Ionization Energy

Table of Contents

1. Periodic Table and Trend of Ionization Energies

1. 1st, 2nd, and 3rd Ionization Energies
2. The Effects of Electron Shells on Ionization Energy
3. Ionization Energy and Electron Affinity--Similar Trend
4. Prediction of Covalent and Ionic Bonds
5. Questions
6. Answers
7. Outside Links
8. References

Ionization energy is the quantity of energy that an isolated, gaseous

atom in the ground electronic state must absorb to discharge an
electron, resulting in a cation.
H(g)→H + (g)+e −

This energy is usually expressed in kJ/mol, or the amount of

energy it takes for all the atoms in a mole to lose one electron each.
When considering an initially neutral atom, expelling the first
electron will require less energy than expelling the second, the
second will require less energy than the third, and so on. Each
successive electron requires more energy to be released. This is
because after the first electron is lost, the overall charge of the
atom becomes positive, and the negative forces of the electron will
be attracted to the positive charge of the newly formed ion. The
more electrons that are lost, the more positive this ion will be, the
harder it is to separate the electrons from the atom.
In general, the further away an electron is from the nucleus, the
easier it is for it to be expelled. In other words, ionization energy is
a function of atomic radius; the larger the radius, the smaller the
amount of energy required to remove the electron from the outer
most orbital.
For example, it would be far easier to take electrons away from the
larger element of Ca (Calcium) than it would be from one where
the electrons are held tighter to the nucleus, like Cl (Chlorine).
In a chemical reaction, understanding ionization energy is
important in order to understand the behavior of whether various
atoms make covalent or ionic bonds with each other. For instance,
the ionization energy of Sodium (alkali metal) is 496KJ/mol (1)
whereas Chlorine's first ionization energy is 1251.1 KJ/mol (2).
Due to this difference in their ionization energy, when they
chemically combine they make an ionic bond. Elements that
reside close to each other in the periodic table or elements that do
not have much of a difference in ionization energy make polar
covalent or covalent bonds. For example, carbon and oxygen make
CO2 (Carbon dioxide) reside close to each other on a periodic
table; they, therefore, form a covalent bond. Carbon and chlorine
make CCl4 (Carbon tetrachloride) another molecule that is
covalently bonded.

Periodic Table and Trend of Ionization Energies

As described above, ionization energies are dependent upon the
atomic radius. Since going from right to left on the periodic
table, the atomic radius increases, and the ionization energy
increases from left to right in the periods and up the groups.
Alkali metals (IA group) have small ionization energies, especially
when compared to halogens or VII A group (see diagram 1). In
addition to the radius (distance between nucleus and the electrons
in outermost orbital), the number of electrons between the nucleus
and the electron(s) you're looking at in the outermost shell have an
effect on the ionization energy as well. This effect, where the full
positive charge of the nucleus is not felt by outer electrons due to
the negative charges of inner electrons partially canceling out the
positive charge, is called shielding. The more electrons shielding
the outer electron shell from the nucleus, the less energy required
to expel an electron from said atom. The higher the shielding effect
the lower the ionization energy (see diagram 2). It is because of
the shielding effect that the ionization energy decreases from top to
bottom within a group. From this trend, Cesium is said to have the
lowest ionization energy and Fluorine is said to have the highest
ionization energy (with the exception of Helium and Neon).

Table 1: showing the increasing trend of ionization energy in

KJ/mol (exception in case of Boron) from left to right in the
periodic table(8)
Li Be B C N O F
520 899 800 1086 1402 1314 1680

Table 2: showing decreasing trend

of ionization energies (Kj/mol) from
top to bottom (Cs is the exception
 in the first group) (8)
Li 520

Na 496

K 419

Rb 408

Cs 376

Fr 398

1st, 2nd, and 3rd Ionization Energies

The symbol I 1 stands for the first ionization energy (energy

required to take away an electron from a neutral atom) and the
symbol I 2 stands for the second ionization energy (energy
required to take away an electron from an atom with a +1 charge.
Each succeeding ionization energy is larger than the preceding
energy. This means that I 1 <I 2 <I 3 <...<I n will always be true.
Example of how ionization energy increases as succeeding
electrons are taken away.
Mg(g)→Mg + (g)+e − I 1 =738kJ/mol

Mg + (g)→Mg 2+ (g)+e − I 2 =1451kJ/mol

See first, second, and third ionization energies of elements/ions in

the table below
 Table 3: Ionization Energies (kJ/mol)
1 2 3 4 5 6 7 8

H 1312

He 2372 5250

Li 520 7297 11810

Be 899 1757 14845 21000

B 800 2426 3659 25020 32820

C 1086 2352 4619 6221 37820 47260

N 1402 2855 4576 7473 9442 53250 64340

O 1314 3388 5296 7467 10987 13320 71320 84070

F 1680 3375 6045 8408 11020 15160 17860 92010

Ne 2080 3963 6130 9361 12180 15240

Na 496 4563 6913 9541 13350 16600 20113 25666

Mg 737 1450 7731 10545 13627 17995 21700 25662

The Effects of Electron Shells on Ionization Energy

Electron orbitals are separated into various shells which have
strong impacts on the ionization energies of the various electrons.
For instance, let us look at aluminum. Aluminum is the first
element of its period with electrons in the 3p shell. This makes the
first ionization energy comparably low to the other elements in the
same period, because it only has to get rid of one electron to make
a stable 3s shell, the new valence electron shell. However, once
you've moved past the first ionization energy into the second
ionization energy, there is a large jump in the amount of energy
required to expel another electron. This is because you now are
trying to take an electron from a fairly stable and full 3s electron
shell. Electron shells are also responsible for the shielding that was
explained above.
Ionization Energy and Electron Affinity--Similar Trend
Both ionization energy and electron affinity have similar trend in
the periodic table. For example, just as ionization energy increases
along the periods, electron affinity also increases. Likewise,
electron affinity decreases from top to bottom due to the same
factor, i.e., shielding effect. Halogens can capture an electron
easily as compared to elements in the first and second group. This
tendency to capture an electron in a gaseous state is termed as
electronegativity. This tendency also determines one of the
chemical differences between Non metallic and metallic elements.
Diagram 3: showing increasing trend of electron affinity from left
to right (9).

B 27 C 123.4 N -7 O 142.5 F 331.4

Diagram 4:showing decreasing pattern of electron affinities of

elements from top to bottom (9)

H 73.5
Li 60.4

Na
53.2

K
48.9

Rb
47.4

Cs 46.0

Fr
44.5

As indicated above, the elements to the right side of periodic table

(diagram 3) have tendency to receive the electron while the one at
the left are more electropositive. Also, from left to right, the
metallic characteristics of elements decrease (4).
Prediction of Covalent and Ionic Bonds
The difference of electronegativity or ionization energies between
two reacting elements determine the fate of the type of bond. For
example, there is a big difference of ionization energies and
electronegativity between Na and. Cl. Therefore, sodium
completely removes the electron from its outermost orbital and
chlorine completely accepts the electron, and as a result we have
an ionic bond (4). However, in cases where there is no
difference in electronegativity, the sharing of electrons produces a
covalent bond. For example, electronegativity of Hydrogen is 2.1
and the combination of two Hydrogen atoms will definitely make a
covalent bond (by sharing of electrons). The combination of
Hydrogen and Fluorine (electronegativity=3.96) will produce a
polar covalent bond because they have small differences between
electronegativity (5).

IONISATION ENERGY

This page explains what first ionisation energy is, and then looks at the
way it varies around the Periodic Table - across periods and down
groups. It assumes that you know about simple atomic orbitals, and
can write electronic structures for simple atoms. You will find a link at
the bottom of the page to a similar description of successive ionisation
energies (second, third and so on).

Important! If you aren't reasonable happy about atomic orbitals and

electronic structures you should follow these links before you go any further.

Defining first ionisation energy

Definition

The first ionisation energy is the energy required to remove the most
loosely held electron from one mole of gaseous atoms to produce 1
mole of gaseous ions each with a charge of 1+.

This is more easily seen in symbol terms.

It is the energy needed to carry out this change per mole of X.

 Worried about moles? Don't be! For now, just take it as a measure of a
particular amount of a substance. It isn't worth worrying about at the moment.

Things to notice about the equation

The state symbols - (g) - are essential. When you are talking about
ionisation energies, everything must be present in the gas state.

Ionisation energies are measured in kJ mol-1 (kilojoules per mole).

They vary in size from 381 (which you would consider very low) up to
2370 (which is very high).

All elements have a first ionisation energy - even atoms which don't
form positive ions in test tubes. The reason that helium (1st I.E. = 2370
kJ mol-1) doesn't normally form a positive ion is because of the huge
amount of energy that would be needed to remove one of its electrons.

Patterns of first ionisation energies in the Periodic Table

The first 20 elements

First ionisation energy shows periodicity. That means that it varies in

a repetitive way as you move through the Periodic Table. For example,
look at the pattern from Li to Ne, and then compare it with the identical
pattern from Na to Ar.

