CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES Flashcards
questions
Who introduced the periodic law of elements for the first time? State the law.
Mendeleev. It states that “the properties of elements are the periodic functions of their
atomic masses.”
State Modern Periodic law. Who proposed this law?
Modern periodic law states that “the properties of elements are the periodic functions of
their atomic numbers.” This is law was proposed by Henry Moseley.
Explain the different blocks in Modern periodic table.
There are 4 blocks in modern periodic table. They are s block, p block, d block and f block.
i. The s block elements: These are elements in which the last electron enters in the outer most s sub shell. They include elements of the groups 1 and 2. Their general outer electronic configuration is ns1 or ns2.
ii. The p block elements: These are elements in which the last electron enters in the outer most p sub shell. They include elements of the groups 13 to 18. Their general outer electronic configuration is ns2 np1 to 6.
iii. The d block elements: These are elements in which the last electron enters in the penultimate d sub shell. They include elements of the groups 3 to 12. Their general outer electronic configuration is (n–1)d1 to 10 ns0 to 2.
iv. The f block elements: These are elements in which the last electron enters in the anti–penultimate f sub shell. They include lanthanides and actinides. Their general outer electronic configuration is (n–2)f1 to 14(n–1)d0 to 1 ns2
Write any 4 properties of transition elements?
They are all metals, they form coloured compounds or ions, they show variable
oxidation states and valencies, they show paramagnetism and catalytic properties
How does atomic radius vary along a group and period and why?
The atomic size decreases from left to right in a period, due to increase in the effective
nuclear charge. Also, the number of shells remain constant in a period.
Down a group, the atomic radius increases from top to bottom, due to the increase in the number of shells and shielding effect.
What are isoelectronic species? Give examples.
Atoms and ions having the same number of electrons are called isoelectronic species.
E.g. O2–, F–, Ne, Na+, Mg2+ etc. (All these contain 10 electrons).
Define ionisation enthalpy.
It is defined as the energy required to remove an electron from the outer most shell of
an isolated gaseous atom in its ground state. Its unit is kJ/mol or J/mol.
How does ionisation enthalpy vary along a period and group? Justify your answer.
: Along a period, ionisation enthalpy increases from left to right, due to decrease in
atomic radius and increase in nuclear charge.
Down a group, ionization enthalpy decreases due to increase in atomic radius and shielding
effect.
The first ionisation enthalpy of Boron is slightly less than that of Beryllium. Why?
: Because of the stable fully filled electronic configuration of B+
ion. ((B+)– 1s22s2 2p1)
The first ionisation enthalpy of Nitrogen is greater than that of Oxygen. Why?
This is because Nitrogen has half–filled p–orbital configuration (1s22s22p3
), which is
more stable and so more energy is required to remove an electron.
Chlorine has higher negative electron gain enthalpy than fluorine. Why?
Or, Electron gain enthalpy of fluorine is less negative than chlorine. Why?
Due to the compactness of the 2p subshell of F, electron–electron repulsion is greater
in F and hence it does not easily add electron. [Or, Due to larger size and less electron–electron repulsion in chlorine].
Define electron gain enthalpy (∆egH).
It is the enthalpy change when an electron is added to the outer most shell of an
isolated gaseous atom. Its unit is kJ/mol.
Phosphorus forms PCl5 while nitrogen cannot form NCl5. Why?
Because of the absence of vacant d-orbitals in nitrogen.
. Electron gain enthalpy values of noble gases are zero. Why?
This is because of their stable fully filled electronic configuration
Define electronegativity.
Electronegativity of an atom in a compound is the ability of the atom to attract shared
pair of electrons.
