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Academic Year 2012 – 2013

PERIODIC PROPERTIES AND VARIATIONS OF PROPERTIES

We have studied in the previous class how the periodic law and the periodic table

evolved.

The periodic law states that the properties of elements are a periodic function of their

atomic numbers. So, when the elements are arranged in increasing order of atomic

number, the properties repeat themselves after a particular interval of elements.

The properties that reappear at regular intervals, or in which there is a gradual

variation (i.e. increase or decrease) at regular intervals are called periodic properties.

The phenomenon that brings about these variations is known as the periodicity of

elements.

The cause of periodicity is the recurrence of similar electronic configuration.

In a particular group, electrons in the outermost orbit remain the same or, in other

words, electronic configuration is similar. Valency also remains the same. So,

elements of the same group have similar properties though the number of shells

increases down the group.

The properties that will be discussed here are:

Atomic radius

Ionisation potential

Electron affinity

Electronegativity

Metallic and Non-metallic character

TEST YOUR KNOWLEDGE

1. Define periodicity. 2. What is the cause of periodicity? 3. Why do elements show periodicity in properties? 4. Mention any three properties of elements which show periodicity. 5. State the reasons for periodicity of elements in periods and groups. 6. Why are the elements sodium and chlorine placed in same period of periodic

table?

SN Kansagra School CHAPTER 1

THE PERIODICITY OF PROPERTIES

CAUSE OF PERIODICITY

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The atomic radius is usually considered as the distance from the centre of the nucleus

to the outermost shell of the atom.

Variation of Atomic Radius in a Period

Atomic radii in picometers are given for the elements of the second and third periods

below.

Following figure shows how atomic radius changes across the periods 3.

Changes in atomic size on moving across a period

In a period, atomic radius generally decreases from left to right. It can be

explained as follows. As we go from left to right, electrons are added, one at a time, to

the same outermost shell. A proton is also added one at a time. The outermost

electrons experience increasingly strong nuclear attraction, so the electrons come

closer to the nucleus and more tightly bound to it. This results in decreasing the

atomic radius.

We have to ignore the noble gas at the end of each period because noble gases do

not form bond under normal conditions. Their van-der-Waals’ radius has been shown

in the above figure which is larger than its atomic or covalent radius.

Variation of atomic radii in a group – Atomic radii in picometers are given below

for the alkali metals and halogens.

Reason: Across the period, the effective nuclear charge increases. This is due to the

fact that the number of electrons increase (in the same subshell), increasing the

number of protons in the nucleus. This pulls the valence shell of electrons in an atom

towards itself, thus decreasing the atomic radius. But as we move down the group, the

number of orbits keeps on increasing along with the number of protons. The space

required to accommodate the extra orbits takes prevalence and therefore the atomic

size increases.

ATOMIC RADIUS

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IONIC RADII: When an atom is converted to an ion, the size of the neutral atom

changes. An anion is bigger than a neutral atom. This is because addition of one or

more electrons increases repulsion among electrons and they move away from each

other. On the other hand, a cation is smaller than the neutral atom. When one or more

electrons are removed, the repulsive force between the remaining electrons decreases

and they come a little closer. This is shown in the following figure.

Figure – Cation is smaller than atom. Anion is bigger than atom

TEST YOUR KNOWLEDGE

1. What happens to atomic radii in a group and period and why? 2. What is the atomic radius of an atom? 3. What is the trend in atomic radius across a period? 4. The trend in atomic radius across a period is caused by _____. 5. What generally happens to atomic radii as one goes down a group or a family? 6. State the factors which affect size of elements in a periodic table. In period 2

from left to right, state which element has the largest atomic size and which has the smallest, giving reasons.

7. Why is cation (Na+) smaller than the parent atom (Na)? 8. Why is anion (Cl-) larger than the parent atom (Cl)? 9. Atomic size of group 18 elements is more than the atomic size of group 17

elements. 10. Which feature of the atomic structure accounts for the similarities in chemical

properties of elements in group 17 (or VIIA)?

Across the period i.e., from left to right:

Atomic radius decreases

Down the group i.e., from top to bottom:

Atomic radius increases

Na Na+

Cl Cl-

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It is the amount of energy required to remove one or more electrons from the valence

shell of an isolated gaseous atoms.

The energy required to remove the first electron is called the first ionisation potential

[E1]. Similarly, the energy required to remove second and third electron from an atom

is called the second ionisation potential [E2] and the third ionisation potential [E3].

The first ionisation potential is least. The second ionisation potential is greater than

the third ionisation potential is still higher, i.e., E3 > E2 > E1.

Reason: As each electron is removed, the effective attraction of the nucleus on the

remaining valence electrons is increased. Hence, more energy is required to pull out

the successive electrons.