These variations in first ionisation energy can all be explained in terms

of the structures of the atoms involved.
Factors affecting the size of ionisation energy

Ionisation energy is a measure of the energy needed to pull a

particular electron away from the attraction of the nucleus. A high
value of ionisation energy shows a high attraction between the electron
and the nucleus.

The size of that attraction will be governed by:

The charge on the nucleus.

The more protons there are in the nucleus, the more positively charged
the nucleus is, and the more strongly electrons are attracted to it.

The distance of the electron from the nucleus.

Attraction falls off very rapidly with distance. An electron close to the
nucleus will be much more strongly attracted than one further away.

The number of electrons between the outer electrons and the

nucleus.

Consider a sodium atom, with the electronic structure 2,8,1. (There's

no reason why you can't use this notation if it's useful!)

If the outer electron looks in towards the nucleus, it doesn't see the
nucleus sharply. Between it and the nucleus there are the two layers of
electrons in the first and second levels. The 11 protons in the sodium's
nucleus have their effect cut down by the 10 inner electrons. The outer
electron therefore only feels a net pull of approximately 1+ from the
centre. This lessening of the pull of the nucleus by inner electrons is
known as screening or shielding.

Warning! Electrons don't, of course, "look in" towards the nucleus - and they
don't "see" anything either! But there's no reason why you can't imagine it in
these terms if it helps you to visualise what's happening. Just don't use these
terms in an exam! You may get an examiner who is upset by this sort of
loose language.
Whether the electron is on its own in an orbital or paired with
another electron.

Two electrons in the same orbital experience a bit of repulsion from

each other. This offsets the attraction of the nucleus, so that paired
electrons are removed rather more easily than you might expect.

Explaining the pattern in the first few elements

Hydrogen has an electronic structure of 1s1. It is a very small atom,

and the single electron is close to the nucleus and therefore strongly
attracted. There are no electrons screening it from the nucleus and so
the ionisation energy is high (1310 kJ mol-1).

Helium has a structure 1s2. The electron is being removed from the
same orbital as in hydrogen's case. It is close to the nucleus and
unscreened. The value of the ionisation energy (2370 kJ mol-1) is much
higher than hydrogen, because the nucleus now has 2 protons
attracting the electrons instead of 1.

Lithium is 1s22s1. Its outer electron is in the second energy level,

much more distant from the nucleus. You might argue that that would
be offset by the additional proton in the nucleus, but the electron
doesn't feel the full pull of the nucleus - it is screened by the 1s2
electrons.
You can think of the electron as feeling a net 1+ pull from the centre (3
protons offset by the two 1s2 electrons).

If you compare lithium with hydrogen (instead of with helium), the

hydrogen's electron also feels a 1+ pull from the nucleus, but the
distance is much greater with lithium. Lithium's first ionisation energy
drops to 519 kJ mol-1 whereas hydrogen's is 1310 kJ mol-1.

The patterns in periods 2 and 3

Talking through the next 17 atoms one at a time would take ages. We
can do it much more neatly by explaining the main trends in these
periods, and then accounting for the exceptions to these trends.

The first thing to realise is that the patterns in the two periods are
identical - the difference being that the ionisation energies in period 3
are all lower than those in period 2.
Explaining the general trend across periods 2 and 3

The general trend is for ionisation energies to increase across a

period.

In the whole of period 2, the outer electrons are in 2-level orbitals - 2s

or 2p. These are all the same sort of distances from the nucleus, and
are screened by the same 1s2 electrons.

The major difference is the increasing number of protons in the

nucleus as you go from lithium to neon. That causes greater attraction
between the nucleus and the electrons and so increases the ionisation
energies. In fact the increasing nuclear charge also drags the outer
electrons in closer to the nucleus. That increases ionisation energies
still more as you go across the period.

Note: Factors affecting atomic radius are covered on a separate page.

In period 3, the trend is exactly the same. This time, all the electrons
being removed are in the third level and are screened by the 1s22s22p6
electrons. They all have the same sort of environment, but there is an
increasing nuclear charge.

Why the drop between groups 2 and 3 (Be-B and Mg-Al)?

The explanation lies with the structures of boron and aluminium. The
outer electron is removed more easily from these atoms than the
general trend in their period would suggest.

Be 1s22s2 1st I.E. = 900 kJ mol-1

B 1s22s22px1 1st I.E. = 799 kJ mol-1

You might expect the boron value to be more than the beryllium value
because of the extra proton. Offsetting that is the fact that boron's
outer electron is in a 2p orbital rather than a 2s. 2p orbitals have a
slightly higher energy than the 2s orbital, and the electron is, on
average, to be found further from the nucleus. This has two effects.

 The increased distance results in a reduced attraction and so a

reduced ionisation energy.
 The 2p orbital is screened not only by the 1s2 electrons but, to
some extent, by the 2s2 electrons as well. That also reduces the
pull from the nucleus and so lowers the ionisation energy.

The explanation for the drop between magnesium and aluminium is

the same, except that everything is happening at the 3-level rather
than the 2-level.

Mg 1s22s22p63s2 1st I.E. = 736 kJ mol-1

Al 1s22s22p63s23px1 1st I.E. = 577 kJ mol-1

The 3p electron in aluminium is slightly more distant from the nucleus

than the 3s, and partially screened by the 3s2 electrons as well as the
inner electrons. Both of these factors offset the effect of the extra
proton.

Warning! You might possibly come across a text book which describes the
drop between group 2 and group 3 by saying that a full s2 orbital is in some
way especially stable and that makes the electron more difficult to remove. In
other words, that the fluctuation is because the group 2 value for ionisation
energy is abnormally high. This is quite simply wrong! The reason for the
fluctuation is because the group 3 value is lower than you might expect for
the reasons we've looked at.
Why the drop between groups 5 and 6 (N-O and P-S)?

Once again, you might expect the ionisation energy of the group 6
element to be higher than that of group 5 because of the extra proton.
What is offsetting it this time?

N 1s22s22px12py12pz1 1st I.E. = 1400 kJ mol-1

O 1s22s22px22py12pz1 1st I.E. = 1310 kJ mol-1

The screening is identical (from the 1s2 and, to some extent, from the
2s2 electrons), and the electron is being removed from an identical
orbital.

The difference is that in the oxygen case the electron being

removed is one of the 2px2 pair. The repulsion between the two
electrons in the same orbital means that the electron is easier to
remove than it would otherwise be.

The drop in ionisation energy at sulphur is accounted for in the

same way.

Trends in ionisation energy down a group

As you go down a group in the Periodic Table ionisation energies

generally fall. You have already seen evidence of this in the fact
that the ionisation energies in period 3 are all less than those in
period 2.

Taking Group 1 as a typical example:

Why is the sodium value less than that of lithium?

There are 11 protons in a sodium atom but only 3 in a lithium

atom, so the nuclear charge is much greater. You might have
expected a much larger ionisation energy in sodium, but
offsetting the nuclear charge is a greater distance from the
nucleus and more screening.

Li 1s22s1 1st I.E. = 519 kJ mol-1

Na 1s22s22p63s1 1st I.E. = 494 kJ mol-1

Lithium's outer electron is in the second level, and only has the
1s2 electrons to screen it. The 2s1 electron feels the pull of 3
protons screened by 2 electrons - a net pull from the centre of 1+.

The sodium's outer electron is in the third level, and is screened

from the 11 protons in the nucleus by a total of 10 inner electrons.
The 3s1 electron also feels a net pull of 1+ from the centre of the
atom. In other words, the effect of the extra protons is
compensated for by the effect of the extra screening electrons.
The only factor left is the extra distance between the outer
electron and the nucleus in sodium's case. That lowers the
ionisation energy.

Similar explanations hold as you go down the rest of this group -

or, indeed, any other group.

Trends in ionisation energy in a transition series

Apart from zinc at the end, the other ionisation energies are all
much the same.

All of these elements have an electronic structure [Ar]3dn4s2 (or

4s1 in the cases of chromium and copper). The electron being
lost always comes from the 4s orbital.

Note: The 4s orbital has a higher energy than the 3d in the transition
elements. That means that it is a 4s electron which is lost from the atom
when it forms an ion. It also means that the 3d orbitals are slightly closer
to the nucleus than the 4s - and so offer some screening.

Confusingly, this is inconsistent with what we say when we use the Aufbau
Principle to work out the electronic structures of atoms.

I have discussed this in detail in the page about the order of filling 3d and
4s orbitals.

If you are a teacher or a very confident student then you might like to
follow this link.

If you aren't so confident, or are coming at this for the first time, I suggest
that you ignore it. Remember that the Aufbau Principle (which uses the
assumption that the 3d orbitals fill after the 4s) is just a useful way of
working out the structures of atoms, but that in real transition metal atoms
the 4s is actually the outer, higher energy orbital.

