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Chemistry > Module 3: Periodic Table and Energy > Flashcards

Module 3: Periodic Table and Energy Flashcards

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1
Q

How was the periodic table first arranged?

A

Mendeleev arranged elements in order of atomic mass, as well as lining up groups of elements with similar properties, leaving gaps if they didn’t fit.
This let him accurately predict the properties of undiscovered elements

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2
Q

How is the periodic table arranged now?

A
  • Elements arranged in increasing atomic number from left to right
  • Elements arranged in groups with the same number of outer-shell electrons
  • Elements arranged in periods - the number of the period gives the number of the highest energy electron shell
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3
Q

What is periodicity?

A

The repeating trends in properties of the elements across each period
- Electron configuration
- Ionisation energy
- Structure
- Melting/boiling points

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4
Q

Describe the trends in electron configuration across the periodic table

A
  • Each period starts with an electron in the new highest energy shell, e.g period 2 has the 2s and 2p shells, period 3 has 3s and 3p
  • Periodic table can be divided into blocks depending on elements’ highest energy sub-shell.
    Groups 1+2 are in the s-block
    Transition metals are the d-block
    Groups 3-8 are the p-block
    The unstable elements are the f-block
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5
Q

What is first ionisation energy?
Give an equation for it
How does this relate to successive ionisation energies?

A

The energy required to remove one electron from each atom in one mole of gaseous atoms of an element forming one mole of gaseous 1+ ions.

Na(g) –> Na+(g) + e-

Definition is the same for each ionisation energy, but turning (n)+ ions into (n+1)+ ions

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6
Q

What electron is lost in the first ionisation energy? Why?

A

The first electron lost is the highest energy level electron because it has the least attraction to the nucleus so is easiest to lose

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7
Q

Explain the factors affecting ionisation energy

A
  • Atomic radius - greater distance between nucleus + electrons = less attraction = lower ionisation energy
  • Nuclear charge -more protons in nucleus = greater attraction = higher ionisation energy
  • Electron shielding - electrons are negatively-charged so inner-shell electrons repel outer-shell electrons, reducing attraction + ionisation energy
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8
Q

How many ionisation energies does an element have?
What happens to the value of successive ionisation energies and why?

A

There are the same amount of ionisation energies as there are electrons

Successive ionisation energies increase because when an electron is lost, the remaining ones get pulled closer to the nucleus = greater attraction so it takes more energy to remove one

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9
Q

How do successive ionisation energies provide evidence for electron configuration?

A
  • Successive ionisation energies increase steadily until there is a large jump
  • The large increase indicates that the next electron is being removed from the next shell closer to the nucleus (smaller distance = greater attraction = increase in I.E)
  • From this you can determine how many electrons are in each shell and so the element’s identity
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10
Q

Describe the trends in first ionisation energies across the periodic table

A
  • Decreases down each group
  • General increase across a period
  • Decreases between groups 2 and 3
  • Decreases between groups 5 and 6
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11
Q

Explain why first ionisation energy:
- Decreases down a group
- Generally increases across a period

A
  • Atomic radius increases, more inner shells so shielding increases, outweighs increasing nuclear charge so nuclear attraction decreases = less energy to remove an electron
  • Nuclear charge increases, similar shielding, atomic radius decreases = nuclear attraction increases = more energy to remove an electron
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12
Q

Explain why first ionisation energy:
- Decreases from group 2-3
- Decreases from group 5-6

A
  • 2-3: in group 2, the outermost electron is in the s sub-shell, but in group 3 it is in the p sub-shell. P sub-shell is higher energy than s so it is easier to remove an electron from it
  • 5-6: in group 5, the p sub-shell (its outermost one) has all orbitals singularly filled, but in group 6, one orbital has paired electrons which repel each other, so it takes less energy to remove one
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13
Q

Describe metallic bonding

A

The electrostatic attraction between cations and delocalised electrons

  • Metal atoms lose their outer shell electrons to a shared pool of delocalised electrons
  • The cations are fixed in position by their electrostatic attraction to the electrons in a giant metallic lattice structure
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14
Q

Explain the properties of metals in reference to structure and bonding

A
  • Conduct electricity - delocalised electrons are mobile so can carry charges through the metallic structure
  • High melting/boiling points (usually) - takes a lot of energy to overcome the strong electrostatic attraction between the cations and electrons
  • Insoluble - interactions between polar solvents and the charges in a metallic lattice lead to a reaction not dissolving
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15
Q

What elements form giant covalent structures?
Describe the structures

A

The non-metals boron, carbon and silicon form giant covalent lattices

Carbon (as diamond) and silicon form 4 covalent bonds, making a tetrahedral structure with 109.5 degree bond angles

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16
Q

Explain the properties of giant covalent structures (excluding graphene and graphite)

A
  • High melting/boiling points - lots of energy needed to break the strong covalent bonds
  • Insoluble - covalent bonds are too strong to be broken by interactions with solvents
  • Non-conductors - there are no mobile charge carriers
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17
Q