Factors Influencing Ionisation Energy:

The ionisation energy of an atom depends upon how tightly the outermost electron is

held by the nucleus. This is influenced mainly by the following factors.

Atomic size - The greater the atomic size, the farther is the outermost electron and so

the easier it is to remove that electron. Hence, as the atomic size increases, the

ionisation energy of the atom decreases.

Nuclear charge - As the nucleus charge increases, the pull of the nucleus on the

outermost electron increases and so it becomes more difficult to remove that electron.

Hence, as the nuclear charge increases, the ionisation energy of the atom increases.

Variation of Ionisation Potential in the Periodic Table

Across the period i.e., from left to right: Ionisation potential increases

Down the group i.e., from top to bottom: Ionisation potential decreases

Reason: Across the period, the effective nuclear charge increases. This causes the

atomic radius to decrease, thus getting the valence shell closer to the nucleus. This

makes it difficult to remove electrons. But as we move down the group, the number of

orbits keeps on increasing along with the number of electrons. The distance from the

IONIZATION ENERGY (POTENTIAL)

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nucleus coupled with the interference of the electron between the nucleus and the

valence shell renders the valence electrons weakly bound to the nucleus.

TEST YOUR KNOWLEDGE

1. What is ionization energy? 2. Write the equation for the ionization of an atom. 3. What is an ion? 4. Why are the ionization energy of elements increases in a period from left to right? 5. Which group or family has the lowest ionization energy? 6. Group 18, the noble gases, have the highest ionization energy (True or False). 7. Elements with a high ionization energy lose electrons easily (True or False). 8. The increase in ionization energy across a period is caused by _____. 9. Why does ionization energy generally decrease going down a group or family? 10. What is the second ionization energy of an atom? 11. Why does fluorine have a higher ionization energy than iodine? 12. A decrease in ionization potential of an element leads to a decrease in non metallic character of the element. 13. What are the factors which influence or affect ionization potential of elements

in a periodic table? 14. Why elements with low ionisation potential exhibit metallic properties?

Another important property that determines the chemical properties of an element is

the tendency to gain an additional electron. This ability is measured by electron

affinity.

It is the amount of energy released when an electron is added to an isolated gaseous

atom. Electron affinity is expressed in electron volt (eV).

X(g) + e– → X- (g) + Energy released

Factors Influencing Electron affinity:

Atomic size – The smaller the atom the greater is the electron affinity.

Reason: This is because the effective attractive force between the nucleus and the

valence electron is greater for the smaller atoms and they can hold the extra electrons

more firmly.

ELECTRON AFFINITY

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For example: Halogens with smaller atomic size have a high electron affinity and

readily forms anions. Alkali metals, with large atomic radii, have low electron affinity

and do not form anions.

Nuclear charge – The greater the nuclear charge the greater is the electron affinity.

Reason: This is because the increase in nuclear charge increases the effective

attractive force on the valence electrons to hold the additional electron in the valence

shell.

For example: Fluorine has a greater electron affinity than oxygen because fluorine

has nine positive charge in its nucleus while oxygen has only eight.

Note: The electron affinity of elements having completely filled orbitals or less than

half-filled orbitals is practically zero.

Reason: Elements with completely filled sub-shells (ex. Noble gases) have no scope

of adding an extra electron. Hence, they have zero electron affinity and do not form

anions.

Variation of Electron Affinity in the Periodic Table

Across the period i.e., from left to right: Electron affinity increases

Reason: This is because both electron affinity and ionization energy are highly related

to atomic size. Large atoms have low ionization energy and low electron affinity.

Therefore, they tend to lose electrons. In general, the opposite is true for small atoms.

Since they are small, they have high ionization energies and high electron affinities.

Therefore, the small atoms tend to gain electrons.

The major exception to this rule is the noble gases. Noble gases follow the general

trend for ionization energies, but do not follow the general trend for electron affinities.

Even though the noble gases are small atoms, their outer energy levels are completely

filled with electrons. Any added electron cannot enter their outer most energy level

and would have to be the first electron in a new (larger) energy level. This causes the

noble gases to have essentially zero electron affinity.

Down the group i.e., from top to bottom: Electron affinity decreases

Reason: Going down a group, the electron affinity generally decreases because of the

increase in size of the atoms. Remember that within a family, atoms located lower on

the periodic table are larger because there are more filled energy levels. When an

electron is added to a large atom, less energy is released because the electron cannot

move as close to the nucleus as it can in a smaller atom. Therefore, as the atoms in a

family get larger, the electron affinity gets smaller.

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TEST YOUR KNOWLEDGE

1. What do you mean by Electron affinity? 2. What is the difference between electron affinity and ionization energy? 3. Write the equation for electron affinity for an exothermic process. 4. Explain the trend in general of electron affinity of elements- a) on moving from left to right across a period. b) on moving down a group. 5. Which group or family gains electrons most easily? 6. Compare the electron affinities of metals and non-metals, in general. 7. What are the factors which influence electron affinity? 8. Give the order of electron affinities of the halogens. 9. Which will have greater electron affinity Oxygen or Fluorine? 10. Electron affinity of noble gas elements is zero.