As you go from one atom to the next in the series, the number of
protons in the nucleus increases, but so also does the number of 3d
electrons. The 3d electrons have some screening effect, and the extra
proton and the extra 3d electron more or less cancel each other out as
far as attraction from the centre of the atom is concerned.
The rise at zinc is easy to explain.

Cu [Ar]3d104s1 1st I.E. = 745 kJ mol-1

Zn [Ar]3d104s2 1st I.E. = 908 kJ mol-1

In each case, the electron is coming from the same orbital, with
identical screening, but the zinc has one extra proton in the nucleus
and so the attraction is greater. There will be a degree of repulsion
between the paired up electrons in the 4s orbital, but in this case it
obviously isn't enough to outweigh the effect of the extra proton.

Note: This is actually very similar to the increase from, say, sodium to
magnesium in the third period. In that case, the outer electronic structure
is going from 3s1 to 3s2. Despite the pairing-up of the electrons, the
ionisation energy increases because of the extra proton in the nucleus.
The repulsion between the 3s electrons obviously isn't enough to outweigh
this either.

I don't know why the repulsion between the paired electrons matters less
for electrons in s orbitals than in p orbitals (I don't even know whether you
can make that generalisation!). I suspect that it has to do with orbital
shape and possibly the greater penetration of s electrons towards the
nucleus, but I haven't been able to find any reference to this anywhere. In
fact, I haven't been able to find anyone who even mentions repulsion in
the context of paired s electrons!

If you have any hard information on this, could you contact me via the
address on the about this site page.

Ionisation energies and reactivity

The lower the ionisation energy, the more easily this change happens:

You can explain the increase in reactivity of the Group 1 metals (Li,
 Na, K, Rb, Cs) as you go down the group in terms of the fall in
ionisation energy. Whatever these metals react with, they have to form
positive ions in the process, and so the lower the ionisation energy, the
more easily those ions will form.

The danger with this approach is that the formation of the positive ion
is only one stage in a multi-step process.

For example, you wouldn't be starting with gaseous atoms; nor would
you end up with gaseous positive ions - you would end up with ions in
a solid or in solution. The energy changes in these processes also vary
from element to element. Ideally you need to consider the whole
picture and not just one small part of it.

However, the ionisation energies of the elements are going to be major

contributing factors towards the activation energy of the reactions.
Remember that activation energy is the minimum energy needed
before a reaction will take place. The lower the activation energy, the
faster the reaction will be - irrespective of what the overall energy
changes in the reaction are.

The fall in ionisation energy as you go down a group will lead to lower
activation energies and therefore faster reactions.

Note: You will find a page discussing this in more detail in the inorganic
section of this site dealing with the reactions of Group 2 metals with water.

Factors affecting the ionization energy.

Factors affecting the ionization energy are:

1. Atomic size: In small atoms electrons remains closer to nucleus and they
feel more nuclear attraction. So more ionization energy is required to
remove electron from small atoms. While in big sized atoms, valence
electrons are away from nucleus , so they experience less nuclear
attraction. Hence it requires less ionization energy to remove electrons from
big atoms.
Therefore, Ionization energy is inversely proportional to Atomic size.

2. Nuclear charge: By increasing the nuclear charge electrons feel more

nuclear attraction. Hence more ionization energy is required.

Therefore, Ionization Energy is directly proportional to Nuclear charge.

3. Penetration Power: Tendency of becoming nearer to the nucleus is called

penetration power.
The order of penetration power of different sub-shells - s > p > d > f.

Therefore, Ionization energy is Directly proportional to Penetration

Power.

4. Stability: In stable configuration we require more energy to release the

electron as compared to non stable configuration.
Therefore, Ionization energy is directly proportional to Stability.

Ionization Energy is more of full-filled shell as compared to half-filled shell.

5. Screening & Shielding effect: Presence of other orbits between nucleus

and last orbit decreases the nuclear attraction. This effect is called
screening effect but electron-electron repulsion is called shielding effect
which also decreases the nuclear attraction. Due to presence of these
effects ionisation energy decreases.
------------------

Downwards in a group ionization energy decreases due to increase in size

of atoms while in a period from left to right ionization energy increases due
to increase in nuclear charge.
------------------------------------------------------------------------
-------------------------------------------------------------------
Ionization Energy
In long form of the periodic table, elements are arranged in 9
periods and 18 groups. In a group, elements have the same number
of valence electrons in their outermost shell. Valence electrons
refer to those electrons of an atom that can form chemical bonds
with other atoms. The arrangement of electrons across the shells is
known as the electronic structure of the element.

In the electronic structure, the number of shells increases steadily

as we compare elements of one group. Or we can say that the
number of shells increases down the group. Now observe the
electronic structure of the elements in the second period. What do
you observe? The elements have different valence electrons but
same number of shells, that is, 2. Notice that the atomic number of
these elements increase by 1 unit, from left to right, the number of
valence shell electrons also increases by 1 unit.

The number of elements in each period is based on how electrons

are filled in various shells. For example,
K Shell – 2 × (1)2 = 2, hence the first period has 2 elements.
L Shell – 2 × (2)2 = 8, hence the second period has 8 elements and
so on.

From all these observations, we can conclude trend of various

properties of elements such as ionisation potential, electron
affinity, electron negativity etc. Let’s discuss the ionisation
potential of elements in the periodic table.
Ionization Energy Definition
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The electrons are raised to higher energy levels by absorption of energy

from an external source. If this process is continued, a stage comes when
the electron goes completely out of the influence of the nucleus and a
positive ion is produced. The electron in an atom is attracted by the
positively charged nucleus.
In order to remove an electron from an atom, energy has to be supplied
to it to overcome the attractive force. This energy is referred to as
ionization energy or ionization potential.
"The amount of energy required to remove the most loosely bound
electron or the outermost electron from an isolated gaseous atom of an
element in its lowest energy state or the ground state to produce a
cation is known as ionization potential or ionization energy of that
element."

A(g) + Energy → A+ (g) + electron

It is generally represented as I or IP and is measured in electron volts (eV)
or kilo calories (K calories) per gram of atom.

1eV atom = 23.06 kcal / mole = 96.3 kj/mole

Ionization potential is a very important property which gives an idea

about the tendency of an atom to form a gaseous positive ion. The
smaller the value of ionization energy, the easier it is to remove the
electron from the atom.
The process by which the element loses an electron, to convert itself into
a cation is called its ionization. This process is an endothermic process,
since the energy is supplied to effect it.

Ionization Energy Trend

Back to Top

In a group, as we move from top to bottom, ionization potential

decreases. As we move down the group, with the increase in atomic
number, new shells ass on and the size of the atom gradually increases.
The distance between the nucleus and the outermost shell increases and
the screening effect also increases.

As a result, the attractive force of the nucleus on the outermost electron

decreases. Thereby, the amount of energy required to remove an
electron from the atom decreases. Ionization potential, thus, decreases
down a group.

Element kJ/mol eV/atom

H 1312.0 13.595

Li 520.1 5.390

Na 459.2 5.138

K 418.7 4.339

Rb 403.0 4.176
 Cs 375.7 3.893

The element with highest ionization potential in the periodic table is

Helium (2372.1 kj/mol), while the element with the lowest ionization
potential is Caesium (375.7 kj/mol).

First Ionization Energy

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First ionization energy of an atom is the energy required to remove the

first electron from the outermost shell of an atom.

Once the first electron has been removed from the gaseous atom, it is
possible to remove second and successive electrons from positive ions
one after the other.

Ionization Energy Periodic Table

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In a period, as we move from left to right, the ionization energy

increases.
As we move along a period, the nuclear charge increases and the
electrons are added into the same shell. Thereby the effective nuclear
charge increases and size decreases. Therefore, the energy required to
remove an electron increases.

Elements Li Be B C N O F Ne
 Ionization Potential 520 900 802 1085 1398 1315 1680 2080

However, some irregularities in the general trend have been noticed.

These are due to half - filled and completely filled configurations which
have extra stability.

To illustrate this, let us consider the ionization energies beryllium and

boron. The ionization energy of boron is slightly less than that of
beryllium even though boron has a greater nuclear charge. This can be
understood by comparing the electronic configurations of Be (1s2, 2s2 )
and B (1s2, 2s2, sp1).

In the case of Beryllium, the electron removed during ionization is an s-

electron whereas the electron removed in the case of boron is p-electron.
The penetration of a 2s electron to the nucleus is more than that of a 2p
electron, hence the 2p - electron of boron is more shielded than the 2s
electron of Beryllium. Therefore, ionization energy of Beryllium is more
than that of Boron.

Second Ionization Energy

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Once the first ionization electron is removed from the gaseous atom, it is
possible to remove second and successive electrons from positive ions
one after the other. Removal of a second electron from an already
ionized gaseous atom is called as second ionization energy.
 A(g) + IE1 → A+(g) + e- [First Ionization]
A+(g) + IE1 → A2+(g) + e- [Second ionization]

The amount of energies required to remove most loosely bound electron

from unipositive, dipositive, tripositive .... ions of the element in gaseous
state are called second, third, fourth, etc ionization energies respectively.