Describe the structure of graphene

A
  • A single layer of graphite - carbon atoms in a planar hexagonal layer with 120 degree bond angles
  • Each carbon atom forms 3 covalent bonds, leaving 1 delocalised electron per atom free to carry a charge to conduct electricity
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18
Q

Describe the structure of graphite

A
  • Giant covalent structure
  • Parallel layers of hexagonally arranged carbon atoms, each layer is bonded by weak London forces
  • Each carbon atom forms 3 covalent bonds, leaving 1 delocalised electron per atom free to carry a charge to conduct electricity
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19
Q

Describe and explain the trends in melting points across the periodic table

A
  • Melting point increases from group 1-14 (4) across a period
  • Sharp decrease between groups 14 and 15 (4 and 5)
  • Melting points low from groups 15 - 18 (5-0)

Sharp decrease = a change from giant to simple molecular structures. Giant structures have strong forces to overcome, so have high melting points and vice versa

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20
Q

Give some general features of the group 2 elements (Alkaline Earth Metals)

A
  • Their metal hydroxides are alkaline
  • They are very reactive so don’t naturally occur in their elemental form
  • They are reducing agents (they are oxidised and donate their outer two electrons to another species)
  • Their most common type of reaction is redox reactions
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21
Q

Give the different types of redox reactions that group 2 elements undergo

A
  • Redox with oxygen:
    Group 2 metal(s) + oxygen(g) –> group 2 metal oxide(s)
  • Redox with water:
    Metal(s) + water(l) –> metal hydroxide(aq) + hydrogen(g)
  • Redox reaction with a dilute acid:
    Metal(s) + acid(aq) –> salt(aq) + hydrogen(g)
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22
Q

Explain the trend in reactivity as you go down group 2

A

Reactivity increases because ionisation energies decrease down the group due to nuclear attraction to the outer 2 electrons decreasing (increasing atomic radius + shielding).
The formation of metal 2+ ions requires the first and second ionisation energies so reactivity decreases

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23
Q

Describe group 2 hydroxides:
- How are they made?
- Give an example of how they are produced
- What happens when they are added to water?

A
  • Group 2 oxides react with water, releasing OH- ions and alkaline solutions of the metal hydroxide
  • CaO(s) + H2O(l) –> Ca2+(aq) + 2OH- (aq)
  • Group 2 hydroxides are only slightly soluble in water. When the solution is saturated, any further metal and hydroxide ions form a solid precipitate
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24
Q

Describe the trend of solubilities of the metal hydroxides as you go down group 2, and its implications

A
  • Solubility of the hydroxides in water increases down the group
  • Their solutions contain more OH- ions as you go down the group, so are more alkaline, so their pH increases
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25
Q

Describe an experiment used to show the trend of the solubilities of group 2 hydroxides

A
  • Add a spatula of each group 2 oxide to water in a test tube
  • Shake the mixture - there is insufficient water to dissolve all the metal hydroxide, so you will have a saturated hydroxide solution with a white precipitate at the bottom
  • Test the pH of each solution, the alkalinity will increase as you go down the group
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26
Q

Describe two uses of group 2 compounds

A
  • Agriculture - calcium hydroxide is added to fields as lime by farmers to increase the pH of acidic soils. Calcium hydroxide neutralises acid in the soil, forming water
  • Medicine - group 2 bases (magnesium hydroxide as milk of magnesia and calcium carbonate) are used as antacids for treating acid indigestion (stomach acid is HCl)
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27
Q

Give some equations that explain how milk of magnesia and calcium carbonate are used as antacids

A

Mg(OH)2 (s) + 2HCl(aq) –> MgCl2(aq) +2H2O(l)

CaCO3(s) + 2HCl(aq) –> CaCl2 (aq) + H2O(l) + CO2(g)

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28
Q

Give some general features of the group 17(7) elements - the halogens

A
  • The most reactive non-metal group so they don’t occur in their elemental form in nature, only as stable halide ions dissolved in sea water or combined with Na/K as salt deposits
  • Astatine is very rare and has never been seen because it is radioactive and decays rapidly
  • They are oxidising agents - they gain an electron lost from other species and are reduces
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29
Q

Explain the trend in boiling points as you go down group 7

A
  • Elements have more electrons as you go down group 7
  • So they have stronger London forces
  • So more energy is required to break their intermolecular forces
  • So boiling points increase
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30
Q

Describe the appearance and state of group 7 elements at RTP

A
  • Fluorine - pale yellow gas
  • Chlorine -pale green gas
  • Bromine - red-brown liquid
  • Iodine - shiny grey-black solid
  • Astatine - n/a - too unstable
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31
Q

What is a displacement reaction?

A

A more reactive halogen displaces a less reactive halogen from an aqueous solution of its salt

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32
Q

Describe an experiment to investigate displacement reactions

A
  • Add an organic non-polar solvent (e.g cyclohexane) to the halogen solutions (Cl, Br, I)
  • This is to make them easier to tell apart, as the non-polar halogens dissolve more readily in the non-polar solvent than in water, making more vibrant colours)
  • In cyclohexane, the top layer of Cl2 is pale green, Br2 is orange and I2 is violet
  • Add every combination of halogen ions (E.g NaCl, NaBr, NaI) to the halogen solutions
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33
Q

Describe the results of the experiment investigating displacement reactions

What does it show?

