Upload
others
View
1
Download
0
Embed Size (px)
Citation preview
1
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
2
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
3
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-
4
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)
5
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
6
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.
7
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
8
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.
9
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.
10
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
11
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
12
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
- 7
-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]
13
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