Classification of Elements and Periodicity in Properties Flashcards
s-block
Groups 1 and 2
p-block
Groups 13-18
d-block
Groups 3-12
f-block
Lanthanoids and Actinoids at the bottom of the periodic table
Normal/respresentative elements
s-block and p-block elements
Atomic radius- definition
The distance between the centre of the nucleus and the valence shell of electrons is known as the atomic radius
Crystal radius
One half of the distance between the centre of the nuclei of two adjacent atoms in a metallic nucleus
Covalent radius
One half of the distance between the centres of two nuclei of two similar atoms bonded together by a single covalent bond.
van der Waals’ radius
One half of the distance between the nuclei of two atoms of the same substance at their closest approach
How does atomic radius vary
1. across a period
2. down a group
+ why?
- Atomic radius decreases because no. of shells remain the same but the nuclear charge increases, pulling the electron cloud closer to the nucleus.
- Atomic radius increases because no. of shells increase and this takes the valence shell further away from the nucleus, decreasing the effective nuclear charge.
Ionic radius- definition
The distance between the centre of the nucleus and the point up to which the nucleus has an influence on the electron cloud of the ion.
Why is the radius of a cation always less than the radius of the atom from which it is formed?
- Cations are positively charged ions and are formed when a neutral atom loses electrons.
- The number of valence electrons decrease but the positive charge in the nucleus remains the same.
- Due to less number of electrons and same nuclear charge, the effective nuclear charge on the remaining electrons is more, causing them to come closer to the nucleus.
Why is the radius of an anion always more than the radius of the atom from which it is formed?
- An anion possesses same nuclear charge but more electrons than the neutral/parent atom.
- This leads to a decrease in the effective nuclear charge i.e. the nucleus exerts less influence on the valence electrons.
- The valence electrons move away and the size increases due to the expansion of the electron cloud.
Isoelectronic ions
Ions of different elements which contain the same number of electrons are called isoelectronic elements
Relationship between the ionic radii of isoelectronic ions and nuclear charge
ionic radii of isoelectronic compounds decrease with increase in nuclear charge.
Ionisation enthalpy or ionisation energy
The amount of energy required to remove a loosely bound electron from the outermost shell of an isolated gaseous atom to form a gaseous ion
Unit of ionisation enthalpy
kJ/mol
Ionisation potential
The voltage at which the ionisation of a gas occurs
Why is the second ionisation energy higher than the first ionisation energy?
This is because the removal of a loosely bound electron from the valence shell of the ion forms a monopositive ion, whose ionic radius is much less than the atomic radius of the parent atom. Due to smaller size of the cation, the remaining electrons will experience a greater pull of the nucleus. Therefore, higher energy will be required to remove another electron.
How does ionisation energy depend upon:
1) atomic radius
2) nuclear charge
1) It decreases with an increase in atomic radius
2) It decreases with a decrease in nuclear charge
What is screening/shielding effect?
In a multielectron atom, the electrons present between the valence electron and the nucleus shield the valence electron from the nucleus.
This decreases the effect of the nucleus on the valence electron i.e. the nucleus exerts less force of attraction on the valence electron.
Dependence of ionisation energy on shielding effect.
If other factors remain the same, ionisation energy increases with a decrease in shielding effect.
What is penetration effect?
In a multielectron atom, the electrons do not maintain distinct boundaries. Instead the electron cloud of one electron penetrates into the electron cloud of another inner electron. This is known as penetration effect.
Due to this effect, electrons move towards the nucleus.
Dependence of ionisation energy on penetration power.
Increases with increase in penetration power
Why are alkali metals highly reactive?
They have extremely low ionisation energies and get easily ionised, making them highly reactive.
Variation of ionisation energy across a period + why
Increases across a period because:
i) nuclear charge increases
ii) size of the atom decreases
Variation of of ionisation energy down a group + why
Ionisation energy decreases down a group because:
i) atomic radius increases
ii) no. of inner electrons increase so shielding effect increases
Electron affinity
The amount of energy released on the addition of an electron to the outermost shell of an isolated gaseous atom to form a gaseous anion.
Dependence of electron gain enthalpy on:
- nuclear charge
- size of the atom
- electron gain enthalpy increases with increase in nuclear charge
- electron gain enthalpy increases with decrease in atomic size
Variation of electron gain enthalpy:
- across a period
- down a group
- becomes more negative as we move across the period
2. becomes less negative as we move down a group
Why do halogens have the most negative electron gain enthalpies?
- Halogens are elements of group 17.
- Their valence electronic configuration is of ns²np⁵ type which is similar to that of noble gases (ns²np⁶)
- Halogens have a very strong tendency to accept electrons.
- Therefore they have the most negative electron gain enthalpies.
Why is the electron gain enthalpy of chlorine more negative than that of fluorine?
- In F, the added electron goes to the 2p subshell while in Cl it goes to the 3p subshell.
- The 2p subshell is more compact compared to the 3p subshell as the 2p subshell is closer to the nucleus.
- Due to the small size of the F atom the inter-electronic repulsion in 2p subshell of F is much larger as compared to that of the 3p subshell of Cl.
- Therefore, the added electrons find it easier to enter into the Cl atom as compared to the F atom.
Why is the electron gain enthalpy of nitrogen positive?
N possesses a half filled shell. Due to this relatively stable electronic configuration, nitrogen atom does not accept the added electron easily.
the addition of an electron to a nitrogen atom is only possible when external energy is supplied to it
Electronegativity
The tendency of an atom to attract the bonding or shared pair of electrons towards itself in a covalent bond is called the electronegativity of that atom.
Dependence of electronegativity on:
- Effective nuclear charge
- Size of the atom
- Oxidation state
- State of hybridization
- More the effective nuclear charge, more the electronegativity
- Less the size of the atom, more the electronegativity
- Greater the oxidation state. more the electronegativity
- Greater the s-character in the hybridization state of an atom, more the electronegativity