The tendency of an atom to attract a bonding pair of electron towards itself when

combined in a compound is called electronegativity.

It is a dimensionless quantity and does not have any unit. It is a relative property.

Electron affinity is a property of gaseous isolated atoms. We normally do not deal

with isolated atoms. Instead, we come across atoms which are bonded to each other.

So electronegativity is a more useful property. It helps to understand the nature of

chemical bond between two atoms.

POLAR AND NON POLAR COVALENT BOND

Electronegativity cannot be directly measured and must be calculated from other

atomic or molecular properties. Several methods of calculation have been proposed

and, although there may be small differences in the numerical values of the

electronegativity, all methods show the same periodic trends between elements. Linus

Pauling’s scale of electronegativity is more in use. On this scale the highest

electronegativity is 4.0 for fluorine and lowest electronegativity is 0.7 for caesium.

When two atoms of an element combine by sharing electrons, the electrons are shared

equally by the two atoms. There is no drift of electrons towards any one of them, and

the bond between the two atoms is said to be nonpolar covalent, as in H2, N2, Cl2, O2

etc.

When two atoms belong to different elements, the shared pair of electron is attracted

by one of the atoms, giving a partial + and – charges to the two atoms. And the bond

formed between the two atoms is said to be polar covalent, as in H2O, HCl, NH3 etc.

ELECTRONEGATIVITY

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Dipole effect in Water molecule

Following table shows electronegativities of elements according to Pauling scale.

The bond between caesium (Cs) and fluorine (F) is an ionic bond and CsF is the

most ionic compound because the difference in the electronegativity between Cs and

F is maximum.

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There is not much difference is the electronegativity of nitrogen (N) and oxygen (O).

Hence, a covalent bond is formed between them.

Variation of Electronegativity in the Periodic Table

Across the period i.e., from left to right: Electronegativity increases

Down the group i.e., from top to bottom: Electronegativity decreases

Reason: Across the period, the effective nuclear charge increases, thus decreasing the

atomic radius. This favours the increase in electronegativity of elements across the

period. But as we move down the group, the number of orbits keeps on increasing and

therefore the atomic size increases and the electronegativity decreases.

Differences between Electronegativity and Electron Affinity

Electronegativity Electron affinity

1. It is the tendency of an atom of an

element to attract shared pair of

electrons towards itself in a molecule.

2. It is the property of the bonded atom.

3. It is simply a number and has no

Units.

1. It is the amount of energy released

when an electron is added to an

isolated neutral gaseous atom

present in the gaseous state so as to

form an anion.

2. It is the property of an isolated atom.

3. It has units, i.e. eV/atom, kJ/mole

and kcal/mole.

TEST YOUR KNOWLEDGE

1. What property of an element is measured by electronegativity? 2. Fluorine is the most electronegative element of the periodic table. 3. Electronegativity of chlorine is higher than that of sulphur. 4. Explain the trend in general of electronegativity of elements across a period and down a group. 5. Write the difference between polar and non-polar bond.

It is also possible that a multi-bond molecule is non-

polar but the individual bonds in the molecule are

polar. If a bond is formed between the different

atoms, it has to be polar.

A molecule of methane is non-polar whereas all the

four bond C-H bonds in the molecule are polar.

Reason: The net effect of the polarity of four

bonds, taking into account their directions, is zero.

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An element which has a tendency to lose electrons when supplied with energy is

considered as a metal.

M → M+ + e-

Metal ion

An element which has a tendency to gain electrons when energy is released is

considered as a non-metal.

N + e- → N-

Non-metal ion

Factors which affect the metallic character

Atomic size: More energy is required to remove an electron from small atom than

from a large atom. Therefore, atoms with larger atomic radii give up electrons easily

and acquire metallic characteristics.

Reason: In small atom, the attraction between the nucleus and the electrons is strong

Hence, more energy is required to overcome the attractive forces between them.

Ionisation potential: Metallic character increases with decrease in ionisation

potential while non-metallic character increases with increase in ionisation potential.

Reason: Lower the ionisation potential, the greater is the tendency of an atom to lose

electrons.

Variation of Metallic and Non-metallic character in periodic table:

In a Period

In going from left to right across a period, the metallic nature decreases and the

nuclear pull increases due to the increase in the atomic number. The atomic size of the

element gradually decreases. Hence, elements cannot lose electrons easily. Therefore,

the metallic nature decreases while the non-metallic nature increases.

Thus, in the third period, sodium, magnesium and aluminium are metallic while

silicon, phosphorus, sulphur and chlorine are non-metallic.

