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A Level Chemistry > Module 3 (chapter 9 and 10) - enthalpy and catalysts > Flashcards

Module 3 (chapter 9 and 10) - enthalpy and catalysts Flashcards

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1
Q

enthalpy

A

(H) a measure of heat energy stored in a system

-a chemical system refers to the atoms, molecules or ions making up the chemicals

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2
Q

enthalpy changes

A

(delta H) the difference between the enthalpy of the reactants and the enthalpy of the products
-enthalpy of the products minus the enthalpy of reactants

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3
Q

heat

A

the process whereby thermal energy (J) is transferred from a hotter object to a cooler one

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4
Q

temperature

A

the direction of energy transferred is determined by the temperature of the objects (K or degrees)

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5
Q

system and surrounding

A
  • system is the chemicals (reactants and products)
  • surroundings are the apparatus and everything that isn’t in the chemical system
  • the universe is everything it includes both the system and the surroundings
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6
Q

exothermic reactions

A
  • where the enthalpy of the products is smaller than the enthalpy of the reactants due to heat loss to the surroundings
  • reactants higher than products
  • the energy lost by the system is balanced by the same energy gain from the surroundings
  • temperature goes up
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7
Q

endothermic reactions

A

where the enthalpy of the products is greater than the enthalpy of the reactants de to heat being taken in from the surroundings

  • temperature goes down
  • any energy gain by the chemical system is balanced by the same energy loss by the surroundings
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8
Q

activation energy

A

the minimum amount of energy required for a reaction to take place by breaking chemical bonds in the reactants
-they determine how reactions happen, if there was no activation energy, every reaction would happen

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9
Q

average bond enthalpy

A

the average enthalpy change that takes place when breaking by homolytic fission 1 mol of a given type of bond in the molecules of a gaseous species

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10
Q

why is the true value for the energy change slightly different from there value calculate using bond enthalpies

A

we use average bond enthalpies and these vary in different environments

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11
Q

standard state

A

refers to the enthalpy change under standard conditions

-standard state refers to the physical state of a substance under standard conditions

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12
Q

standard conditions

A
  • standard pressure (100kPa, 1atm)
  • standard temperature (298K, 25 degrees)
  • standard concentration (1moldm-3)
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13
Q

standard enthalpy change of reaction

A

the enthalpy change that accompanies a reaction in the molar quantities shown in a chemical equation under standard contains with all reactants and products in their standard states
delta r H sigma

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14
Q

enthalpy change of formation

A

the enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions with all reactants and products in their standard states
-elements have an enthalpy change of formation of 0 as it is the formation of one mole of an element from its element (no change)
delta f H sigma

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15
Q

enthalpy change of combustion

A

the enthalpy change of combustion is the enthalpy change that takes place when one mole of a substance reacts completely with oxygen under standard conditions with all reactants and products in standard states.
delta c H sigma
-when a substance reacts completely with oxygen the products are the oxides of the elements in their substance

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16
Q

enthalpy change of neutralisation

A

the enthalpy change that accompanies the reaction of an acid by a base to form one mole of H20 (l) under standard conditions with all reactants and products in their standard states.
-involves the reaction of H+ with OH- to form one mole of H20 and so the value is the same for all neutralisation reactions

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17
Q

how do you calculate energy change

A

Q = m x c x delta T

  • (m = mass of the surroundings (g) and is used to identify the materials that are changing temperature)
  • (c = specific heat capacity changes for each substance but for water is 4.18 Jg-1K-1)
  • (delta T = change in temperature by thermometer (final - initial)
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18
Q

specific heat capacity

A

the energy required to raise the temperature of 1g of a substance by 1K

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19
Q

how accurate is the experimental enthalpy change value?

A
  • heat loss to the surroundings other than the water including the beaker but mainly the air surrounding and flame
  • incomplete reaction (e.g. combustion of methanol to form carbon monoxide rather than dioxide)
  • evaporation of methanol from the wick (burner must be weighed as soon as possible after extinguishing the flame, there’s usually a cover on the spirit burner to reduce this)
  • non standard conditions (the conditions for this experiment are unlikely to be identical the the standard conditions)
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20
Q

how to measure enthalpy using spirit burners?

A
  • measure out a set volume of water and pour into beaker, measure initial temp of this
  • add methanol to spirit burner and weigh it
  • light it, place the spirit burner under the water and stir with the thermometer
  • after three minutes extinguish flame and record temp of water
  • reweigh the spirit burner (assume the wick hasn’t been burnt)
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21
Q

what is a way of mitigating heat loss

A

use a plastic cup made from polystyrene, these are cheap, waterproof and light weight offering some insulation

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22
Q

what are the surroundings for enthalpy

A
  • the solution itself is the immediate surroundings
  • the chemical particles within the solution may react when they collide and any energy transfer is between the chemical particles and water molecules in the solution
  • a thermometer will record any temperature change
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23
Q

averaged bond enthalpy

A

the energy requited to break one mole of a specified type of bond in gaseous molecule

