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IB Chemistry > Energetics > Flashcards

Energetics Flashcards

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1
Q

Definition of heat.

A

a form of energy that flows from something at a higher temperature to something at a lower temperature

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2
Q

Definition of temperature.

A

a measure of the average kinetic energy of particles

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3
Q

definition of an exothermic reaction

A

heat energy is transferred from a system (chemical reaction) to the surroundings — the surroundings get hotter.

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4
Q

definition of endothermic reaction

A

a system (chemical reaction) takes in heat energy from the surroundings — the surroundings get cooler

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5
Q

What is the enthalpy change of a system equal to?

A

the amount of heat taken in/given out in a chemical reaction

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6
Q

What is the symbol for enthalpy change?

A

∆H

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7
Q

What is the enthalpy change?

A

heat exchange with the surroundings at constant pressure.

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8
Q

How do you measure the enthalpy (H) of a system?

A

It is not possible. You can only measure the enthalpy change (∆H)

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9
Q

What is the ∆H for an exothermic and endothermic reaction?

A

Exo: negative
Endo: positive

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10
Q

Are the products more or less stable than the reactants in exothermic reactions? Why?

A

the products are more stable than the reactants.

the enthalpy of products is less than that of the reactants

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11
Q

Are the products more or less stable than the reactants in endothermic reactions? Why?

A

the products are less stable than the reactants.

the enthalpy of products is greater than that of the reactants

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12
Q

total energy of the reactants=

A

total energy of the products + heat given out

(exothermic)

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13
Q

total energy of products=

A

total energy of reactants + heat taken in

(endothermic)

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14
Q

Heat energy change (Q) =

A

mass (m) (of what is heated) x specific heat capacity (c) x temp. change (∆T)

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15
Q

What is the specific heat capacity of water?

A

4.18 J g^-1 K^-1

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16
Q

enthalpy change (∆H) =

A

n

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17
Q

Why would a value in an experiment be less exothermic than the literature value?

A

some heat dissipates to surroundings — not all heat energy from burning of fuel goes to heating up the water.

Incomplete combustion gives out less energy

energy required to heat calorimeter (copper can)

water evaporating

fuel evaporating

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18
Q

How would you get more accurate values in “a molar enthalpy change when a liquid hydrocarbon or alcohol is burnt”?

A

use a bomb calorimeter

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19
Q

Draw the exothermic reaction: CH4 (g) + 2 O2 (g) –> CO2 (g) + 2 H2O (l) ∆H= -890 kJ mol^-1

A

see page 2

| reactants

reactants
|___________________
| CH4 (g) + 2 O2 (g).
| 890 kJ mol^-1
| .
| . products
| __________________
| CO2 (g)+2 H2O (l)
|________________________________________

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20
Q

Draw the endothermic reaction: N2 (g) + O2 (g) –> 2 NO (g) ∆H= +180 kJ mol^-1

A

see page 2

| _______________________

products
| __________________
| . 2 NO (g)
| 180 kJ mol^-1 .
| .
| reactants .
| __________________
| N2 (g) + O2 (g)
|________________________________________

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21
Q

What are enthalpy changes of neutralisation (∆Hn)? Is is it exorthermic or endothermic?

A

enthalpy changes when 1 mole of H2O molecules are formed when an acid (H+) reacts with an alkali (OH-) under standard conditions.

e.g: H+ (aq) + OH- (aq) –> H2O (l)

always exothermic

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22
Q

What are enthalpy changes of solution (∆Hsol)

A

the enthalpy change when one mole of solute is dissolved in excess solvent to form a solution of “infinite dilution” under standard condition.

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23
Q

What is the pressure and temperature under standard conditions?

A

100 kPa
298.15 K / 25°C

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24
Q

What is the standard enthalpy change of reaction (∆Hr)? Give an example.

A

The enthalpy change when molar amounts of reactants as shown in the stoichiometric equation react together under standard conditions to give products.

N2 (g) + 3 H2(g) –> 2 NH3 (g) ∆H= -92 kJ mol^-1

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25
Q

What is the standard enthalpy change of combustion (∆Hc)? Give an example.

A

the enthalpy change when one mole of a substance is completely burnt in oxygen under standard conditions.

e.g: CH₄ (g) + 2O₂ (g) –> CO₂ (g) + 2H₂O (l) ∆H- -890 kJ mol⁻¹

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26
Q

What is the standard enthalpy change of formation (∆Hf)? Give an example.

A

the enthalpy change when one mole of the substance is formed from its elements in their standard states under standard conditions.

e.g: C (s) + 2H₂ (g) –> CH₄ (g) ∆Hf= -75 kJ mol⁻¹

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27
Q

What is the ∆Hf for any element in its standard state?

A

0

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28
Q

What is the standard enthalpy change of formation of CH₃COOH (l)?

A

2C (s) + 2H₂ (g) + O₂ (g) –> CH₃COOH (l)

29
Q

What is Hess’s law?

A

The enthalpy change accompanying a chemical reaction is independent of the pathway between the initial and final states.

30
Q

How to find the enthalpy change from enthalpies of combustion?

