Module 2

Question 1

Define relative atomic mass

Answer 1

Relative atomic mass is the weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of carbon 12

Question 2

Define relative isotopic mass

Answer 2

Relative isotopic mass is the mass of an isotope relative to 1/12th of the mass of an atom of carbon 12

Question 3

What are isotopes

Answer 3

Isotopes are atoms of the same element with diff no. of neutrons and different masses

Question 4

What is meant by 0.5 mol dm-3 HCL?

Answer 4

0.5 mol HCL is dissolved in each 1dm3 of solution

Question 5

What can you expect to observe when a carbonate reacts with an acid?

Answer 5

• bubbles

- sold dissolves

Question 6

Why volume of CO2 is likely to be less than calculated?

Answer 6

-CO2 is slightly soluble in water

Question 7

Comment on low atom economy in terms of sustainability

Answer 7

• low atom economy means a lot of waste
• use for other product needs to be developed to increase atom economy
• alternative process needs to be developed with higher atom economy

Question 8

How do you know if it is a neutralisation reaction?

Answer 8

An acid is neutralised by a base to form water

Question 9

Explain reduction and oxidation in terms of e- and oxidation numbers

Answer 9

• reduction is a gain of e- and a decrease in oxidation no.

- oxidation is a loss of e- and an increase in oxidation no.

Question 10

Why is —— an example of a salt?

Answer 10

-the H+ ion in acid has been replaced by a metal ion

Question 11

Explain how the hydroxide ion acts as a base when it neutralises an acid

Answer 11

Accepts H+

Question 12

What is the ionic charge for phosphide?

Answer 12

3-

Question 13

What is the ionic charge for phosphate?

Answer 13

PO4 3-

Question 14

What is the ionic charge for sulfite?

Answer 14

SO3 2-

Question 15

What is the ionic charge for Manganate (VII)

Answer 15

MnO4-

Question 16

Explain how ionic bonding holds together particles in an ionic compound

Answer 16

Electrostatic attraction between oppositely charged ions

Question 17

Explain the electrical conductivities when ionic compound solid and when dissolved in water

Answer 17

• when solid, ions not free to move as fixed in giant ionic lattice structure + thus cannot carry charge
• when in solution, ions free to move + can carry charge

Question 18

Explain the pattern in melting points and ionic compounds

Answer 18

• as ionic charge increases, melting point increases

- greater the ionic charge, greater the attraction between ions + the more energy needed to break the ionic bonds

Question 19

Explain in terms of structure and bonding why MgO has a high melting point

Answer 19

• MgO has a giant ionic lattice
• there is strong binding between 2+ and 2- ions
• large amount of energy needed to break ionic bonds

Question 20

What is a covalent bond

Answer 20

The strong electrostatic attraction between a shared pair of e- and the nuclei of the bonded atoms

Question 21

What is meant by isoelectronic?

Answer 21

X and Y have the same no. of e-

Question 22

X molecules are polar. Why?

Answer 22

• x y bonds are polar

- molecule is not symmetrical + dipoles do not cancel

Question 23

Origin of London forces?

Answer 23

• movement of e- produces changing dipole in molecule, creating instantaneous dipole
• instantaneous dipole induces dipole on neighbouring molecule
• induced dipole induced further drilled on neighbouring molecules which then attract one another

Question 24

Define electronegativity

Answer 24

-ability of an atom to attract e- in a covalent bond

Question 25

What is a hydrogen bond?

Answer 25

An attraction between a lone pair of e- on an electronegative atom in one molecule and a H atom in another molecule attached to an electronegative atom

Question 26

What are water molecules polar?

Answer 26

- O more electronegative than H | - O attracts bonded pair of e- in covalent bond between O and H more than H

Question 27

State and explain properties of ice that are a result of hydrogen bonding

Answer 27

- Ice is less dense than water. Because H20 molecules in ice are held apart by H bonds in open lattice structure - ice has higher melting points than expected. Because H bonds are stronger than other intermolecular forces so more energy needed to break H bonds

Question 28

Explain how scientists are able to predict the shape of a molecule from no. Of e- pairs surrounding central atom.

Answer 28

- pairs of e- surrounding central atom repel | - shape determined by no. of bond pairs + no. Of lone pairs

Question 29

Bonding and structure of graphite?

Answer 29

- giant covalent structure - hexagonal layers - good conductor- delocalised e- - high mp- strong covalent bonds - soft- layers can slide

Question 30

How London forces arise?

Answer 30

- uneven distribution of e- - causes an instantaneous dipole - causes induced dipoles on neighbouring molecules

Question 31

Anhydrous?

Answer 31

No waters of crystallisation

Question 32

Why chlorine shouldn’t be added to water?

Answer 32

-chlorinated hydrocarbons are carcinogens

Question 33

Metallic bonding?

Answer 33

Electrostatic attraction between positive ions and delocalised e-

Question 34

BF3 has polar bonds but molecules are non polar. Why?

Answer 34

- symmetrical | - dipoles cancel each other out

Question 35

Ionic lattice?

Answer 35

-repeating pattern of oppositely charged ions

Question 36

Ice has higher melting point than expected?

Answer 36

-H bonds are strong and need to be broken

Question 37

Diamond structure?

Answer 37

Giant covalent