  • bond enthalpies are always endothermic and so have a positive energy value
  • this is because energy is always required to break bonds
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24
Q

limitations of average bond enthalpies

A
  • average bond enthalpy can very on the chemical environment of the bond.
  • you will therefore often be given an average bond enthalpy rather than the actual bond enthalpy of an individual bond. this may lead to slight variations
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25
Q

enthalpy of bond making and bond breaking

A

bond breaking is endothermic (positive)

bond making is exothermic (negative)

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26
Q

how do you calculate enthalpy changes from average bond enthalpies

A

sum of bond enthalpies in reactants - sum of bond enthalpies in products

27
Q

how to measure enthalpy changes

A
  • using a calorimeter
  • themometer
  • bon enthalpies
28
Q

Hess’s law

A 
if a reaction can take place by more than one route and the initial and final conditions are the same, the total enthalpy change is the same 
delta H (route A) = delta H (route B) - delta H (route C)
-when doing it if it stays as it is, things cancel (e.g. oxygen in excess you can ignore them)  
-allows the determination if enthalpy changes indirectly, coming from the idea of conservation of energy with there being two routes to covert reactants into products
29
Q

problems with hess’s law

A
  • high activation energy
  • a slow reaction rate
  • more than one reaction taking place
30
Q

indirect routes - enthalpy change of formation

A

delta reaction H = (sum of the enthalpy change of products) - (sum of enthalpy change of formation reactants)

31
Q

indirect routes - enthalpy change of combustion

A

delta reaction H = (sum of the enthalpy change of reactants) - (sum of the enthalpy change of products)

32
Q

rate of chemical reaction

A

how fast a reactant is being used up or how fast a product is being formed
-rate of reaction is the change in concentration of a reaction of product in a given time
(moldm-3s-1)

33
Q

how does rate of reaction change

A
  • fastest at the start of a reaction as each reactant is at its highest concentration
  • the rate of reaction slows down ad the reaction proceeds because the reactants are being used up and their concentration decrease
  • once one of the reactants has been used up, the concentrations stop changing and the rate of reaction is completely zero.
34
Q

why are some collisions effective whilst others are not?

A
  • the particles must collide with the correct orientation

- they must have sufficient energy to overcome the activation energy barrier of the reaction

35
Q

how does increasing the concentration effect the reaction?

A

increasing the concentration increases the number of particles in the same volume

  • the particles are closer together and therefore collide more frequently
  • in a given period of time there will therefore be more effective collisions and an increased rate of reaction
36
Q

how doe increasing the pressure of a gas affect the rate of reaction?

A
  • when a gas is compressed into a smaller volume, the concentration of the gas molecules increases as the same number of gas molecules are in a smaller volume
  • the gas molecules are closer together and collide more frequently leading to more effective collisions in the same time.
37
Q

how do you monitor the progress of a reaction?

A
  • monitoring the removal (decrease in concentration) of a reactant
  • following the formation (increase in concentration) of a product
  • this will depend on the properties and physical states of the reactants and products in the reaction
38
Q

how do you measure reactions that produce gases?

A
  • monitoring the volume of gas produced a regular time intervals using gas collection
  • monitoring the loss of mass of reactants using a balance
  • both are proportional to the change in concentration of reactants and products and so can give a measure of the rate of reaction
39
Q

how do you monitor reactions using a balance

A
  • the mass of the flash and contents is initially recorded and at regular time intervals
  • the reaction is complete when no more gas is produced so no more mass is then lost
40
Q

catalysts

A
  • substance that changes the rate of reaction without undergoing any permanent change itself
  • not used up during the reaction
  • may react with a reactant to form an intermediate or may provide a surface on which the reaction can take place
  • at the end the catalyst is regenerated
41
Q

how does a catalyst work

A

-increases the rate of reaction by providing an alternate reaction pathway of lower activation energy

42
Q

energy profile diagrams

A
  • endothermic = products higher than reactants

- exothermic = reactants higher than products

43
Q

homogeneous catalysts

A
  • has the same physical state as the reactants

- the catalysts reacts to form an intermediate which then breaks down to give the products and regenerates the catalyst

44
Q

examples of homogeneous catalysts

A
  • making ester with sulphuric acid as a catalyst
  • the reactants (ethanol and ethnic acid) and the catalyst are all liquids
  • ozone depletion with chlorine radicals as catalysts
  • the reactant (O3) and the catalyst (Cl.) are both gases
45
Q

heterogeneous catalysts

A
  • has a different physical state from the reactants
  • usually solids in contact with gaseous reactants or reactants in solution
  • reactant molecules are absorbed onto the surface of the catalyst where the reaction takes place. after reaction, the product molecules leave the surface of the catalyst by desorption
46
Q

examples of heterogenous catalysts

A
  • iron is a solid in the hater process (reactants are gases)
  • nickel is a solid in the hydrogenation of alkene (reactants are gases)
47
Q

