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After this lesson you should be able to:
•Recall the definitions for 1st and successive ionization energies.
•Understand the factors affecting the sizes of ionization energies.
•Show how ionization energies provide evidence for shells
Starter
At the very end of the last lesson we endeavoured to come up with a
set of principles to
help explain
ways inalways
which electrons
fill orbitals.
I
•Aufbau/building
principle:
electrons
fill the lowest
energy
am nowfirst.
going to tell three simples rules to help you understand the
orbitals
process.
Electron
arrangement
orbitals
•Hund's rule:
electronsinnever
pair up in the same orbital until all
There
rulesenergy
which are
determine
the way inand
which
electrons
orbitalsare
of three
the same
singly occupied,
all unpaired
fill
the orbitals
electrons
have parallel spin.
•Pauli exclusion principle: only two electrons may occupy the
same orbital, and they must do so with opposite spin.
You have 5 minutes to come up with a quick way to remember
these rules, any mnemonic will do but I would prefer it if you
would try to construct it as a rap/ poem/ song etc.
Here is my bad attempt- surely yours has to be better-surely!!
His name was Aufbau, for sure he had a principle
He told us how the orbitals would always be filled
Electrons he said- well they really have a thirst,
To fill the lowest energy orbitals,
and they wanna fill them first!
So you get the idea! Try it yourselves, best one gets homework
off, lamest one gets homework off!
Now,and
time
to get serious!
Ions
ionization
energy
Ions:
An atom can either lose or gain an electron to form an ion: a charged
atom. A positive ion is formed when an atom loses electron(s). For
example a lithium atom forms a positive ion with a 1+ charge by losing
an electron
Li
3p+,4n,3e-
Li
+
+
3p+, 4n, 2e-
e
_
A negative ion is formed when an atom gains electrons. For
example, an oxygen atom forms a negative ion with a 2- charge by
gaining two electrons.
Write what the electron configurations would look like - you have 2
minutes!
O + 2e-
6p+, 6n, 6e-
O
2-
6p+, 6n, 8e-
Ionization energy
Ionization energy measures the ease with which electrons are lost
in the formation of positive ions. An element has as many energies
as there are electrons, therefore:
The first ionization energy of an element is the energy required to
remove one electron from each atom in 1 mole of free gaseous atoms
to form 1 mole of gaseous 1+ ions.
The equation representing the first ionisation energy of sodium would be shown as:
Na(g)
Na+(g) + e-
First ionization energy = +496kj mol-1
Factors affecting ionization energies
Electrons are held in their shells by attraction from the nucleus, the 1st
electron lost will be from the highest occupied energy level. This
electron experiences least attraction from the nucleus.
There are three factors that affect the size of the attraction
Nuclear
•Atomic charge:
radius:
Electron
shielding
(screening):
The
the
of
protonsthe
in the
nucleus,
the
The greater
greater
the number
distance
between
nucleus
and the
the greater
outer electrons,
The
outer
shell
electrons
areAttraction
repelled by
anyrapidly
inner shells
between
attractive
force.
the less
the
attractive
force.
falls
with increasing
the
electrons
thelike
nucleus.
This repelling
effect known
as electron
distance
(a lotand
most
long distance
relationships)!
The distance
is a
shielding,
reduces
theand
overall
force experienced by the
very important
factor
has attractive
a big effect.
outer electrons.
1st task:
The three factors we just looked at are very important and may well
help to explain many, many chemical ideas throughout the course. They
are a basic grounding in the process of chemistry.
In pairs you will explain in simple terminology the three factors affecting
ionization energy (from your head, you may not use the books). Your
partner must then create a pamphlet on the three factors, while your
partner is doing this you must design a power point presentation on the
same topic. At the end of the lesson, you will swap these resourcesthey will get yours and you will get theirs. – you have 20 mins to
complete your task.
The second ionization energy of an element is the energy required
remove 1 electron from each atom in 1 mole of gaseous 1+ ions to form
1 mole of gaseous 2+ atoms
The equation representing the second ionization energy of sodium would be shown
as:
Na+ (g)
Notice thendpattern!
