in their physical properties
• Boiling Points of Noble Gases:
- steady increase in boiling points down the group
- caused by the increasing atomic radius down the
group
- as atoms increase in size temporary dipoles are
produced due to the larger electron clouds
- Van der Waals forces are stronger between bigger
atoms - giving rise to higher boiling points
• A similar trend can be noted in the halogens as you
go down the group