Periodicity

Question 1

what is nuclear charge and why does it increase down a group?

Answer 1

refers to the number of protons in the nucleus of an atom

going down a group, successive elements contain more protons = increasing nuclear charge

Question 2

why does nuclear charge increases across a period (left to right)?

Answer 2

going across a period, successive elements contain more protons = increasing nuclear charge

Question 3

what is shielding

Answer 3

the attraction between outer shell electrons and the nuclear charge is shielded by the inner core electrons

Question 4

why does the shielding effect increase down a group?

Answer 4

going down a group, the number of inner shell electrons increases = more shielding

Question 5

why does the shielding effect remain constant across a period (left to right)?

Answer 5

going across a period, electrons are only added to the valence shell. The number of inner shell electrons remains constant = constant shielding

Question 6

what is core charge (effective nuclear charge)

Answer 6

core charge or effective nuclear charge is the effective pull of the nucleus on the outer shell electrons

determined by:
core charge = nuclear charge – number of the shielding electrons

Question 7

why does core charge (effective nuclear charge) increase across a period (left to right)

Answer 7

going across a period, the nuclear charge increase but the number of shielding electrons remains the same = increasing core charge

Question 8

why do atomic radii increase down a group?

Answer 8

going down a group, each successive element contains another shell whilst core charge remains the same = larger atom

Question 9

why do atomic radii decrease across a period (left to right)?

Answer 9

going across a period, core charge increases whilst the number of occupied shells remains the same , drawing in electrons = smaller atom

Question 10

what is the pattern in the ionic radii of the Group 1 elements

Answer 10

ionic radii increases down the group

Question 11

compare the atomic and ionic radii of Group 1 elements. Explain the difference in sizes

Answer 11

the ions are smaller. This is because the ions contain one less shell

Question 12

compare the atomic and ionic radii of the Group 17 elements. Explain the difference in sizes

Answer 12

the ions are larger. This is because the ions contain more electrons in the outer shell which results in more repulsion between electrons = larger ion

Question 13

describe and explain the trend in ionic radii for the elements in period 3

Answer 13

Ions decreases in size from Na+ to Si4+. There is then a sudden increase in size (P3– is larger than Na+). Then ions decrease in size from P3- to Cl-.
Explanation:
- Decrease in ionic radii is due to increasing core charge.
- The negative ions contain an extra shell so they’re larger

Question 14

First ionisation energy definition

Answer 14

First ionization energy – The energy required to remove one mole of electrons from one mole of gaseous atoms
to form one mole of unipositive ions

(an atom of an element has as many ionisation energies as there are electrons)

Question 15

Why does first IE decrease down a group?

Answer 15

going down a group, valence electrons are further from the nucleus, so less attracted to nucleus so less energy is required to remove a valence electron

Question 16

Why does first IE generally increase across a period (left to right)?

Answer 16

going across a period, core charge increases, so valence electrons are more attracted to the nucleus so more energy is required to remove a valence electron

Question 17

First electron affinity definition

Answer 17

First electron affinity – The energy required to add one mole of electrons to one mole of gaseous atoms to form
one mole of uninegative ions
(elements can have multiple electron affinities)

Since the electron is naturally attracted to the positively charged nucleus of an atom, this is an exothermic process

Question 18

Why does first EA generally decrease down a group?

Answer 18

going down a group, valence electrons are further from the nucleus so there is less attraction for an electron

Question 19

Why does first EA generally increase across a period (left to right)?

Answer 19

going across a period, core charge increases so there is greater attraction for an electron

Question 20

Why do most of the Period 2 elements have a lower first EA than their equivalent Period 3 elements?

Answer 20

the electrons in shell 2 are closer together, so there is a slightly greater repulsion between them compared to the electrons in shell 3. This decreases the attraction for an electron

Question 21

Electronegativity definition

Answer 21

Electronegativity – A measure of the tendency of an atom in a covalent molecule to attract a pair of shared electrons towards itself. The difference in electronegativity of two atoms in a covalent bond gives an indication of the bond’s polarity

Question 22

How can the pattern for each group’s electronegativity be explained in terms of atomic structure

Answer 22

going down a group, bonding electrons are further from the nucleus so they are less strongly attracted to it

Question 23

ow can the pattern for each group’s electronegativity across Period 3 be explained in terms of atomic structure

Answer 23

going across a period, core charge increases so bonding electrons are closer to the nucleus so they are more strong attracted to it

Question 24

up to pg 10 SL