A

∆H = ∑∆Hc(reactants) - ∑∆Hc(products)

31
Q

How to find the enthalpy change from enthalpies of formation?

A

∆H = ∑∆Hf(products) - ∑∆Hf(reactants)

32
Q

Definition of bond enthalpy. Give an example.

A

the enthalpy change when one mole of covalent bonds, in a gaseous molecule, are broken under standard conditions.

e.g: H₂ (g) –> 2H (g) ∆H= +436 kJ mol⁻¹

(can only be calculated for substances in the gaseous state)

33
Q

does bond breaking require or release energy?

A

requires

34
Q

does bond making require or release energy

A

releases

35
Q

Why are the bond energies given in tables only average values? Why could this be a problem?

A

they are slightly different with all compounds.

Can give inaccuracies in calculations.

36
Q

∆Hr =

A

∑bonds broken - ∑bonds made

37
Q

In bond energy, when will a reaction (in the gas phase) be exothermic?

A

more energy is released when bonds are formed than is required to break bonds.

38
Q

In bond energy, when will a reaction (in the gas phase) be endothermic?

A

if less energy is released when bonds are formed than is required to break bonds.

39
Q

is breaking bonds endothermic or exothermic?

A

endothermic

40
Q

is making bonds endothermic or exothermic?

A

exothermic

41
Q

What is the standard enthalpy change of atomisation (∆Hat)? Give an example.

A

the enthalpy change when one mole of gaseous atoms is formed from the element under standard conditions.

e.g: ½H₂ (g) –> H (g) ∆Hat= +218 kJ mol⁻¹

42
Q

Is the standard enthalpy change of atomisation exothermic or endothermic?

A

endothermic

43
Q

How is the standard enthalpy change of atomisation linked to bond energies?

A

∆Hat values are half the bond energy values.

44
Q

What is the first ionisation energy?

A

the enthalpy change when 1 electron is removed from each atom in 1 mol of gaseous atoms under standard conditions.

[must have the (g) symbols]

45
Q

Write the first ionisation energy (use M).

A

M (g) –> M⁺ (g) + e⁻

46
Q

Write the second ionisation energy (use M).

A

M⁺ (g) –> M²⁺ (g) e⁻

47
Q

Write the sum of the first and second ionisation energies (use M).

A

M (g) –> M²⁺ (g) 2e⁻

48
Q

What is the first electron affinity?

A

enthalpy change when 1 electron is added to each atom in 1 mol of gaseous atoms under standard conditions.

49
Q

Write the first electron affinity (use X)

A

X (g) + e⁻ –> X⁻ (g)

50
Q

Is the first electron affinity exothermic or endothermic? Why?

A

exothermic

it is a favourable process to bring an electron from infinity to where it feels the attractive force of the nucleus in an atom.

51
Q

Is the second electron affinity exothermic or endothermic? Why?

A

endothermic

it is an unfavourable process to add an electron to an ion which is already negatively charged, due to repulsion between the negative charges.

52
Q

Write the second electron affinity (use O).

A

O⁻ (g) + e⁻ –> O²⁻ (g)

53
Q

What is the lattice enthalpy (∆Hlatt)? Give an example.

A

the enthalpy change when 1 mol of ionic compound is broken apart into its constituent gaseous ions under standard conditions.

NaCl (s) –> Na⁺ (g) + Cl⁻ (g) ∆Hlatt = +771 kJ mol⁻¹

54
Q

When is lattice enthalpy exothermic and endothermic?

A

if it breaks the lattice, it is endothermic.
if it makes the lattice, it is exothermic.

(I think)

55
Q

What order does Born-Haber Cycles go in?

A

enthalpy change of formation ∆Hf
enthalpy change of atomisation ∆Hat
1st ionisation energy
(2nd ionisation energy…)
enthalpy change of atomisation ∆Hat
1st electron affinity
(2nd electron affinity…)
lattice enthalpy ∆Hlatt

56
Q

What factors does lattice enthalpy rely on?

A

charge on ions
size of ions (ionic radii)

57
Q

What happens if there is a higher charge on ions?

A

the ions will have a greater electrostatic attraction to each other, and therefore the lattice enthalpy will be higher (more energy required to separate ions)

58
Q

What happens to the lattice enthalpy if the ions are small?

A

the attraction is greater, and so is the lattice enthalpy.

59
Q

What does an exothermic reaction involve?

A

a decrease in the enthalpy of the system.

60
Q

What does “more stable” mean?

A

lower energy

61
Q

Are products or reactants more stable than each other?

A

products more stable than reactants.

62
Q

Why is graphite more stable than diamond?

A

it is at a lower enthalpy

63
Q

Why doesn’t a mixture of methane and oxygen spontaneously combust when they are both unstable?

A

they need activiation energy

64
Q

what is activation energy?

A

the minimum energy that colliding species must have before collision results in chemical reaction.

65
Q

Draw an exothermic and endothermic energy profile.

A

See page 26

66
Q

When there is higher activation energy, is the reaction slower or faster?

A

slower

67
Q

What does how exothermic/endothermic a reaction is tell us about how quickly the reaction occurs?

A

is doesn’t tell us anything.

68
Q
A

IB Chemistry (6 decks)

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