catalysts -sustainability and economic importance

A
  • reduces the temperature needed for the process and the energy requirements
  • if a chemical process requires less energy, then less electricity of fossil fuel is used
  • making the product faster and using less energy can cut costs and increase profitability
  • the economic advantages outweigh any costs associated with developing a catalytic process
  • the modern focus on sustainability requires industry to operate processes with high atom economies and fewer pollutants
  • using less fossil fuels will cut carbon dioxide emissions, mitigating against global warming
48
Q

particles energy

A
  • some molecules more slowly with low energy, other faster with higher energy
  • the spread of molecules energy is known as the Bolzmann distribution
49
Q

features of a Boltzmann distribution

A
  • no molecules have zero energy (the curve starts at the origin)
  • the area under the curve is equal to the total number of molecules
  • there is no maximum energy for a molecule (the curve doesn’t meet the x axis at high energy, it would need to reach infinite energy to do s0)
  • y axis is number of molecules with a given energy
  • x axis is energy
50
Q

Boltzmann distribution and higher temperatures

A
  • more molecules have an energy greater than or equal to the activation energy
  • therefore a greater proportion of collisions will lead to a reaction, increasing the rate of reaction
  • collisions will also be more frequent as the molecules are moving faster, but the increased energy of the molecules is much more important than the increased frequency of collisions
  • the peak is lower and shifted to the right meaning a greater proportion of molecules can overcome the activation energy
51
Q

the Bolzmann distribution and catalysts

A

-activation energy shifts left meaning a greater proportion of molecules exceed the new lower activation energy

52
Q

reversible reactions

A
  • reactions that take place in both the forwards and reverse directions
  • e.g. haber process
53
Q

dynamic equilibrium

A
  • the rate of the forwards reaction is equal to the rate of the reverse reaction
  • the concentrations of the reactants and products do not change
  • must be a closed system, isolated from the surroundings meaning components aren’t effected from outside influences
54
Q

le Chatelier’s principal

A
  • the position of the equilibrium indicates the extent of the reaction
  • states that when a system in equilibrium is subjected to an external change the system readjusts itself to minimise the effect of the change
55
Q

how does concentration effect equilibrium?

A
  • if there are more products formed, the position of the equilibrium will shift right
  • if there are more reactants formed the position of the equilibrium will shift left
56
Q

investigating changes in equilibrium concentration

A

chromate ions are yellow and dichromate are orange

  • add yellow potassium chromate to a beaker
  • add dilute sulphuric acid dropwise until there is no fitter change (solution turns orange)
  • add aq sodium hydroxide until there is no further change in colour (solution goes back to yellow)
57
Q

how does the concentration investigation work?

A
  • when you add sulphuric acid you are increasing the concentration of H+ ions. This increases the rate of the forwards reaction and so causes the position of equilibrium to shift right to minimising the H+ concentration
  • the equation shifts right making more products turning solution orange
  • when sodium hydroxide added, OH- ions react with H+ ions decreasing concentration of H+ ions
  • this shift increases the concentration of the reactant that has been removed, H+
  • this decreases the rate the forwards reaction causing the equilibrium to shift left, turning it yellow
58
Q

effect of pressure on equilibrium

A
  • only if there are more gases molecules on one side of the equation than on the other
  • pressure is proportional to concentration
  • increasing pressure will shift the equilibrium to the side with the fewer molecules, reducing the pressure of the system
  • decreasing pressure shifts the position in the opposite direction
59
Q

effect of catalyst on equilibrium

A
  • does not change the posit of equilibrium
  • it speeds up both reactions equally
  • it will increase the rate at which an equilibrium is established
60
Q

the equilibrium law

A

-used to calculate the exact position of equilibrium
-equilibrium constant = Kc
Kc = [C]c [D]d / [A]a [B]b
(products/reactants)
-square brackets are short for concentration of
-abcd are balancing numbers in the overall equation
-[A] are the equilibrium concentrations of the reactants and products of this equilibrium

61
Q

what does Kc tell us?

A
  • a Kc value of 1 indicates a position of equilibrium that is halfway between reactants and products
  • a Kc value > 1 indicates a position of equilibrium that is towards the products
  • a Kc value < 1 indicates a position of equilibrium that is towards the reactants
  • larger Kc the further right the equilibrium
62
Q

effect of temperature of equilibrium

A
  • increasing temperature shifts the equilibrium positions in the endothermic direction
  • decreasing temperature shifts it in the exothermic direction
63
Q

how do you investigate the effects of temperature on equilibrium?

A
  • dissolve cobalt chloride in water in a boiling tube. add a small quantity of HCl and place in iced water.
  • solution is pink at this point
  • set up water bath and place the boiling tube into the boiling water and the solution turns blue
  • transfer boiling tube back into iced water and solution turns pink
64
Q

how does the investigation of temperature of equilibrium work?

A

cobalt chloride (blue) dissolves in water to form a pink solution and the forwards reaction

  • [Co(H20)6]2+ + 4Cl- –> [CoCl4]2- + 6H20
  • forwards reaction is endothermic so the position shifts to the right to take heat energy in and minimise the increase in temperature
  • this turns the solution blue
  • decreasing temperature shifts the position of equilibrium ledt to give out energy

A Level Chemistry (21 decks)

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