DF4: Bond enthalpies

Question 1

Define “bond enthalpy”

Answer 1

The energy needed to break one mole of a bond to give separate atoms all in the gaseous state

Question 2

Define “average bond enthalpy”

Answer 2

The energy needed to break one mole of a bond to give separate atoms all in the gaseous state, averaged over many different compounds

Question 3

Are bond enthalpies always positive/negative and why?

Answer 3

Positive

Bond breaking is endothermic

Question 4

Is bond making exothermic/endothermic?

Answer 4

exothermic

Question 5

What is the relationship between bond strength and bond enthalpy and why?

Answer 5

The stronger a bond, the more energy needed to break it so the higher its bond enthalpy - greater attraction between nuclei and electron pair(s) in bonds

Question 6

Explain the forces which determine a bonds length in covalent molecules(3)

Answer 6

• The positive nucleui are attracted to the shared electrons so the atoms move together
• The two positive nucleui also repel each other, these forces become greater until the atoms stop moving together
• The distance between the 2 nucleui is the distance where the attractive and repulsive forces balance each other out : equilibrium bond length

Question 7

What is the relationship between bond length and bond enthalpy and why?

Answer 7

The shorter the bond length, the higher the bond enthalpy as there is a stronger attraction between the two atoms in the covalent bond.

Question 8

Order these types of bonds in order of increasing bond enthalpy:

• Double bond
• Single Bond
• Triple Bond

Answer 8

• Triple bond
• Double bond
• Single Bond

Question 9

Why do triple bonds have a higher bond enthalpy than double and single bonds?(2)

Answer 9

• 2 pairs of electrons are shared with a double bond so electron density between the two atoms is greater
• Therefore the positive nucleui are more attracted to the electrons and so the atoms move closer together

Question 10

What is the formula for bond enthalpy change?

Answer 10

Bond enthalpy change = total amount of bonds broken + total amount of bonds made

Question 11

What do all reactions initially need and why?

Answer 11

Activation Energy - to stretch and break bonds as some must be broken before product molecules can begin to form

Question 12

Why do some reactions such as neutralisation need barely be heated before they start?

Answer 12

They need only a little energy with enough available in their surroundings of room temperature

Question 13

Why do you not need to break all bonds before a reaction occurs?

Answer 13

Once a few bonds have been broken, new bonds start to form and this gives out enough energy to keep reactions going

Question 14

Why do some reactions need continuous heating?

Answer 14

They are only slightly exothermic or could be endothermic

Question 15

Why is there some variation between experimental and theoretical bond enthalpies? (2)

Answer 15

• The bond enthalpy value is not actual standard value, e.g as under standard conditions water is a liquid but when calculating bond enthalpies you must work in a gaseous state
• Bond enthalpies in theory are given out as averages of several different bonds from different compounds

Question 16

Why are bond enthalpies useful?

Answer 16

They enable enthalpy changes to be measured when there is little specific data for the compound

Question 17

Why is it difficult to measure bond enthalpy? (2)

Answer 17

• There is often more than one type of bond in a compound

- Difficult to make measurements when everything is in the gaseous state.

Question 18

How are bond enthalpies measured?

Answer 18

Indirectly using enthalpy cycles (Hess Cycles)