Bonding & Structure

Question 1

Define ionic bonding

Answer 1

Electrostatic attraction between oppositely charged ions

Question 2

Define metallic bonding

Answer 2

Electrostatic attraction between positive metal cations and delocalised electrons

Question 3

Define covalent bonding

Answer 3

Electrostatic attraction between protons in two nuclei and a shaired pair of electrons between them

Question 4

Explain why water has a low boiling point.

Answer 4

Weak forces between the H2O molecules (weak “intermolecular forces”)
Little energy is needed to overcome them.

Question 5

What is required for a substance to conduct electricity? Using this, explain why water cannot conduct electricity.

Answer 5

For a substance to conduct electricity, it must have freely moving charges. Although the H2O molecules can move, they are neutral.

Question 6

What does covalent bond strength depend on?

Answer 6

Bond length - shorter bonds, stronger bonds
Bond order (more e- pairs, stronger attraction)

Question 7

Define allotropes.

Answer 7

Forms of an element with different structures.

Question 8

What are the 4 main carbon allotropes?

Answer 8

Diamond
Graphite
Graphene
Nanotubes & fullerenes

Question 9

State the properties of diamond

Answer 9

V. high MP
Hard - good for cutting
Electrical insulator
Great conductor of sound/heat

Question 10

State the properties of graphite

Answer 10

High MP
Soft
Electrical conductor

Question 11

What does VSEPR stand for?

Answer 11

Valence-Shell Electron-Pair Repulsion

Question 12

What is the shape and bond angle around a molecule with two bond pairs?

Answer 12

Linear, 180*

Question 13

What is the shape and angle around an atom with three bond pairs?

Answer 13

Trigonal planar, 120*

Question 14

What is the shape and angle around an atom with four bond pairs?

Answer 14

Tetrahedral, 109.5

Question 15

What is the shape and angle around an atom with five bond pairs?

Answer 15

Trigonal bipyramid, 120*

Question 16

What is the shape and angle around an atom with six bond pairs?

Answer 16

Octahedral, 90*

Question 17

Explain why specific numbers of bond/electron pairs result in their respective angles & shapes.

Answer 17

Electron pairs separate as far as possible
To minimise electron pair repulsion

Question 18

What is the shape and angle around an atom with three bond pairs and one lone pair?

Answer 18

Trigonal pyramid, 107*

Question 19

What is the shape and angle around an atom with two bond pairs and two lone pairs?

Answer 19

V-shaped / Bent, 104-5*

Question 20

Why do the angles change when lone electron pairs are introduced?

Answer 20

Lone electron pairs repel bonding electron pairs more than bonding electron pairs repel bonding electron pairs.

Question 21

Why are some bonds polar?

Answer 21

A difference in electronegativity between atoms in a bond result in a polar bond.

Question 22

How can you tell experimentally if a molecule is polar?

Answer 22

Deflecting jet test:
polar liquids will deflect towards a charged rod

Question 23

What are the limitations to a delfecting jet test?

Answer 23

Doesn’t work with (s), (g) and (aq)
but… could dissolve into a non-polar solvent.

Question 24

What does the strength of london forces depend on?

Answer 24

Number of electrons per molecule - more e-, stronger attraction
Shape of molecule - closer packing, better attraction
Extent of delocalisation

Question 25

What are permanenet dipole-dipole forces?

Answer 25

Permanent dipole-dipole forces are the weak intermolecular forces of attraction that arise between permanently polar molecules. These forces are between the delta positive end of one polar bond with the delta negative end of another polar bond.

Question 26

What is hydrogen bonding?

Answer 26

Hydrogen bonding is electrostatic force of attraction between a hydrogen atom which is covalently bound to a more electronegative "donor" atom or group, and another electronegative atom bearing a lone pair of electrons—the hydrogen bond acceptor

Question 27

Why can hydrogen bonding only occur with certain elements?

Answer 27

Only occurs with nitrogen, oxygen or fluorine as the atom needs to be electronegative enough to make the X-H bond polar enough to expose the H nucleus.

Question 28

List different IMFs in order of strength (strongest to weakest)

Answer 28

Hydrogen bonding Dipole-dipoles London forces

Question 29

Define miscibility.

Answer 29

Miscibility is the property of two substances to mix in all proportions, forming a homogeneous mixture.

Question 30

How can you tell if two substances are miscible?

Answer 30

A substance is miscible if the bonds formed are equal to or greater than, in terms of strength, than the bonds broken.

Question 31

What affects the strength of ionic bonds and how?

Answer 31

Ionic radius - increasing radius decreases strength Ionic change - increasing change increases attraction

Question 32

What is formed when an atom losses or gains an electron? What process does this?

Answer 32

Ions are formed Cations from losing electrons in oxidation Anions from gaining electrons in reduction

Question 33

Describe the general properties of ionic compounds.

Answer 33

High melting and boiling points Soluble in polar solvents Electrically conductive when molten or dissolved in water

Question 34

Define electronegativity.

Answer 34

Tendency of an atom to attract bonding electrons

Question 35

Describe the different types of intermoleuclar forces and what affects the strength of them.

Answer 35

London forces (Instantaneous dp - induced dp) - greater electron density, increases strength Permanent dipoles - strength increases with increasing polarity in the molecule Hydrogen bonds - increasing charge increases bond strength

Question 36

How does hydrogen bonding arise in molecules such as H20, NH3 or liquid HF?

Answer 36

When bonded to hydrogen, oxygen and nitrogen and flourine have a large enough difference in electronegativity such that a polar bond is formed. This allows the hydrogen to be positively charged enough to form hydrogen bonds to lone pairs on other atoms.

Question 37

How does the strength of hydrogen bonds compared to other IMFs?What properties does this effect and how?

Answer 37

Hydrogen bonds are far stronger than other IMFs and therefore take more energy to break. This results in an increase in melting/boiling points for substances that contain hydrogen bonds.

Question 38

How do hydrogen bonds affect the density of water/ice?

Answer 38

When freezing ice, more hydrogen bonds form between molecules which causes the molecules to space out and therefore decrease the density of water.

Question 39

How does bonding vary with different shapes of chains and chain length in alkanes? How does this affect melting/boiling points?

Answer 39

As chain length increase, london forces become stronger therefore increasing melting/boiling points. Straighter thinner chains have stronger london forces compared to branched chains due to their ability to pack closer together.

Question 40

How does the boiling point vary between alcohols and alkenes with similar numbers of electrons? (e.g. methanol and ethane (both 18 e-))

Answer 40

Alcohols are polar and are able to form hydrogen bonds Both alcohols and alkenes have london forces DPDP and hydrogen bonds are much stronger therefore alcohols have higher boiling points

Question 41

Describe and explain the trend in boiling points between different hydrogen halides.

Answer 41

As the size of the halogen increases, the greater the boiling point due to the increase in london forces.