The second, third and fourth, etc ionization energies are collectively
called as successive ionization energies.

It is also seen that IE3 > IE2 > IE1.

Ionization Energy Chart

Back to Top

Ionization chart of II A group elements are as follows.

Element Number of inner shells eV/atom

Be 1 9.3

Ma 2 7.6

Ca 3 6.1

Sr 4 5.7

Ba 5 5.2
Factors Affecting Ionization Potential
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The magnitude of Ionization potential depends upon:

1. Charge in the nucleus

The greater the charge on the nucleus, the more difficult it will be to
remove an outermost electron from that atom. Thus, greater the nuclear
charge of an atom, greater will be its ionization potential. Halogens have
highest ionization potential in their respective periods.

2. Atomic radius

With the increase in the atomic radius, the ionization potential decreases.
This is because of the fact that in case of larger atoms the attraction
between the nucleus and the outer most electron is less and hence it is
easier to remove an electron from a larger atom than from a smaller one.

This can be seen from the fact that as we move down a group, the size
increases, while ionization potential decreases.

3. Completely and partially filled in shells

The atoms having a completely filled or partially filled outer most shell
are comparatively more stable than the atoms with incompletely filled
outer shells. This, the ionization potential of such atoms are relatively
higher than expected normally from their positions in the periodic table.
Successive Ionization Energies
Back to Top

The amount of energy necessary for removing the subsequent electrons

of a gaseous atom is better known as successive ionization energies.
These are basically termed as 1st , 2nd, 3rd or 4th …… ionization energy
and these depends completely on the removal of the 1st, 2nd, 3rd
electrons respectively.

M (g) → IE 1 M + (g) +e −

M + (g) → IE 2 M 2+ (g) +e −

M 2+ (g) → IE 3 M 3+ (g) +e −

 2nd ionization energy amount are higher compared to 1st because

after the first electron is removed the atom changes into a positive
monovalent ion.
 The nuclear charge remains same although there is a decrease in
electron number and this ultimately leads to a situation where the
remaining electrons getting held more tightly by the nucleus which
ultimately makes it very difficult to remove the second electron.
 The value of the 2nd ionization energy (IE2) is more than the 1st
ionization energy (IE1). Removing 2nd electron would ultimately
lead to forming a di-positive ion which results in a stronger
attraction between the nucleus and remaining electrons. Higher
value of ionization energy results from this.
Atomic and Ionic Radius
Table of Contents

1. Atomic Radius
2. Trends in atomic radius in the Periodic Table
1. Trends in atomic radius in Periods 2 and 3
2. Trends in atomic radius down a group
3. Trends in atomic radius across periods
3. Trends in the transition elements
4. Ionic Radius
1. Trends in ionic radius in the Periodic Table
2. Trends in ionic radius across a period
5. Trends in ionic radius for some more isoelectronic ions
6. Contributors
This page explains the various measures of atomic radius, and then looks at the way it varies
around the Periodic Table - across periods and down groups. It assumes that you understand
electronic structures for simple atoms written in s, p, d notation.

Atomic Radius
Unlike a ball, an atom does not have a fixed radius. The radius of an atom can only be found
by measuring the distance between the nuclei of two touching atoms, and then halving that
distance.

As you can see from the diagrams, the same atom could be found to have a different radius
depending on what was around it. The left hand diagram shows bonded atoms. The atoms
are pulled closely together and so the measured radius is less than if they are just touching.
This is what you would get if you had metal atoms in a metallic structure, or atoms covalently
bonded to each other. The type of atomic radius being measured here is called the metallic
radius or the covalent radius depending on the bonding.
The right hand diagram shows what happens if the atoms are just touching. The attractive
forces are much less, and the atoms are essentially "unsquashed". This measure of atomic
radius is called the van der Waals radius after the weak attractions present in this situation.

Trends in atomic radius in the Periodic Table

The exact pattern you get depends on which measure of atomic radius you use - but the
trends are still valid. The following diagram uses metallic radii for metallic elements, covalent
radii for elements that form covalent bonds, and van der Waals radii for those (like the noble
gases) which don't form bonds.

Trends in atomic radius in Periods 2 and 3

Trends in atomic radius down a group

It is fairly obvious that the atoms get bigger as you go down groups. The reason is equally
obvious - you are adding extra layers of electrons.

Trends in atomic radius across periods

You have to ignore the noble gas at the end of each period. Because neon and argon don't
form bonds, you can only measure their van der Waals radius - a case where the atom is
pretty well "unsquashed". All the other atoms are being measured where their atomic radius is
being lessened by strong attractions. You aren't comparing like with like if you include the
noble gases.

Leaving the noble gases out, atoms get smaller as you go across a period. If you think about
it, the metallic or covalent radius is going to be a measure of the distance from the nucleus to
the electrons which make up the bond. (Look back to the left-hand side of the first diagram on
this page if you aren't sure, and picture the bonding electrons as being half way between the
two nuclei.)

From lithium to fluorine, those electrons are all in the 2-level, being screened by the 1s2
electrons. The increasing number of protons in the nucleus as you go across the period pulls
the electrons in more tightly. The amount of screening is constant for all of these elements.
In the period from sodium to chlorine, the same thing happens. The size of the atom is
controlled by the 3-level bonding electrons being pulled closer to the nucleus by increasing
numbers of protons - in each case, screened by the 1- and 2-level electrons.

Trends in the transition elements

Although there is a slight contraction at the beginning of the series, the atoms are all much
the same size. The size is determined by the 4s electrons. The pull of the increasing number
of protons in the nucleus is more or less offset by the extra screening due to the increasing
number of 3d electrons.

Ionic Radius
Ionic radii are difficult to measure with any degree of certainty, and vary according to the
environment of the ion. For example, it matters what the co-ordination of the ion is (how many
oppositely charged ions are touching it), and what those ions are. There are several different
measures of ionic radii in use, and these all differ from each other by varying amounts. It
means that if you are going to make reliable comparisons using ionic radii, they have to come
from the same source.

What you have to remember is that there are quite big uncertainties in the use of ionic radii,
and that trying to explain things in fine detail is made difficult by those uncertainties. What
follows will be adequate for UK A level (and its various equivalents), but detailed explanations
are too complicated for this level.

Trends in ionic radius in the Periodic Table

Trends in ionic radius down a group: This is the easy bit! As you add extra layers of electrons
as you go down a group, the ions are bound to get bigger. The two tables below show this
effect in Groups 1 and 7.

electronic ionic radius

structure (nm)
 of ion

Li+ 2 0.076

Na+ 2, 8 0.102

K+ 2, 8, 8 0.138

Rb+ 2, 8, 18, 8 0.152

Cs+ 2, 8, 18, 18, 8 0.167

electronic
ionic radius
structure
(nm)
of ion

F- 2, 8 0.133

Cl- 2, 8, 8 0.181

Br- 2, 8, 18, 8 0.196

I- 2, 8, 18, 18, 8 0.220

Trends in ionic radius across a period

Let's look at the radii of the simple ions formed by elements as you go across Period 3 of the
Periodic Table - the elements from Na to Cl.

Na+ Mg2+ Al3+ P3- S2- Cl-

 no of protons 11 12 13 15 16 17

electronic structure of ion 2,8 2,8 2,8 2,8,8 2,8,8 2,8,8

ionic radius (nm) 0.102 0.072 0.054 (0.212) 0.184 0.181

The table misses out silicon which doesn't form a simple ion. The phosphide ion radius is in
brackets because it comes from a different data source, and I am not sure whether it is safe
to compare it. The values for the oxide and chloride ions agree in the different source, so it is
probably OK. The values are again for 6-co-ordination, although I can't guarantee that for the
phosphide figure.

First of all, notice the big jump in ionic radius as soon as you get into the negative ions. Is this
surprising? Not at all - you have just added a whole extra layer of electrons.

Notice that, within the series of positive ions, and the series of negative ions, that the ionic
radii fall as you go across the period. We need to look at the positive and negative ions
separately.

 The positive ions: In each case, the ions have exactly the same electronic structure -
they are said to be isoelectronic. However, the number of protons in the nucleus of the
ions is increasing. That will tend to pull the electrons more and more towards the centre
of the ion - causing the ionic radii to fall. That is pretty obvious!
 The negative ions: Exactly the same thing is happening here, except that you have an
extra layer of electrons. What needs commenting on, though is how similar in size the
sulphide ion and the chloride ion are. The additional proton here is making hardly any
difference.

The difference between the size of similar pairs of ions actually gets even smaller as you go
down Groups 6 and 7. For example, the Te2- ion is only 0.001 nm bigger than the I- ion.

As far as I am aware there is no simple explanation for this - certainly not one which can be
used at this level. This is a good illustration of what I said earlier - explaining things involving
ionic radii in detail is sometimes very difficult.

