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Chemistry > 9 - Enthalpy > Flashcards

9 - Enthalpy Flashcards

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1
Q

enthalpy H

A

the measure of heat energy in a chemical system
-> cannot be measured but enthalpy changes can

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2
Q

enthalpy change ΔH

A

ΔH= H(products) - H(reactants)

+ = products have more energy thanreactnts

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3
Q

conservation of energy

A

energy cannot be created or destroyed

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4
Q

where is heat transferred between in a reaction

A

the system and surroundings
the system- reactants and products
surroundings- apparatus, the lab
universe= system + surrounding

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5
Q

exothermic

A

ΔH is negative
chemical system loses energy
surroundings gain energy
temp of surroundings increase

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6
Q

endothermic

A

ΔH is positive
chemical system gains energy
surroundings loses energy
temp of surroundings decreases

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7
Q

activation energy

A

minimum energy required for a reaction to take place

-> the energy required to break bonds acts as an energy barrier to the reaction

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8
Q

standard conditions

A

100 kPa
298 K (25’c)
1 moldm-3
standard state = the physical state of a substance under standard conditions

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9
Q

standard enthalpy change of reaction ΔrH°

A

enthalpy change that accompanies a reaction in the molar quantities shown in a chemical equation under standard conditions and states

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10
Q

ΔrH° of Mg and O2

A

Mg (s) + 1/2O2(g) -> MgO (s) ΔrH°= -602

or

2Mg (s) + O2(g) -> 2MgO (s) ΔrH°= -602 x 2= -1204

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11
Q

standard enthalpy change of formation ΔfH°

A

enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions and states

NOTE; writing this formula must show that one mole of the substance is produced

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12
Q

ΔrH° of MgO(s)

A

Mg (s) + 1/2O2(g) -> MgO (s) ΔfH°= -602

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13
Q

standard enthalpy change of combustion ΔcH°

A

the enthalpy change that takes place when one mole of a substance reacts completely with oxygen under standard conditions and states

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14
Q

ΔcH° of butane

A

C4H10 + 13/2 O2 -> 4CO2 + 5H20

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15
Q

standard enthalpy change of neutralisation ΔneutH°

A

enthalpy change when an acid and a base react to form one mole of H2O(l) under standard conditions and states

  • this value is the same for all neutralisation reactions
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16
Q

Calculating an energy change - the mass of the surroundings

A

m = of the mass that changes temperature in grams

NOTE - not the mass of the reactants

17
Q

Calculating an energy change - specific heat capacity

A

the energy required to raise the temperature of 1g of a substance by 1K
-> in most experiments we will be measuring c for water as this is usually the substance that will have a temp change
c= 4.18 Jg-1K-1

18
Q

Calculating an energy change - temperature change

A

ΔT=T(final) - T(initial)

19
Q

formula for heat energy

A

q = mcΔT
m in grams
ΔT - doesn’t matter as long final and initial are the same units

unit of q usually calculated is J as C= Jg-1K-1
J-> KJ divide by 1000

20
Q

method to determine ΔcH° of methanol

A

CH3OH(l) + 3/2O2(g) -> CO2(g) + H2O(l)

Materials and Apparatus:

Methanol (fuel)
Spirit burner- wick and methanol
Thermometer
Measuring cylinder
clamp
Water
Calorimeter (insulated container)

Experimental Setup:

Set up the spirit burner on a stable surface.
Measure the initial mass of the spirit burner (with methanol) and record it.
Measure a known volume of water using the measuring cylinder and pour it into the calorimeter.
Measure the initial temperature of the water in the calorimeter and record it.

Combustion of Methanol:

Ignite the spirit burner containing methanol and let it burn
Stir the methanol with the thermometer

Temperature Measurement:

After the combustion, measure the final temperature of the water in the calorimeter and record it.
Measure the weight of the spirit burner

use q=mcΔT to calculate energy change of the water

then calculate the mole of CH3OH burnt

calculate ΔcH
->ΔcH is the enthalpy change of the complete combustion of 1 mole of methanol
-> in the experiment x mol of methanol transfers y KJ of energy to the water
-> the water gained y KJ of energy from the combustion of x of methanol
->therefore 1 mole of methanol has lost y/x kj of energy
-> so ΔcH of methanol is -y/x

NOTE= ΔcH is negative as system has lost energy due to the combustion of methanol (exothermic0

21
Q

reasons why the experimental ΔcH is not accurate

A
  • heat loss to the surroundings other than water. such as the beaker and the air surrounding the flame
  • incomplete combustion of methanol- black layer of soot
  • evaporation of methanol from the wick - the burner must be weight soon as possible after extinguishing the flame otherwise methanol may have evaporated
  • non-standard conditions

all but the last reasons would lead to a value of ΔcH that is less exothermic

22
Q

average bond enthalpy

A

the average energy required to break a particular type of bond in a gaseous molecule, averaged over a range of different compounds

23
Q

bond enthalpies are always exo or end?

A

endothermic
-> bond enthalpies is breaking bonds which requires energy
-> always a positive value

24
Q

to break bonds energy is __________

A

required

25
Q

to form bonds energy is __________

A

released

26
Q

limitations of average bond enthalpies

A

actual bond enthalpies can vary depending on its chemical environment

27
Q

calculate ΔrH°

A

Σ(bond enthalpies in reactants) - Σ(bond enthalpies in products)

28
Q

calculations using average ______ _______ need all species to be _________ molecules. If you produce H2O(g) rather than H2O(l), this means that the calculate _____ is not a ______ enthalpy change.

what do you do to overcome this?

A

bond
enthalpies
gaseous
ΔrH°
standard

work out the standard enthalpy change but consider the enthalpy change for H2O(g) condensing into H2O(l)

29
Q

Hess’ Law

A

is a reaction can take place by two routes, and the starting and finishing conditions are the same, the total enthalpy change is the same for each route

30
Q

why is Hess’ law helpful

A

enthalpy changes of reactions are often hard to determine directly

31
Q
A

Chemistry (30 decks)

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