Periodicity

Question 1

what is periodicity

Answer 1

it is a pattern of repeating trends

Question 2

explain why across a period, atomic radius decreases

Answer 2

across a period, the number of electrons in the outershell increases. This increases the positive charge. Because of this positive charge, there is an increased attraction between the protons in the nucleus and the electrons. This draws the electrons closer to the nucleus, hence, causing atomic radius to decrease

Question 3

Across a period, what happens to shielding effect and why?

Answer 3

All the elements in period two have only one inner electron electron shell. This means that shielding effect is constant/ the same across a period

Question 4

what happens to the atomic radius down a group and why?

Answer 4

Down a group, atomic radius increases. This is because the number of electron shells increases as we move down the group, so the outer electron shell is far from the nucleus. Also, each element has one more inner electron shell, so shielding is increased. This reduces the attraction of the nucleus to the outer electrons

Question 5

what happens to first ionization energy across a period and why?

Answer 5

it generally increases across a period. Shielding effect is constant and the nuclear attraction increases due to the addition of electrons

Question 6

In period 2, boron and oxygen do not fit into this trend. why?

Answer 6

BORON:
Looking at the electronic configuration of boron and beryllium.
Be= 1s2 2s2.
B= 1s2 2s2 2p1
Boron has only one electron in its p orbital, hence less energy is required to remove this electron than in beryllium, whose orbital is full/stable

OXYGEN:
looking at the electronic configuration of nitrogen and oxygen, we need to take a closer look at the 2p subshell

In nitrogen, all the electrons are in separate orbitals, but in oxygen, there is a pair of electrons. These electrons repel each other, hence less energy is required to remove these electrons

Question 7

List 3 properties of metals

Answer 7

1. They are good conductors of electricity when they are in solids or liquids due to their delocalized electrons being free to move.
2. Most metals have relatively high melting and boiling points.This is because of the strength of the metallic bond. The electrostatic attraction between the delocalized electrons to the cations is very strong, so great amounts of energ are required to overcome this energy. This results in their high boiling and melting point
3. Metals do not dissolve. They react with water, not dissolve

Question 8

Melting and boiling point generally increases across a period. why?

Answer 8

for example, in period 2, lithium and beryllium have high melting point due to their metallic bonding. Boron and Carbon have high melting point because they form giant covalent structures. N,O,F are diatomicand Ne is a noble gas. They all form simple molecular structures

Question 9

explain the giant covalent structures in Boron and Carbon

Answer 9

In a giant covalent structure, the numerous atoms are joined by strong covalent bonds and together they form a giant covalent lattice. It takes a lot tof energy to break all the covalent bonds in a giant covalent structure. This explains their high melting and boiling point

Question 10

List three properties of giant covalent structures (diamond)

Answer 10

1.Has high melting and boiling point
2.Does not conduct electricity. There are no delocalized electron
3. They are insoluble