Trends in ionic radius for some more isoelectronic ions

This is only really a variation on what we have just been talking about, but fits negative and
positive isoelectronic ions into the same series of results. Remember that isoelectronic ions
all have exactly the same electron arrangement.
 N3- O2- F- Na+ Mg2+ Al3+

no of protons 7 8 9 11 12 13

electronic structure of ion 2, 8 2, 8 2, 8 2, 8 2, 8 2, 8

ionic radius (nm) (0.171) 0.140 0.133 0.102 0.072 0.054

Note: The nitride ion value is in brackets because it came from a different source, and I don't
know for certain whether it relates to the same 6-co-ordination as the rest of the ions. This
matters. My main source only gave a 4-coordinated value for the nitride ion, and that was
0.146 nm.

You might also be curious as to how the neutral neon atom fits into this sequence. Its van der
Waals radius is 0.154 or 0.160 nm (depending on which source you look the value up in) -
bigger than the fluoride ion. You can't really sensibly compare a van der Waals radius with the
radius of a bonded atom or ion.

Periodic Trends in Ionic Radii

An understanding of periodic trends is necessary when analyzing
and predicting molecular properties and interactions. Common
periodic trends include those in ionization energy, atomic radius,
and electron affinity. One such trend is closely linked to atomic
radii -- ionic radii. Neutral atoms tend to increase in size down a
group and across a period. When a neutral atom gains or loses an
electron, creating an anion or cation, the atom's radius increases or
decreases. This module explains how this occurs and how this
trend differs from that of atomic radii.
Shielding and Penetration
Electromagnetic interactions between electrons in an atom modify
the effective nuclear charge (Z eff ) on each electron. Penetration
refers to the presence of an electron inside the shell of an inner
electron, and shielding is the process by which an inner electron
masks an outer electron from the full attractive force of the
nucleus, decreasing Z eff . Differences in orbital characteristics
dictate differences in shielding and penetration. Within the same
energy level (indicated by the principle quantum number, n), due
to their relative proximity to the nucleus, s-orbital electrons both
penetrate and shield more effectively than p-orbital electrons, and
p electrons penetrate and shield more effectively than d-orbital
electrons. Shielding and penetration along with the effective
nuclear charge determine the size of an ion. An overly-simplistic
but useful conceptualization of effective nuclear charge is given by
the following equation:
Z eff =Z−S

where Z is the number of protons in the nucleus of an atom or ion

(the atomic number), and
 S is the number of core electrons.

Figure 1 illustrates how this equation can be used to estimate the

effective nuclear charge of
sodium:
 Figure 1: Shielding of the valence electrons by the core electrons

Cations and Anions

Neutral atoms that have lost an electron exhibit a positive charge
and are called cations. These cations are smaller than their
respective atoms; this is because when an electron is lost, electron-
electron repulsion (and therefore, shielding) decreases and the
protons are better able to pull the remaining electrons towards the
nucleus (in other words, Z eff increases). A second lost electron
further reduces the radius of the ion. For instance, the ionic radius
of Fe2+ is 76 pm, while that of Fe3+ is 65 pm. If creation of an ion
involves completely emptying an outer shell, then the decrease in
radius is especially great.
Neutral atoms that have gained an electron are called anions, and
they are much larger than their respective atoms. As an additional
electron occupies an outer orbital, there is increased electron-
electron repulsion (and hence, increased shielding) which pushes
the electrons further apart. Because the electrons now outnumber
the protons in the ion, the protons can not pull the extra electrons
as tightly toward the nucleus; this results in decreased Z eff
Figure 2 shows an isoelectric series of atoms and ions (each has the
same number of electrons, and thus the same degree of electron-
electron repulsion and shielding) with differing numbers of protons
(and thus different nuclear attraction), giving the relative ionic
sizes of each atom or ion.

Figure 2
The Periodic Trend
Due to each atom’s unique ability to lose or gain an electron,
periodic trends in ionic radii are not as ubiquitous as trends in
atomic radii across the periodic table. Therefore, trends must be
isolated to specific groups and considered for either cations or
anions.
Consider the s- and d-block elements. All metals can lose electrons
and form cations. The alkali and alkali earth metals (groups 1 and
2) form cations which increase in size down each group; atomic
radii behave the same way. Beginning in the d-block of the
periodic table, the ionic radii of the cations do not significantly
change across a period. However, the ionic radii do slightly
decrease until group 12, after which the trend continues (Shannon
1976). It is important to note that metals, not including groups 1
and 2, can have different ionic states, or oxidation states, (e.g. Fe2+
or Fe3+ for iron) so caution must be employed when generalizing
about trends in ionic radii across the periodic table.
All non-metals (except for the noble gases which do not form ions)
form anions which become larger down a group. For non-metals, a
subtle trend of decreasing ionic radii is found across a period
(Shannon 1976). Anions are almost always larger than cations,
although there are some exceptions (i.e. fluorides of some alkali
metals).

Figure 3
Measurement and Factors Affecting Ionic Radii
The ionic radius of an atom is measured by calculating its spatial
proportions in an ionic bond with another ion within a crystal
lattice. However, it is to consistently and accurately determine the
proportions of the ionic bonds. After comparing many compounds,
chemist Linus Pauling assign a radius of 140 pm to O2- and use this
as a reference point to determine the sizes of other Ionic Radii
(Jensen 2010).
Ionic radius is not a permanent trait of an ion, but changes
depending on coordination number, spin state, and other variables
(Shannon 1976). For a given ion, the ionic radius increases with
increasing coordination number and is larger in a high-spin state
than in a low-spin state.
According to group theory, the idea of ionic radii as a
measurement of spherical shapes only applies to ions that form
highly-symmetric crystal lattices like Na+ and Cl-. The point group
symmetry of a lattice determines whether or not the ionic radii in
that lattice can be accurately measured (Johnson 1973). For
instance, lattices with Oh and Td symmetries are considered to
have high symmetry; thus the electron densities of the component
ions occupy relatively-spherical regions and ionic radii can be
measured fairly accurately. However, for less symmetrical and
more polar lattices such as those with Cn, Cnh, and Cnv symmetries,
significant changes in the electron density can occur, causing
deviations from spherical shape; these deviations make ionic radii
more difficult to measure

Ionic Radius
Similarly charged ions tend to decrease in size
across a period (row) and increase in size down
a group (column).
Learning Objectives[ edit ]
 Identify the general trends of the ionic radius size for the periodic
table.

 Describe the change in radius of an atom upon ionization.

Key Points[ edit ]
 The ionic radius is the distance between the nucleus and the
electron in the outermost shell of an ion.

 When an atom loses an electron to form a cation, the lost

electron no longer contributes to shielding the other electrons
from the charge of the nucleus; consequently, the other
electrons are more strongly attracted to the nucleus, and the
radius of the atom gets smaller.

 When an electron is added to an atom, forming an anion, the

added electron repels other electrons, resulting in an increase
in the size of the atom.

 The trend observed in size of ionic radii is due to shielding of

the outermost electrons by the inner-shell electrons so that the
outer shell electrons do not "feel" the entire positive charge of
the nucleus.

Terms[ edit ]
 ion
An atom or group of atoms bearing an electrical charge, such as the
sodium and chlorine atoms in a salt solution.

 cation
A positively charged ion, as opposed to an anion.

 anion
A negatively charged ion, as opposed to a cation.
In chemistry, periodic trends are the tendencies of certain elemental
characteristics to increase or decrease along a period (row) or group
(column) of the periodic table of elements. Ionic radius (rion) is the radius
of an ion, regardless of whether it is an anion or a cation. Although
neither atoms nor ions have sharp boundaries, it is useful to treat them
as if they are hard spheres with radii. In this way, the sum of ionic radii of
a cation and an anion can give us the distance between the ions in a
crystal lattice. Ionic radii are typically given in units of either picometers
(pm) or Angstroms (Å), with 1 Å = 100 pm. Typical values range from 30
pm (0.3 Å) to over 200 pm (2 Å).

Trends in Ionic Radii

Ions may be larger or smaller than the neutral atom, depending on the
ion's charge. When an atom loses an electron to form a cation, the lost
electron no longer contributes to shielding the other electrons from the
charge of the nucleus; consequently, the other electrons are more
strongly attracted to the nucleus, and the radius of the atom gets smaller.
Similarly, when an electron is added to an atom, forming an anion, the
added electron repels other electrons, resulting in an increase in the size
of the atom.
The ionic radius is not a fixed property of a given ion; rather, it varies with
/chemistry/definition/coordinationcoordination number, spin state, and
other parameters. For our purposes, we are considering the ions to be as
close to their ground state as possible. Nevertheless, ionic radius values
are sufficiently transferable to allow periodic trends to be recognized.
ELECTRON AFFINITY

This page explains what electron affinity is, and then looks at the
factors that affect its size. It assumes that you know about simple
atomic orbitals, and can write electronic structures for simple atoms.

