3.2.1- Periodicity (PAPER 1)

Question 1

What is periodicity?

Answer 1

A repeating pattern in properties of elements across periods of the periodic table.

Question 2

Where are s,d,p,f blocks on the periodic table?

Answer 2

s= Group 1 and 2
d= middle section between Group 2 and 3
p= Group 3- Group 0
f= transition metals at bottom

Question 3

What classifies an element as being in a specific block?

Answer 3

Their highest energy/ outer electron(s) are in the block.

Question 4

What was Mendeleev’s statement?

Answer 4

When elements are arranged in order of atomic mass, there are recurring patterns in certain properties.

Question 5

Why is atomic radius of an element difficult to define?

Answer 5

There is uncertainty over the size of the electron cloud.

Question 6

What is one definition of atomic radius?

Answer 6

Half the shortest internuclear distance found in the structure of the element.

Question 7

What is the covalent radius of non metallic elements?

Answer 7

Half the internuclear distance between 2 identical atoms in a single covalent bond.

Question 8

Draw a diagram of 2 circles + equation to show how you could calculate atomic radius of 2 covalently bonded atoms.

Answer 8

2 circles next to eachother, dot in middle for nucleus: line between 2 dots labelled d.

Equation = d/2

Question 9

What is the van der waal’s radius for non-bonded adjacent atoms?

Answer 9

Half the shortest internuclear distance between 2 similar non-bonded atoms- eg in a covalent crystal of a non-metallic element.

Question 10

What is the metallic radius for metallic elements?

Answer 10

Half the shortest internuclear distance between 2 adjacent atoms in a metallic bond.

Question 11

What is the trend in atomic radius across Period 3? (proton n +shielding)

Answer 11

Across P3, atomic radius decreases:

-Proton number increases, increasing the nuclear charge of the atom, the nucleus has a stronger attraction for the electrons and pulls them in closer.

-The extra electrons are added to the outer energy level so they do not provide extra shielding effect ( as shielding occurs when inner electron shells repel outer electron shells).

Question 12

What increases when proton number increases and shielding remains constant?

Answer 12

Effective nuclear charge

Question 13

Why does argon not follow general trend along Period 3?

Answer 13

Argon particles are single atoms with electrons closer to nucleus so are not easily polarised: no permanent dipole-dipole attractions, only induced vDWs.

The orbitals are not changing size as there is no movement of electrons, true atomic radius.

Question 14

What is the 1st ionisation energy of an element?

Answer 14

The energy required to remove one mole of electrons from one mole of gaseous atoms.

Question 15

What is the trend in 1st ionisation energy across period 3? (proton n, shielding)

Answer 15

General increase in 1st ionisation energies across Period 3:

-Proton number increases but shielding remains relatively constant so effective nuclear charge increases, greater attraction between nucleus and outer electrons so more energy required to remove an electron.

Question 16

Explain why Al has a lower 1st ionisation energy than Mg?

Answer 16

Although Al has higher nuclear charge, electron removed when it is ionised is in a 3p sub-level, which is higher in energy than the 3s sub-level the electron in Mg is removed from.

Question 17

Explain why S has a lower first ionisation energy than P?

Answer 17

Although S has higher nuclear charge, both elements outer electrons are in same 3p sub-level but electron is paired in P, mutual repulsion between paired electrons means less energy is needed to remove one.

Question 18

What is electronegativity?

Answer 18

The power of an atom to attract the 2 electrons in a covalent bond.

Question 19

What is the trend in electronegativity across a period?

Answer 19

Electronegativity increases as:

there are more protons in the nucleus, so elements have smaller atomic radius and there is a stronger attraction between nucleus and 2 electrons in covalent bond.

Question 20

Why do the melting points of Na, Mg and Al increase in Period 3? (charge density, e-)

Answer 20

There is an increase in their metallic bond strength:

1) Charge density- Na+ ions are large with small charge so low charge density, Al3+ ions smaller with larger charge so higher charge density and more attracted to delocalised electrons.

2) Number of delocalised electrons- Number of protons/ electrons increases across the period- eg Al has 3 free electrons per metal ion so more attractions that must be broken, increasing positive charge also decreases atomic radius.

Question 21

Why does silicon have a higher m+b point than other Period 3 elements?

Answer 21

Silicon has macromolecular structure: each Si atom strongly covalently bonded to 4 other Si atoms.

These covalent bonds must be broken in order to melt Si and requires a lot of energy.

Question 22

Why do P,S, C, Ar have lower m+b point than silicon?

Answer 22

They have a simple molecular structure with weak VdW forces holding the molecules together, require less energy to overcome than covalent bonds.

Question 23

What is the structure of phosphorous?

Answer 23

Forms tetrahedral P4 molecules which are the second largest molecules of Period 3 non metals.

Second highest m+b points of P, S, Cl, Ar.

Question 24

What is the structure of sulphur?

Answer 24

Forms crown-shaped S8 molecules which are the largest molecules of Period 3 non-metals so there are more VdW forces between them.

Highest m+b points of P, S, Cl, Ar.

Question 25

What is the structure of chlorine?

Answer 25

Forms diatomic Cl2 molecules.

Although it has more electrons than P and S, the small size of its molecules means it has weaker VdW forces so lower m+b point.

Question 26

What is the structure of argon?

Answer 26

Monatomic, only weak VdW forces holding the atoms together so very weak m+b point.