Reason: The atomic radii of the elements gradually decease across a period and the

ionisation potential increases. Hence, the tendency of an atom to lose electrons

decreases.

METALLIC AND NON-METALLIC

NATURE

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In a Group

The metallic character increases and non-metallic character decreases in descending

from top to bottom. Thus, in Group 15, nitrogen and phosphorus are typical non-

metals; arsenic is a metalloid; while antimony and bismuth are well defined

metals.

Reason: With increase in atomic number, the metallic character or the atomic size

increases and, though the nuclear charge also increases, yet the effect of the increased

atomic size is greater compared to the increased nuclear charge. Hence, metallic

nature increases and non-metallic nature decreases down the group from top to

bottom.

TEST YOUR KNOWLEDGE

1. Explain the trends from metallic to non-metallic character of the different

elements in the first three periods. 2. With reference to any one group of the periodic table explain with reasons the

trends in metallic and non-metallic character down a group. 3. State the factors which affect the metallic and the non-metallic character of

elements in a periodic table.

GENERAL TENDENCY OF PERIODIC PROPERTIES

PERIODIC PROPERTIES ALONG A PERIOD DOWN THE GROUP

ATOMIC RADIUS Decreases Increases

IONIZATION ENERGY Increases Decreases

ELECTRON AFFINITY Increases Decreases

ELECTRONEGATIVITY Increases Decreases

METALLIC PROPERTY Decreases Increases

NON-METALLIC PROPERTY Increases Decreases

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STUDY OF SELECTED GROUPS

ALKALI METALS – GROUP 1 [1A] AND HALOGENS – GROUP 17 [VIIA]

PROPERTIES ALKALI METALS HALOGENS

PHYSICAL

1. Elements

2. Valence electrons

3. Nature

4. Conduction of

heat and electricity

5. Melting Point and

Boiling point

6. Atomic size

7. Ionization Potential

CHEMICAL

1. Reactions with Non-

Metals

2. Reactions with

Hydrogen

3. Reaction with Water

4. Nature of oxides

5. Reducing /

Oxidizing nature

-Lithium, sodium, potassium,

rubidium, caesium, francium

- 1

-Highly reactive, highly

electropositive, light, soft

metals.

-Good conductors

-Decreases down the group.

Ex. Li-13300C and Cs-6900C

-Largest in their periods

except noble gases

-Decreases down the group as

a result the tendency to lose

electrons and form positive

ions. Francium is the most

electropositive element.

-Electrovalent compounds

formed [Ex: NaCl, KBr]

-Ionic hydrides formed

[Ex: LiH, NaH]

-Reacts vigorously to form

hydroxides and liberating

hydrogen.

-They form basic oxides.

Strong reducing agents

[Alkali metals-electron donor]

Fluorine, chlorine, bromine,

iodine, astatine

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-Highly reactive, highly

electropositive, Non-metals,

Gaseous – F, Cl, Liquid- Br,

Solid-I

-Bad conductors

-Increases with increase in

atomic number.

-Smallest in their periods.

-Increases down the group.

High; lower only than the

noble gases.

-Covalent compounds

formed [Ex. HCl, PCl3]

-Covalent hydrides

[Ex. HF, HCl]

-They do not liberate

hydrogen.

-They form acidic oxides.

Strong oxidizing agent

[Halogens-electron

acceptors]

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TEST YOUR KNOWLEDGE

Q. Name the following:

1. The non-metal is liquid at room temperature.

2. The non-metal which strong oxidizing agent and destroys germs.

3. The gas which is yellowish green in colour.

4. The non-metal which is lustrous.

5. The non-metal which is a subliming substance.

6. The metal that burns in air with a golden yellow flame.

7. An ionic hydride.

8. A covalent hydride.

9. A metal which is stored under kerosene oil.

10. The smallest atom in the periodic table.

11. A non-metal which is solid at room temperature.

12. The most electropositive element in the periodic table.

13. The largest atom in the periodic element.

14. The radioactive element in group VII A.

15. A metal which reacts violently with water.

References / Figures / Diagrams –

Method of measurement of atomic radius http://www.chemguide.co.uk/atoms/properties/atradius.html

Atomic radius decreases across the periods 2 and 3

http://www.chemguide.co.uk/atoms/properties/atradius.html

Cation is smaller than atom. Anion is bigger than atom.

http://www.chemguide.co.uk/atoms/properties/atradius.html

Periodic table of electron affinities of elements http://www.wikipedia.org/wiki/Electron_affinity

Periodic table of electronegativities of elements according to Pauling Scale www.knowledgerush.com/kr/encyclopedia/Pauling_Electronegativity_Scale

Periodic properties

http://net.mkcl.org/WebFiles/Periodic%20Classification%20of%20Elements.pdf