Important! If you aren't reasonable happy about atomic orbitals and

electronic structures you should follow these links before you go any
further.

First electron affinity

Ionisation energies are always concerned with the formation of
positive ions. Electron affinities are the negative ion equivalent, and
their use is almost always confined to elements in groups 6 and 7 of
the Periodic Table.

Defining first electron affinity

The first electron affinity is the energy released when 1 mole of

gaseous atoms each acquire an electron to form 1 mole of gaseous
1- ions.

This is more easily seen in symbol terms.

It is the energy released (per mole of X) when this change happens.

First electron affinities have negative values. For example, the first
electron affinity of chlorine is -349 kJ mol-1. By convention, the
negative sign shows a release of energy.

The first electron affinities of the group 7 elements

F -328 kJ mol-1
Cl -349 kJ mol-1
Br -324 kJ mol-1
I -295 kJ mol-1

Note: These values are based on the most recent research. If you are
using a different data source, you may have slightly different numbers.
That doesn't matter - the pattern will still be the same.

Is there a pattern?

Yes - as you go down the group, first electron affinities become less
(in the sense that less energy is evolved when the negative ions are
formed). Fluorine breaks that pattern, and will have to be accounted
for separately.

The electron affinity is a measure of the attraction between the

incoming electron and the nucleus - the stronger the attraction, the
more energy is released.

The factors which affect this attraction are exactly the same as
those relating to ionisation energies - nuclear charge, distance and
screening.

Note: If you haven't read about ionisation energy recently, it might be a

good idea to follow this link before you go on. These factors are
discussed in more detail on that page than they are on this one.
The increased nuclear charge as you go down the group is offset by
extra screening electrons. Each outer electron in effect feels a pull
of 7+ from the centre of the atom, irrespective of which element you
are talking about.

For example, a fluorine atom has an electronic structure of

1s22s22px22py22pz1. It has 9 protons in the nucleus.

The incoming electron enters the 2-level, and is screened from the
nucleus by the two 1s2 electrons. It therefore feels a net attraction
from the nucleus of 7+ (9 protons less the 2 screening electrons).

By contrast, chlorine has the electronic structure

1s22s22p63s23px23py23pz1. It has 17 protons in the nucleus.

But again the incoming electron feels a net attraction from the
nucleus of 7+ (17 protons less the 10 screening electrons in the first
and second levels).

Note: If you want to be fussy, there is also a small amount of screening

by the 2s electrons in fluorine and by the 3s electrons in chlorine. This
will be approximately the same in both these cases and so doesn't affect
the argument in any way (apart from complicating it!).

The over-riding factor is therefore the increased distance that the

incoming electron finds itself from the nucleus as you go down the
group. The greater the distance, the less the attraction and so the
less energy is released as electron affinity.

Note: Comparing fluorine and chlorine isn't ideal, because fluorine

breaks the trend in the group. However, comparing chlorine and bromine,
say, makes things seem more difficult because of the more complicated
electronic structures involved.

What we have said so far is perfectly true and applies to the fluorine-
chlorine case as much as to anything else in the group, but there's
another factor which operates as well which we haven't considered yet -
 and that over-rides the effect of distance in the case of fluorine.

Why is fluorine out of line?

The incoming electron is going to be closer to the nucleus in fluorine

than in any other of these elements, so you would expect a high
value of electron affinity.

However, because fluorine is such a small atom, you are putting the
new electron into a region of space already crowded with electrons
and there is a significant amount of repulsion. This repulsion
lessens the attraction the incoming electron feels and so lessens
the electron affinity.

A similar reversal of the expected trend happens between oxygen

and sulphur in Group 6. The first electron affinity of oxygen (-142 kJ
mol-1) is smaller than that of sulphur (-200 kJ mol-1) for exactly the
same reason that fluorine's is smaller than chlorine's.

Comparing Group 6 and Group 7 values

As you might have noticed, the first electron affinity of oxygen (-142
kJ mol-1) is less than that of fluorine (-328 kJ mol-1). Similarly
sulphur's (-200 kJ mol-1) is less than chlorine's (-349 kJ mol-1).
Why?

It's simply that the Group 6 element has 1 less proton in the nucleus
than its next door neighbour in Group 7. The amount of screening is
the same in both.

That means that the net pull from the nucleus is less in Group 6
than in Group 7, and so the electron affinities are less.
First electron affinity and reactivity

The reactivity of the elements in group 7 falls as you go down the

group - fluorine is the most reactive and iodine the least.

Often in their reactions these elements form their negative ions. At

GCSE the impression is sometimes given that the fall in reactivity is
because the incoming electron is held less strongly as you go down
the group and so the negative ion is less likely to form. That
explanation looks reasonable until you include fluorine!

An overall reaction will be made up of lots of different steps all

involving energy changes, and you cannot safely try to explain a
trend in terms of just one of those steps. Fluorine is much more
reactive than chlorine (despite the lower electron affinity) because
the energy released in other steps in its reactions more than makes
up for the lower amount of energy released as electron affinity.

Second electron affinity

You are only ever likely to meet this with respect to the group 6
elements oxygen and sulphur which both form 2- ions.

Defining second electron affinity

The second electron affinity is the energy required to add an

electron to each ion in 1 mole of gaseous 1- ions to produce 1 mole
of gaseous 2- ions.

This is more easily seen in symbol terms.

It is the energy needed to carry out this change per mole of X-.

Why is energy needed to do this?

You are forcing an electron into an already negative ion. It's not
 going to go in willingly!

1st EA = -142 kJ mol-1

2nd EA = +844 kJ mol-1

The positive sign shows that you have to put in energy to perform
this change. The second electron affinity of oxygen is particularly
high because the electron is being forced into a small, very
electron-dense space.

Electron Affinity
Table of Contents

1. Introduction
2. First Electron Affinity
3. Nonmetals vs. Metals
4. Patterns in Electron Affinity
5. Why is Fluorine an Anomaly?
6. Comparing Group 16 and Group 17 values
7. Second Electron Affinity
8. Practice Problems
9. Answers
10. References
11. ChemWiki Links
12. Outside Links
13. Contributors

Electron affinity is defined as the change in energy (in kJ/mole) of

a neutral atom (in the gaseous phase) when an electron is added to
the atom to form a negative ion. In other words, the neutral atom's
likelihood of gaining an electron.
Introduction
Energy of an atom is defined when the atom loses or gains energy
through chemical reactions that cause the loss or gain of electrons.
A chemical reaction that releases energy is called an exothermic
reaction and a chemical reaction that absorbs energy is called
an endothermic reaction. Energy from an exothermic reaction is
negative, thus energy is given a negative sign; whereas, energy
from an endothermic reaction is positive and energy is given a
positive sign. An example that demonstrates both processes is
when a person drops a book. When he or she lifts a book, he or she
gives potential energy to the book (energy absorbed). However,
once the he or she drops the book, the potential energy converts
itself to kinetic energy and comes in the form of sound once it hits
the ground (energy released).
When an electron is added to a neutral atom, i.e first electron
affinity, energy is released; thus, the first electron affinities are
negative. However, when an electron is added to a negative ion,
i.e. second electron affinity, more energy is required. Thus, more
energy is released to add the electron into an ion because the
negative ion has to force the electron to go into its electron orbital;
thus, the second electron affinities are positive.
First Electron Affinity (negative energy because energy released):
X (g) +e − →X − (g) (1)

Second Electron Affinity (positive energy because energy needed

is more than gained):
 X − (g) +e − →X 2− (g) (2)

First Electron Affinity

Ionization energies are always concerned with the formation of
positive ions. Electron affinities are the negative ion equivalent,
and their use is almost always confined to elements in groups 16
and 17 of the Periodic Table. The first electron affinity is the
energy released when 1 mole of gaseous atoms each acquire an
electron to form 1 mole of gaseous -1 ions. It is the energy released
(per mole of X) when this change happens. First electron affinities
have negative values. For example, the first electron affinity of
chlorine is -349 kJ mol-1. By convention, the negative sign shows a
release of energy.
When an electron is added to a metal element, energy is needed to
gain that electron (endothermic reaction). Metals have a less likely
chance to gain electrons because it is easier to lose their valance
electrons and form cations. It is easier to lose their valence
electrons because metals' nuclei do not have a strong pull on their
valence electrons. Thus, metals are known to have lower electron
affinities.
Example 1: Group 1 Electron Affinities
This trend of lower electron affinities for metals is described by the
Group 1 metals:
 Lithium (Li): 60 KJ mol-1
 Sodium (Na): 53 KJ mol-1
 Potassium (K): 48 KJ mol-1
 Rubidium (Rb): 47 KJ mol-1
 Cesium (Cs): 46 KJ mol-1
Notice that electron affinity decreases down the group.
When nonmetals gain electrons, the energy change is usually
negative because they give off energy to form an anion
(exothermic process); thus, the electron affinity will be negative.
Nonmetals have a greater electron affinity than metals because of
their atomic structures: first, nonmetals have more valence
electrons than metals do, thus it is easier for the nonmetals to gain
electrons to fulfill a stable octet and secondly, the valence electron
shell is closer to the nucleus, thus it is harder to remove an electron
and it easier to attract electrons from other elements (especially
metals). Thus, nonmetals have a higher electron affinity than
metals, meaning they are more likely to gain electrons than atoms
with a lower electron affinity.
electrons than atoms with a lower electron affinity.
Example 2: Group 17 Electron Affinities
For example, nonmetals like the elements in the halogens series in
Group 17 have a higher electron affinity than the metals. This trend
is described as below. Notice the negative sign for the electron
affinity which shows that energy is released.
 Fluorine (F) -328 kJ mol-1
 Chlorine (Cl) -349 kJ mol-1
 Bromine (Br) -324 kJ mol-1
 Iodine (I) -295 kJ mol-1