Na2+ (g) + e-
2 ionization energy = +4563kj mol-1
It makes sense then that Sodium with 11 electrons, has 11 successive
ionization energies. Successive ionization energies provide evidence for
the different energy levels of the principal quantum numbers
Taking this a step further, we can now look at the 1st ionization energies
of the first 20 elements in the periodic table :
first ionisation energy (kJ
per mole)
Variation of first ionisation energy with atomic
number for the first twenty elements
2500
2000
1500
1000
500
0
0
5
10
atomic number
15
20
There are various trends in this graph which can be explained by
reference to the proton number and electronic configuration of the
various elements. A number of factors must be considered:
•nuclear charge
•shielding
•effective nuclear charge
•electron repulsion
•See handout for further clarification:
The trends in first ionization energies amongst elements in the periodic
table can be explained on the basis of variations in one of the four
factors mentioned.
Trend across period 1:
Trends across period 2
Compare the first ionization energies of H and He. Neither have inner
In
general,
the first
ionization
energy
increases
across
a period
because
shells,
so
there
is
no
shielding.
He
has
two
protons
in
the
nucleus;
2
Compare now the first ionization energies of He (1s ) and Li (1s22s1H). Li
the
nuclear
charge
increases
but theelectrons
shieldingare
remains
the
same.
only
has
one.
Therefore
the
helium
more
strongly
has an extra proton in the nucleus (3) but two inner-shell electrons.
attracted
to the nucleus
andcancel
henceout
more
These inner-shell
electrons
thedifficult
charge to
of remove.
two of the protons,
The
first ionization
energy
of He
is thusonhigher
that of
reducing
the effective
nuclear
charge
the 2sthan
electron
to H.
+1. This is
Since
andthe
He effective
are the only
atoms
whose
electrons
are not
lowerHthan
nuclear
charge
onouter
the He
1s electrons,
+2, and
shielded
from theare
nucleus,
it follows
has the
highest first
so the electrons
less strongly
heldthat
andHe
easier
to remove.
ionization energy of all the elements. All elements (except H) have outer
electrons
which areenergy
shielded
from that
the nucleus
The first ionization
of to
Li issome
thusextent
lower than
of He. and thus
are easier to remove.
So Helium has the highest first ionization energy of all the elements.
Trends down a group:
Successive ionization energies
The second ionization energy of an atom is the energy required to
On descending
a group,
nuclear
the
remove
one electron
from the
eacheffective
of a mole
of free charge
gaseousstays
unipositive
same but the number of inner shells increases. The repulsion
ions.
2+(g) +inner
between
these
M+(g)  M
e shells and the outer electrons makes them
less stable,
pushes
themcan
further
from the
nucleus
and makes
Other
ionization
energies
be defined
in the
same way:
themthird
easier
to remove.
The
ionization
energy of an atom is the energy required to
remove one electron from each of a mole of bipositive ions.
M2+(g)  M3+(g) + e
The nth ionization energy can be defined as the energy required for
the process
M(n-1)+(g)  Mn+(g) + e
It always becomes progressively more difficult to remove successive
electrons from an atom; the second ionization energy is always
greater than the first, the third always greater than the second and
so on.
Variations in 1st ionization energies across a period in the periodic
table provides evidence for the existence of sub-shells. There are
trends we should notice with a little more academic training.
For example we might spot that there are some points where we
would notice a sharp decrease in 1st ionization energy between the
end of one period and the start of the next
(He
Li);
(Ne
Na);
(Ar
K).
What might this indicate? Homework off for the right answer..
This reflects the addition of a new outer shell with the resulting increase
in shielding and distance
We can use
trends such
as differences
between
ionisation
Variation
of ionisation
energy
with number
of
energy values to indicate
directly the
electronic configuration of
ionisations
in silicon
an atom
1000000
ionisation energy
For example: lets look at Silicon
100000
10000
1000
100
1
3
5
7
9
11
13
number of ionisations
Large jumps occur between 4th and 5th and between 12th and 13th.
Therefore there are three shells: The first contains 2 electrons, the second 8 and the
third 4.
Finally
•The successive ionization energies of an atom always increase. The
more electrons that are removed, the fewer the number electrons that
remain. There is therefore less repulsion between the electrons in the
resulting ion. The electrons are therefore more stable and harder to
remove.
•By far the largest jumps between successive ionization energies come
when the electron is removed from an inner shell. This causes a large
drop in shielding, a large increase in effective nuclear charge and a large
increase in ionization energy
Last task:
I am about to show you a chart of Ionisation energies, you will be given
this chart on a handout (going around now). I want you to explain what
is happening on the chart and continue to explain what you have
learned about ionisation energy values in this lesson.
Wrap up:
Today was a long information heavy lesson, in your own words let me
know what you learned today. Was it information overload or how did
you cope?