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CLS3003-B: Chemistry (Semester 2) > L1 - Chemical Reactivity > Flashcards

L1 - Chemical Reactivity Flashcards

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1
Q

Define ‘energy’.

A

The ability to do work.

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2
Q

Define ‘work’.

A

The distance moved against an opposing force.

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3
Q

What is the equation for work?

A

Work = Force x Distance

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4
Q

What units is energy measured in?

A

Joules (J)

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5
Q

Define ‘1 Joule’.

A

The amount of energy required to raise a 1kg substance 10cm against the force of gravity.

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6
Q

Define ‘1 Calorie’.

A

The amount of heat necessary to raise the temperature of exactly one gram of water by one degree Celsius.

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7
Q

What are the conversions into kcal and cal for 1 calorie.

A

1 calorie = 1 kcal = 1000 cal

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8
Q

Describe potential energy.

A
  • Stored energy

- Energy due to position

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9
Q

What is the equation for potential energy?

A

PE = mgh

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10
Q

Describe kinetic energy.

A
  • Energy of motion

- Depends on mass and velocity

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11
Q

What is the equation for kinetic energy?

A

KE = 1/2mv^2

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12
Q

Describe chemical energy.

A
  • Energy stored in bonds
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13
Q

State the 5 different types of energy.

A
  • Potential energy
  • Kinetic energy
  • Electromagnetic energy
  • Nuclear energy
  • Chemical energy
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14
Q

State the ‘First Law of Thermodynamics’ / Law of Conservation of Energy.

A

“Energy cannot be created or destroyed by any physical and chemical changes - it can only be converted from one form to another.”

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15
Q

What are the conditions for chemical reactions to occur?

A
  • Reactants energetic

- Reactants oriented correctly

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16
Q

Energy is either ________ or ________.

A

Energy is either released or absorbed.

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17
Q

What is the change in energy?

A

The difference between bond energies of reactants and products.

Enthalpy change of reaction (ΔH)

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18
Q

Define ‘enthalpy’.

A

A measure of the heat content of a substance at constant pressure.

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19
Q

What does ΔH° mean?

A

Heat released / absorbed during a chemical reaction at standard conditions

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20
Q

What is the equation for ΔH°?

A

ΔH° = H.products - H.reactants

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21
Q

What can you not measure?

A

Cannot measure the actual enthalpy of a substance (but can measure enthalpy change)

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22
Q

What is the standard state of an element?

A
  • 1 atmosphere

- 25°C

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23
Q

What is the standard state of a compound?

A
  • Gaseous substance = 1 atmosphere
  • Pure substance in a condensed state (liquid / solid) = state it would have at 1 atmosphere
  • Substance in solution = concentration of 1M
24
Q

What is used to indicate physical state?

A

Subscripts:

  • (g)
  • (l)
  • (s)
25
Q

Describe the 3 characteristics of exothermic reactions.

A
  • Negative ΔH
  • Heat given out
  • Enthalpy of reactants > products
26
Q

Describe the 3 characteristics of endothermic reactions.

A
  • Positive ΔH
  • Heat absorbed
  • Enthalpy of products > reactants
27
Q

What will happen to the temperature of an exothermic reaction?

A

Temperature of the system will be observed to rise.

28
Q

What will happen to the temperature of an endothermic reaction?

A

Temperature of the system will fall (in practice, reactants will be heated to speed up the reaction and provide the absorbed heat).

29
Q

What happens during bond breaking?

A

Energy is added

30
Q

What happens during bond making?

A

Energy is released

31
Q

Define ‘entropy’.

A

Measures the amount of energetic disorder in a system.

32
Q

When is entropy higher?

A

When the disorder / randomness of particles in a substance / mixture are greater.

33
Q

What is the symbol of entropy?

A

S

34
Q

What units is entropy measured in?

A

J mol^-1 K^-1

35
Q

What is the equation for ΔS?

A

ΔS = S.final - S.initial

36
Q

What is the equation for ΔS.total?

A

ΔS.total = ΔS.system + ΔS.surroundings

37
Q

State the ‘Second Law of Thermodynamics’.

A

“Entropy tends to a maximum.”

38
Q

What do all spontaneously occurring chemical and physical changes involve?

A

An overall increase in entropy.

39
Q

Describe scenarios in which entropy is increased.

A
  • Solids melting
  • Liquids boiling
  • Number of molecules increasing
  • Ionic solids dissolving
  • Temperature increasing
40
Q

Define ‘Gibbs Free Energy’.

A

Energy from a reaction free to do work.

41
Q

What is the equation for ΔG?

A

ΔG = ΔH - TΔS

42
Q

What happens if ΔG < 0?

A

Reaction will be spontaneous

43
Q

What happens if ΔG > 0?

A

Reaction needs energy input to occur

44
Q

What happens if ΔG = 0?

A

System is in equilibrium

45
Q

What is the quantity of ΔG if the reaction is spontaneous?

A

ΔG < 0

46
Q

What is the quantity of ΔG if the reaction needs energy input to occur?

A

ΔG > 0

47
Q

What is the quantity of ΔG if the system is in equilibrium?

A

ΔG = 0

48
Q

What happens during catabolic reactions?

A

High energy compounds → Simple molecules

49
Q

What happens during anabolic reactions?

A

Simple subunit → Complex molecule

50
Q

How do anabolic reactions take place?

A

In steps, coupled with ATP hydrolysis or another exergonic reaction.

51
Q

State the ΔG of the following reaction:

glucose + 6O2 → 6CO2 + 6H2O

A

ΔG = -2870 kJ mol^-1

52
Q

State the ΔG of the following reaction:

ATP + H2O → ADP + Pi

A

ΔG = -30.5 kJ mol^-1

53
Q

State the ΔG of the following reaction:

glucose + fructose → sucrose

A

ΔG = +29.3 kJ mol^-1

54
Q

State the ΔG of the following reaction:

glucose + ATP → glucose,p + ADP

A

ΔG = -16.7 kJ mol^-1

55
Q

State the ΔG of the following reaction:

fructose + ATP → fructose,p + ADP

A

ΔG = -14.2 kJ mol^-1

56
Q

State the ΔG of the following reaction:

glucose,p + fructose,p → sucrose + 2Pi

A

ΔG = -0.8 kJ mol^-1

CLS3003-B: Chemistry (Semester 2) (12 decks)

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