Part 1- Periodicity, trend, bonding

Question 1

What is the name if group 1 in the periodic table?

Answer 1

Alkali metals

Question 2

What is the name of group 2 in the periodic table?

Answer 2

Alkali Earth metals

Question 3

What is the name of group 7 in the periodic table?

Answer 3

Halogens

Question 4

What is the name of group 0/8 in the periodic table?

Answer 4

Noble gases

Question 5

Describe metallic bonding

Answer 5

Positive charged ions with delocalised electrons.

Question 6

Explain why metals have high melting and boiling point.

Answer 6

STRONG ATTRACTION force between these positive ions and delocalised electrons. A lot of energy is required to break metallic bonds.

Question 7

What are monoatomic elements

Answer 7

Consist of single atoms, they are not bonded, all noble gases are monoatomic elements

Question 8

Why monoatomic elements have a very low melting and boiling point?

Answer 8

Monoatomic elements can form very weak forces between the atoms when cooled. These forces are called London Dispersion Forces and are very weak. They can be easily broken, therefore monoatomic element have a very low MP and BP.

Question 9

Which electrons can move about easily in the metallic lattice?

Answer 9

The electrons from the outershell.

Question 10

What happen to the melting point of alkalis metals going down the group? (use databook) explain why.

Answer 10

Melting point decrease as we go down group 1. Because the delocalised electrons that pull the positive ion together is further away from the nucleus (due to an increase in energy level) so the attraction between the delocalised electron and the positive nucleus decrease as we go down a group.

Question 11

Which type of bonding is holds the atoms together in solid Neon?

Answer 11

London dispersion forces

Question 12

At what Kelvin temperature does the solid lattice Neon break up and turn liquid? (Note: 273 K = 0 °C) (use databook)

Answer 12

At -246°C Neon goes from Gas to Liquid. -246°C = 27K

Question 13

What state of matter would Argon be at -187 °C

Answer 13

Liquid

Question 14

What do the δ+ and δ- signs stand for?

Answer 14

Slightly positive and slightly negative

Question 15

What elements can form covalent network?

Answer 15

Carbon, Boron and Silicon

Question 16

Give an example of a colalent network

Answer 16

Diamond, Graphite, Boron and Silicon

Question 17

Why covalent network have high melting point and boiling point?

Answer 17

Structures consist of many thousands of atoms joined together by covalent bonds. High MP and BP as strong covalent bonds must be broken for the solid to melt/boil.

Question 18

What are the diatomic molecules in the first 20 elements of the periodic table?

Answer 18

H₂; N₂; O₂; F₂; Cl₂

Question 19

What is a diatomic molecule?

Answer 19

2 atoms joined together by covalent bond.

Question 20

What is the definition of Covalent Radii?

Answer 20

Is a measure of the size of an atom. It is half the distance between the nuclei of 2 covalently bonded atoms.

Question 21

What happen to the atomic radius going down a group and why?

Answer 21

In groups, covalent radius increases going down the group. More electron layers = higher covalent radius

Question 22

What happen to the atomic radius across periods and why?

Answer 22

In a period, covalent radius decreases going across the period. All atoms have the same number of electron layers but the number of protons increases across the period. Greater positive charge = greater attraction for electrons, the electron shells are pulled closer => Covalent radius is smaller.

Question 23

Give the definition of ionisation energy.

Answer 23

Is the energy required to remove 1 mole of electrons from one mole of gaseous atom.

Question 24

Why the second ionisation energy of sodium is much higher than the first?

Answer 24

Because we need to break into a new, complete electron shell which is closer to the nucleus. Attraction between the nucleus is therefore greater.

Question 25

What happen to the ionisation energy going down a group and why?

Answer 25

it decreases because there are extra electron shells meaning that the electron (-ve) being removed is further from the nucleus (+ve) so the force of attraction is lower.

Question 26

What happen to the ionisation energy across periods and why?

Answer 26

It increases. The electron is being removed from the same electron shell. Across the period there are more protons (+ve) and the covalent radius is decreasing. The electron (-ve) being removed is closer to the nucleus, which has more protons (+ve) so the force of attraction is higher.

Question 27

Explain why only the 1^st ionisation energy is quoted for hydrogen.

Answer 27

Hydrogen have only one electron it can get rid of.