Notice that electron affinity decreases down the group, but

increases up with the period.
Nonmetals vs. Metals
To summarize the difference between the electron affinity of
metals and nonmetals:
 Metals: Metals like to lose valence electrons to form cations to have
a fully stable octet. They absorb energy (endothermic) to
lose electrons. The electron affinity of metals is lower than that of
nonmetals.
 Nonmetals: Nonmetals like to gain electrons to form anions to have
a fully stable octet. They release energy (exothermic) to gain
electrons to form an anion; thus, electron affinity of nonmetals is
higher than that of metals.

Patterns in Electron Affinity

Electron affinity increases upward for the groups and from left to
right across periods of a periodic table because the electrons added
to energy levels become closer to the nucleus, thus a stronger
attraction between the nucleus and its electrons. Remember that
greater the distance, the less of an attraction; thus, less energy is
released when an electron is added to the outside orbital. In
addition, the more valence electrons an element has, the more
likely it is to gain electrons to form a stable octet. The less valence
electrons an atom has, the least likely it will gain electrons.
Electron affinity decreases down the groups and from right to left
across the periods on the periodic table because the electrons are
placed in a higher energy level far from the nucleus, thus a
decrease from its pull. However, one might think that since the
number of valence electrons increase going down the group, the
element should be more stable and have higher electron affinity.
One fails to account for the shielding affect. As one goes down the
period, the shielding effect increases, thus repulsion occurs
between the electrons. This is why the attraction between the
electron and the nucleus decreases as one goes down the group in
the periodic table.
As you go down the group, first electron affinities become less (in
the sense that less energy is evolved when the negative ions are
formed). Fluorine breaks that pattern, and will have to be
accounted for separately. The electron affinity is a measure of the
attraction between the incoming electron and the nucleus - the
stronger the attraction, the more energy is released. The factors
which affect this attraction are exactly the same as those relating to
ionization energies - nuclear charge, distance and screening. The
increased nuclear charge as you go down the group is offset by
extra screening electrons. Each outer electron in effect feels a pull
of 7+ from the center of the atom, irrespective of which element
you are talking about.
Example 3: Fluorine vs. Chlorine
A fluorine atom has an electronic structure of 1s22s22px22py22pz1.
It has 9 protons in the nucleus.The incoming electron enters the 2-
level, and is screened from the nucleus by the two 1s2 electrons. It
therefore feels a net attraction from the nucleus of 7+ (9 protons
less the 2 screening electrons).
In contrast, chlorine has the electronic structure
1s22s22p63s23px23py23pz1 with 17 protons in the nucleus. But again
the incoming electron feels a net attraction from the nucleus of 7+
(17 protons less the 10 screening electrons in the first and second
levels). There is also a small amount of screening by the 2s
electrons in fluorine and by the 3s electrons in chlorine. This will
be approximately the same in both these cases and so does not
affect the argument in any way (apart from complicating it!).
The over-riding factor is therefore the increased distance that the
incoming electron finds itself from the nucleus as you go down the
group. The greater the distance, the less the attraction and so the
less energy is released as electron affinity.
Comparing fluorine and chlorine is not ideal, because fluorine
breaks the trend in the group. However, comparing chlorine and
bromine, say, makes things seem more difficult because of the
more complicated electronic structures involved. What we have
said so far is perfectly true and applies to the fluorine-chlorine case
as much as to anything else in the group, but there's another factor
which operates as well which we haven't considered yet - and that
over-rides the effect of distance in the case of fluorine.
Why is Fluorine an Anomaly?
The incoming electron is going to be closer to the nucleus in
fluorine than in any other of these elements, so you would expect a
high value of electron affinity. However, because fluorine is such a
small atom, you are putting the new electron into a region of space
already crowded with electrons and there is a significant amount of
repulsion. This repulsion lessens the attraction the incoming
electron feels and so lessens the electron affinity.
A similar reversal of the expected trend happens between oxygen
and sulfur in Group 16. The first electron affinity of oxygen (-142
kJ mol-1) is smaller than that of sulfur (-200 kJ mol-1) for exactly
the same reason that fluorine's is smaller than chlorine's.
Comparing Group 16 and Group 17 values
As you might have noticed, the first electron affinity of oxygen
(−142kJmol −1 ) is less than that of fluorine (−328kJmol −1 ).
Similarly sulfur's (−200kJmol −1 ) is less than chlorine's
(−349kJmol −1 ). Why? It's simply that the Group 16 element has 1
less proton in the nucleus than its next door neighbor in Group 17.
The amount of screening is the same in both. That means that the
net pull from the nucleus is less in Group 16 than in Group 17, and
so the electron affinities are less.
The reactivity of the elements in group 17 falls as you go down the
group - fluorine is the most reactive and iodine the least. Often in
their reactions these elements form their negative ions. The first
impression that is sometimes given that the fall in reactivity is
because the incoming electron is held less strongly as you go down
the group and so the negative ion is less likely to form. That
explanation looks reasonable until you include fluorine!
An overall reaction will be made up of lots of different steps all
involving energy changes, and you cannot safely try to explain a
trend in terms of just one of those steps. Fluorine is much more
reactive than chlorine (despite the lower electron affinity) because
the energy released in other steps in its reactions more than makes
up for the lower amount of energy released as electron affinity.
Second Electron Affinity
You are only ever likely to meet this with respect to the group 16
elements oxygen and sulfur which both form -2 ions. The second
electron affinity is the energy required to add an electron to each
ion in 1 mole of gaseous 1- ions to produce 1 mole of gaseous 2-
ions. This is more easily seen in symbol terms.
X − (g) +e − →X −2 (g) (3)

It is the energy needed to carry out this change per mole of X − .

Why is energy needed to do this? You are forcing an electron into
an already negative ion. It's not going to go in willingly!
O g +e − →O − (g) 1st EA = -142 kJ mol −1 (4)

O − g +e − →O 2− (g) 2nd EA = +844 kJ mol −1 (5)

The positive sign shows that you have to put in energy to perform
this change. The second electron affinity of oxygen is particularly
high because the electron is being forced into a small, very
electron-dense space.

ELECTRONEGATIVITY

This page explains what electronegativity is, and how and why it varies
around the Periodic Table. It looks at the way that electronegativity
differences affect bond type and explains what is meant by polar bonds
and polar molecules.

If you are interested in electronegativity in an organic chemistry context,

you will find a link at the bottom of this page.

What is electronegativity
Definition

Electronegativity is a measure of the tendency of an atom to attract a

bonding pair of electrons.

The Pauling scale is the most commonly used. Fluorine (the most
electronegative element) is assigned a value of 4.0, and values range
down to caesium and francium which are the least electronegative at
0.7.

What happens if two atoms of equal electronegativity bond

together?

Consider a bond between two atoms, A and B. Each atom may be

forming other bonds as well as the one shown - but these are irrelevant
to the argument.

If the atoms are equally electronegative, both have the same tendency
to attract the bonding pair of electrons, and so it will be found on
average half way between the two atoms. To get a bond like this, A and
B would usually have to be the same atom. You will find this sort of
bond in, for example, H2 or Cl2 molecules.

Note: It's important to realise that this is an average

picture. The electrons are actually in a molecular orbital,
and are moving around all the time within that orbital.

This sort of bond could be thought of as being a "pure" covalent bond -

where the electrons are shared evenly between the two atoms.

What happens if B is slightly more electronegative than A?

B will attract the electron pair rather more than A does.

That means that the B end of the bond has more than its fair share of
electron density and so becomes slightly negative. At the same time,
the A end (rather short of electrons) becomes slightly positive. In the
diagram, " " (read as "delta") means "slightly" - so + means "slightly
positive".

Defining polar bonds

This is described as a polar bond. A polar bond is a covalent bond in

which there is a separation of charge between one end and the other -
in other words in which one end is slightly positive and the other slightly
negative. Examples include most covalent bonds. The hydrogen-
chlorine bond in HCl or the hydrogen-oxygen bonds in water are typical.

What happens if B is a lot more electronegative than A?

