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Chemistry > 1 Chemical reactivity > Flashcards

1 Chemical reactivity Flashcards

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1
Q

Define ‘energy’

A

the ability to work

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2
Q

Define ‘work’

A

the distance moved against an opposing force

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3
Q

What are both work and energy measured in?

A

Joules

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4
Q

What is 1 joule the same as?

A

the amount of energy required to raise a 1kg substance 1 cm up against the force of gravity

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5
Q

What is 1 calorie the same as?

A

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

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6
Q

What is one food calorie equivalent to?

A

1 Kcal or 1,000 calories

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7
Q

What is potential energy?

A

the stored energy of position

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8
Q

State the formula for potential energy.

A

PE = mgh

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9
Q

State the formula for kinetic energy

A

KE = 1/2mv^2

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10
Q

What is ‘electromagnetic’ energy?

A

a form of energy that can be reflected/emitted from objects through electric or magnetic waves travelling through space

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11
Q

What is ‘nuclear’ energy?

A

the energy released during nuclear fission or fusion, especially when used to generate electricity

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12
Q

What is ‘chemical’ energy?

A

energy stored in bonds

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13
Q

State the ‘first law of thermodynamics.’

A

energy cannot be created or destroyed by any physical or chemical changes, it can only be converted from one form to another (except nuclear)

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14
Q

State and explain Einstein’s equation of special relativity related to the ‘first law of thermodynamics.’

A

e = mc²

if matter is destroyed, energy is created, and if energy is destroyed, mass is created

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15
Q

State 2 energy transformations

A

chemical, motion, radiant, chemical

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16
Q

How do reactions occur?

A

By interaction of molecules

Collision theory: reactions occur when particles collide, BUT not all collisions lead to a reaction

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17
Q

Which 2 conditions must reactants meet in order for a chemical reaction to occur?

A

must be energetic

and orientated correctly

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18
Q

State the 3 main stages of a successful chemical reaction

A
  1. collision between reactant particles
  2. breaking of chemical bonds
  3. right orientation of reactants will lead to a new product
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19
Q

What happens to energy in chemical reactions?

A

It is either released or absorbed

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20
Q

What is the ‘enthalpy change of reaction, ΔH?’

A

change in energy which is the difference between bond energies of reactants and products

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21
Q

What is meant by ‘enthalpy?’

A

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

22
Q

What type of energy is enthalpy and how can it be measured?

A

It is internal energy

the actual enthalpy of a substance cannot be measured but the enthalpy change can

23
Q

Define ‘enthalpy change.’

A

heat released or absorbed during a chemical reaction at standard conditions

24
Q

State the formula for enthalpy change.

A

ΔH = ΔH(products) - ΔH(reactants)

25
Q

Why might enthalpy values vary?

A

according to conditions

26
Q

What are standard conditions for an element?

A

form the element exists in under conditions of 1 atm and 25 Degrees Celcius

27
Q

What are standard conditions for a compound if:

a) a gas
b) a pure substance
c) a substance present in solution

A

a) at a pressure of exactly 1 atm
b) in a condensed state (liquid/solid), at 1 atm
c) at a concentration of exactly 1M

28
Q

For an exothermic reaction:

a) how does enthalpy of reactants and products compare
b) what is the value
c) what occurs with heat energy
d) draw a graph labelling ΔH and Ea

A

a) enthalpy of reactants > products
b) ΔH is negative
c) heat is given out/released to surroundings
d) see document

29
Q

Give 3 examples of an exothermic reaction.

A
  1. burning
  2. respiration
  3. production of quicklime
30
Q

State what happens to the temperature of the system in exothermic reactions.

A

It will be observed to rise

31
Q

For an endothermic reaction:

a) how does enthalpy of reactants and products compare
b) what is the value
c) what occurs with heat energy
d) draw a graph labelling ΔH and Ea

A

a) enthalpy of products > reactants
b) ΔH is positive
c) heat is absorbed/taken in from surroundings
d) see document

32
Q

Give 3 examples of an endothermic reaction.

A
  1. photosynthesis
  2. thermal decomposition of limestone
  3. an electron being excited from n=1 to n=3
33
Q

What happens in terms of energy in bond-making and bond-breaking?

A 
bond-breaking = add energy
bond-making = release energy
34
Q

Define ‘entropy.’

A

measures the amount of disorder of a system

35
Q

Finish the following sentence:

The greater the amount of disorder, or randomness of the particles in a substance/mixture …

A

… the higher the entropy

36
Q

State the formula with units for change in entropy.

A 
ΔS = S(final) - S(initial)
S = measured in J/mol/K
37
Q

State the formula for total entropy change.

A

ΔTotal(total) = ΔS(system) - ΔS(surroundings)

38
Q

State the ‘second law of thermodynamics’ and explain why entropy will always decrease

A

entropy tends towards a maximum

39
Q

What do all spontaneously-occurring chemical and physical changes involve?

A

an overall increase in entropy

40
Q

Give 5 examples where entropy increases.

A
  1. solids melting
  2. liquids boiling
  3. number of molecules increases
  4. ionic solids dissolve
  5. temperature increases
41
Q

What is ‘Gibbs free energy?’

A

energy from a reaction free to do work

42
Q

State the formula for Gibbs free energy (including units).

A 
ΔG = ΔH - TΔS
ΔG in J/mol
ΔH in KJ/mol
T in Kelvin (K)
ΔS J/K/mol
43
Q

Interpret the spontaneity of each ΔG value:

a) ΔG < 0
b) ΔG > 0
c) ΔG = 0

A

a) reaction will be spontaneous
b) reaction needs energy input to happen
c) the system is in equilibrium

44
Q

What ΔG value corresponds to exergonic and endergonic reactions?

A 
exergonic = ΔG negative 
endergonic = ΔG positive
45
Q

Draw a table to show how ΔH, ΔS and -TΔS affect the value of ΔG.

A

see document

46
Q

State the formula for the temperature for a reaction to occur spontaneously and state units required.

A

T = ΔH/ΔS
T = Kelvin (K)
ΔH and ΔS = must be in same units
(here, ΔG is assumed to be zero)

47
Q

Define a ‘catabolic’ reaction.

A

high-energy, complex, compounds are broken down into simple ones

48
Q

Give 2 examples of catabolic reactions.

A
  1. Glucose + 6O^2 → 6CO^2 + 6H^2O

2. ATP + H^2O → ADP + Pi (phosphate transfer)

49
Q

Define an ‘anabolic’ reaction.

A

building-up of complex molcules from smaller ones

50
Q

Give 2 examples of an anabolic reaction.

A

takes place in steps coupled with ATP hydrolysis or another exergonic reaction

51
Q

Give 4 reactions that are coupled in metabolism.

A

glucose + fructose → sucrose
glucose + ATP → glucose-P + ADP
fructose + ATP → Fructose-P + ADP
glucose-P + fructose-P → sucrose + 2Pi

Chemistry (12 decks)

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