Question 28

Would the melting point of Li be higher or lower that that of sodium? Why?

Answer 28

The atoms in a metal are held together by the attraction of the nuclei to electrons which are delocalized over the whole metal mass. As the atoms increase in size, the distance between the nuclei and these delocalized electrons increases; therefore, attractions fall. The atoms are more easily pulled apart to form a liquid, and then a gas. Li has a smaller covalent radius than Na, so it has more attraction to the nuclei, therefore stronger force and more energy is needed to break these forces. Therefore, Li will have a higher MP than Na.

Question 29

Xenon was the first noble gas to be made into a compound, explain why this would be the case.

Answer 29

Xenon have electrons not as tightly bond to the nuclei as the other noble gas, with smaller atomic number. Its inner electrons screen the outer electrons from the nucleus.

Question 30

A potassium ion is larger than a sodium ion because….

Answer 30

it has an extra shell of electrons.

Question 31

Explain why the Al³⁺ ion is much smaller than the P^3- ion even though they are very close to each other in the periodic table.

Answer 31

Al³⁺ have 13 protons and 10 electrons and therefore will have 2 shell of electrons, P^3- have 15 protons and 18 electrons and therefore will have 3 shells of electrons, therefore P^3- will be larger than Al^3+.

Question 32

Explain why potassium has a lower first ionisation energy than lithium.

Answer 32

K has a nuclear charge of 19+ and electron arrangement 2,8,8,1 Li has a nuclear chare of 3+ and electron arrangement 2,1 Potassium has two extra filled layers of electrons compared with lithium. These extra layers screen the outer electrons from the nucleus. Consequently, the outer electron in K is further from the nucleus and is not as strongly attracted. It therefore takes less energy to remove the outer electron from K than it does to remove the outer electron from Li.

Question 33

Why there is no electronegativity value provided for Argon?

Answer 33

Argon is a noble gas and do not form bonds.

Question 34

Al and P are close in the periodic table but P^3- ion is much larger than Al³⁺ ion, why?

Answer 34

P need to gain 3 electrons to become P^3- and will have an electron arrangement: 2,8,8 while aluminium need to loose 3 electrons to become Al³⁺ and have an electron arrangement 2.8. P^3- has an extra shell of electrons compare to Al³⁺ therefore it's larger.

Question 35

The P^3- ion and the Ca²⁺ ion have the same electron arrangement but Ca²⁺ ion is smaller the P^3- ion, why?

Answer 35

Calcium ion has a higher nuclear charge, so the electrons are pulled closer to the nucleus, making the calcium ion smaller than the phosphorus ion.

Question 36

What is the bonding in the first 20 elements?

Answer 36

**metallic**: Lithium, beryllium, sodium, magnesium, aluminium, potassium, calcium **covalent molecular**: hydrogen, nitrogen, oxygen, fluorine, chlorine, phosphorus (P₄), sulfur (S₈) carbon (fulerenes) **Covalent network:** Boron, carbon (graphite and diamond), silicon **monoatomic:** helium, neon, argon

Question 37

What are the trends on covalent radius across a period and down a group?

Answer 37

Across a period: decrease due to increaseing nuclear charge. Down a group: increase due to shielding effect.

Question 38

What is the first ionisation energy?

Answer 38

The energy required to remove one mole of electrons from one mole of gaseous atom. e.g: Cl**(g)** --\> Cl⁺**(g)** + e^- PS: state symbol **(g)** is required.

Question 39

What is the second ionisation energy?

Answer 39

The removal of one mole of electrons from one mole of positive ions in gaseous state. e.g: Cl⁺(g) --\> Cl²⁺(g) + e^-

Question 40

What are the trends in ionisation energy down a group and across a period?

Answer 40

Across period: increases due to increased in nuclear charge Down a group: decreases due to shielding effect

Question 41

What is electronegativity?

Answer 41

Electronagetivity is measure of the attraction an atom has for elecrons of a bond

Question 42

What are the trend in electronegativity down a group and across a period?

Answer 42

Across a period: increased due to increased in nuclear charge Down a group: decreased due to shielding effect

Question 43

What is a covalent bond?

Answer 43

Two positive nuclei held together by common attraction for shaired pair of electrons

Question 44

What is a polar covalent bond?

Answer 44

a covalent bond where the electronegativities of the atoms are different, giving rise to a dipole e.g