In this case, the electron pair is dragged right over to B's end of the
bond. To all intents and purposes, A has lost control of its electron, and
B has complete control over both electrons. Ions have been formed.

A "spectrum" of bonds

The implication of all this is that there is no clear-cut division between

covalent and ionic bonds. In a pure covalent bond, the electrons are
held on average exactly half way between the atoms. In a polar bond,
the electrons have been dragged slightly towards one end.

How far does this dragging have to go before the bond counts as ionic?
There is no real answer to that. You normally think of sodium chloride
as being a typically ionic solid, but even here the sodium hasn't
completely lost control of its electron. Because of the properties of
sodium chloride, however, we tend to count it as if it were purely ionic.

Note: Don't worry too much about the exact cut-off point
between polar covalent bonds and ionic bonds. At
A'level, examples will tend to avoid the grey areas - they
 will be obviously covalent or obviously ionic. You will,
however, be expected to realise that those grey areas
exist.

Lithium iodide, on the other hand, would be described as being "ionic

with some covalent character". In this case, the pair of electrons hasn't
moved entirely over to the iodine end of the bond. Lithium iodide, for
example, dissolves in organic solvents like ethanol - not something
which ionic substances normally do.

Summary

 No electronegativity difference between two atoms leads to a pure

non-polar covalent bond.
 A small electronegativity difference leads to a polar covalent bond.
 A large electronegativity difference leads to an ionic bond.

Polar bonds and polar molecules

In a simple molecule like HCl, if the bond is polar, so also is the whole
molecule. What about more complicated molecules?

In CCl4, each bond is polar.

Note: Ordinary lines represent bonds in the plane of the

screen or paper. Dotted lines represent bonds going
away from you into the screen or paper. Wedged lines
represent bonds coming out of the screen or paper
towards you.
The molecule as a whole, however, isn't polar - in the sense that it
doesn't have an end (or a side) which is slightly negative and one which
The molecule as a whole, however, isn't polar - in the sense that it
doesn't have an end (or a side) which is slightly negative and one which
is slightly positive. The whole of the outside of the molecule is
somewhat negative, but there is no overall separation of charge from
top to bottom, or from left to right.

By contrast, CHCl3 is polar.

The hydrogen at the top of the molecule is less electronegative than

carbon and so is slightly positive. This means that the molecule now
has a slightly positive "top" and a slightly negative "bottom", and so is
overall a polar molecule.

A polar molecule will need to be "lop-sided" in some way.

Patterns of electronegativity in the Periodic Table

The most electronegative element is fluorine. If you remember that fact,

everything becomes easy, because electronegativity must always
increase towards fluorine in the Periodic Table.

Note: This simplification ignores the noble gases.

Historically this is because they were believed not to form
bonds - and if they don't form bonds, they can't have an
electronegativity value. Even now that we know that some of
them do form bonds, data sources still don't quote
electronegativity values for them.
Trends in electronegativity across a period

As you go across a period the electronegativity increases. The chart

shows electronegativities from sodium to chlorine - you have to ignore
argon. It doesn't have an electronegativity, because it doesn't form
bonds.

Trends in electronegativity down a group

As you go down a group, electronegativity decreases. (If it increases up

to fluorine, it must decrease as you go down.) The chart shows the
patterns of electronegativity in Groups 1 and 7.
Explaining the patterns in electronegativity

The attraction that a bonding pair of electrons feels for a particular

nucleus depends on:

 the number of protons in the nucleus;

 the distance from the nucleus;
 the amount of screening by inner electrons.

Note: If you aren't happy about the concept of screening or

shielding, it would pay you to read the page on ionisation
energies before you go on. The factors influencing ionisation
energies are just the same as those influencing
electronegativities.

Use the BACK button on your browser to return to this page.

Why does electronegativity increase across a period?

Consider sodium at the beginning of period 3 and chlorine at the end

(ignoring the noble gas, argon). Think of sodium chloride as if it were
covalently bonded.

Both sodium and chlorine have their bonding electrons in the 3-level.
The electron pair is screened from both nuclei by the 1s, 2s and 2p
electrons, but the chlorine nucleus has 6 more protons in it. It is no
wonder the electron pair gets dragged so far towards the chlorine that
ions are formed.

Electronegativity increases across a period because the number of

charges on the nucleus increases. That attracts the bonding pair of
electrons more strongly.

Why does electronegativity fall as you go down a group?

Think of hydrogen fluoride and hydrogen chloride.

The bonding pair is shielded from the fluorine's nucleus only by the 1s2
electrons. In the chlorine case it is shielded by all the 1s22s22p6
electrons.

In each case there is a net pull from the centre of the fluorine or chlorine
of +7. But fluorine has the bonding pair in the 2-level rather than the 3-
level as it is in chlorine. If it is closer to the nucleus, the attraction is
greater.

As you go down a group, electronegativity decreases because the

bonding pair of electrons is increasingly distant from the attraction of the
nucleus.

Diagonal relationships in the Periodic Table

What is a diagonal relationship?

At the beginning of periods 2 and 3 of the Periodic Table, there are

several cases where an element at the top of one group has some
similarities with an element in the next group.

Three examples are shown in the diagram below. Notice that the
similarities occur in elements which are diagonal to each other - not
side-by-side.
For example, boron is a non-metal with some properties rather like
silicon. Unlike the rest of Group 2, beryllium has some properties
resembling aluminium. And lithium has some properties which differ
from the other elements in Group 1, and in some ways resembles
magnesium.

There is said to be a diagonal relationship between these elements.

There are several reasons for this, but each depends on the way atomic
properties like electronegativity vary around the Periodic Table.

So we will have a quick look at this with regard to electronegativity -

which is probably the simplest to explain.

Explaining the diagonal relationship with regard to

electronegativity

Electronegativity increases across the Periodic Table. So, for example,

the electronegativities of beryllium and boron are:

Be 1.5
B 2.0

Electronegativity falls as you go down the Periodic Table. So, for

example, the electronegativities of boron and aluminium are:

B 2.0
Al 1.5

So, comparing Be and Al, you find the values are (by chance) exactly
the same.

The increase from Group 2 to Group 3 is offset by the fall as you go

down Group 3 from boron to aluminium.

Something similar happens from lithium (1.0) to magnesium (1.2), and

from boron (2.0) to silicon (1.8).

In these cases, the electronegativities aren't exactly the same, but are
very close.

Similar electronegativities between the members of these diagonal pairs

means that they are likely to form similar types of bonds, and that will
affect their chemistry. You may well come across examples of this later
on in your course.

Questions to test your understanding

If this is the first set of questions you have done, please read the introductory
page before you start. You will need to use the BACK BUTTON on your
browser to come back here afterwards.

questions on electronegativity

answers

There are no questions on the rest of this page.

Warning! As far as I am aware, none of the UK-based A

level (or equivalent) syllabuses any longer want the next bit.
It used to be on the AQA syllabus, but has been removed
from their new syllabus. At the time of writing, it does,
however, still appear on at least one overseas A level
syllabus (Malta, but there may be others that I'm not aware
of). If in doubt, check your syllabus.

Otherwise, ignore the rest of this page. It is an alternative

(and, to my mind, more awkward) way of looking at the
formation of a polar bond. Reading it unnecessarily just risks
confusing you.
 The polarising ability of positive ions
What do we mean by "polarising ability"?

In the discussion so far, we've looked at the formation of polar bonds

from the point of view of the distortions which occur in a covalent bond if
one atom is more electronegative than the other. But you can also look
at the formation of polar covalent bonds by imagining that you start from
ions.

Solid aluminium chloride is covalent. Imagine instead that it was ionic

Solid aluminium chloride is covalent. Imagine instead that it was ionic. It

would contain Al3+ and Cl- ions.

The aluminium ion is very small and is packed with three positive charges -
the "charge density" is therefore very high. That will have a considerable
effect on any nearby electrons.

We say that the aluminium ions polarise the chloride ions.

In the case of aluminium chloride, the electron pairs are dragged back
towards the aluminium to such an extent that the bonds become covalent.
But because the chlorine is more electronegative than aluminium, the
electron pairs won't be pulled half way between the two atoms, and so the
bond formed will be polar.

Factors affecting polarising ability

Positive ions can have the effect of polarising (electrically distorting) nearby
negative ions. The polarising ability depends on the charge density in the
positive ion.
Polarising ability increases as the positive ion gets smaller and the number
of charges gets larger.

As a negative ion gets bigger, it becomes easier to polarise. For example, in

an iodide ion, I-, the outer electrons are in the 5-level - relatively distant from
the nucleus.

A positive ion would be more effective in attracting a pair of electrons from

an iodide ion than the corresponding electrons in, say, a fluoride ion where
they are much closer to the nucleus.

Aluminium iodide is covalent because the electron pair is easily dragged

away from the iodide ion. On the other hand, aluminium fluoride is ionic
because the aluminium ion can't polarise the small fluoride ion sufficiently to
form a covalent bond.