A
  • Br2 added to NaCl solution has no reaction
  • I2 added to NaCl has no reaction
  • Cl2 added to NaBr turns orange due to Br2 formation
  • I2 added to NaBr has no reaction
  • Cl2 added to NaI turns violet from I2 formation
  • Br2 added to NaI turns violet from I2 formation
  • Chlorine reacts with Br- and I-
  • Bromine reacts only with I-
  • Iodine doesn’t react at all
  • Shows that reactivity decreases down the group
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34
Q

Explain the reactivity trend as you go down group 17(7)

A

Reactivity decreases down the group because atomic radius and shielding increase, so there is less nuclear attraction to capture an electron from another species

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35
Q

What is a disproportionation reaction?

Give two examples

A

A redox reaction in which the same element is both oxidised and reduces

  • Chlorine with water
  • Chlorine with cold, dilute, aqueous NaOH
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36
Q

Describe the reaction of chlorine with cold, dilute, aqueous sodium hydroxide and its uses

A
  • Disproportionation reaction
  • Cl2(aq) +2NaOH(aq) –> NaClO(aq) + NaCl(aq) +H2O(l)
  • The reaction of chlorine with water is limited by its low solubility in water
  • If the water contains NaOH, much more dissolves + they react
  • The resulting solution has a large concentration of ClO- ions from the sodium chlorate (I) so it is used as a household bleach
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37
Q

Describe the risks and benefits of using chlorine in water purification

A

Risks:
- It is an extremely toxic gas - a respiratory irritant in small concentrations, fatal in large concentrations
- In drinking water it can react with organic hydrocarbons formed from decaying vegetation (e.g CH4), forming chlorinated hydrocarbons, which are a suspected carcinogen

Benefits:
- The risk of not adding chlorine (diseases like cholera) is greater than the risk from chlorinated hydrocarbons

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38
Q

What is qualitative analysis?

A

Relies on simple observations, not measurements
Can be carried out quickly on a test-tube scale

39
Q

Describe and explain the test for carbonate ions

A
  • Carbonates react with acids, forming CO2
  • Add nitric acid (not HCl or sulfuric acid as you’d be adding halide/sulfate ions)
  • If it fizzes, bubble the gas through limewater (saturated aqueous solution of calcium hydroxide)
  • If it turns cloudy , CO2 has reacted to form a white precipitate of CaCO3, so carbonate ions are present
40
Q

Describe and explain the test for sulfate ions

A
  • Barium sulfate is very insoluble
  • Aqueous barium ions are added to the solution (usually barium nitrate or barium chloride if you’re not going to test for halogens)
  • If sulfate ions are present, they react with the barium forming barium sulfate, which is a white precipitate
41
Q

Describe and explain the halide test (just precipitation)

A
  • Aqueous silver ions react with aqueous halide ions to form silver halide precipitates
  • Add aqueous silver nitrate (AgNO3) to the solution
    Silver chloride = white precipitate
    Silver bromide = cream precipitate
    Silver iodide = yellow precipitate
42
Q

Describe the way to improve the halide test
Why is it necessary?

A

The three silver precipitates have similar colours and are different to tell apart
- Add dilute ammonia, then concentrated ammonia to the solution(s) to test their solubilities
- Chlorine is soluble in dilute ammonia
- Bromine is soluble in only concentrated ammonia
- Iodine is insoluble in both

43
Q

Explain the correct order for the tests for anions

A

1) Carbonate test - only carbonates produce bubbles in acid, so there is no possibility of an incorrect conclusion
2) Sulfate test - barium carbonate also forms a white precipitate, so you must ensure that there are no carbonate ions first
3) Halide test - silver carbonate and sulfate are also insoluble and form precipitates, so you must rule out those possibilities first

44
Q

Describe how you would test for the identities of a mixture of ions

A

Test all on the same sample of solution
1) Carbonate test - add nitric acid (any other would add sulfate/chloride ions). If it bubbles, keep adding it until it stops to remove all carbonate ions
2) Sulfate test - add an excess of barium nitrate to the solution leftover. Any sulfate ions precipitate out as barium sulfate - filter the solution to remove it
3) Halide test - add silver nitrate to the solution leftover. Any precipitate must have halide ions as carbonate/sulfate ions have been removed. Add NH3 to confirm which halide you have

45
Q

Describe the test for cations (ammonium ions, NH4+)

A
  • When heated together, aqueous ammonium ions and aqueous hydroxide ions react to form ammonia gas, NH3
  • NH4+(aq) + OH-(aq) –> NH3(g) + H2O(l)
  • Add aqueous NaOH to the solution
  • Ammonia gas is produced but it dissolves in the water
  • Warm the mixture and the NH3 is released
  • You may smell the ammonia but if you test the gas with moist pH paper it turns blue (NH3 is alkaline)
46
Q

What is enthalpy (H)?

A

The measure of the heat energy in a chemical system. Sometimes thought of as the energy stored in the bonds. Enthalpy can’t be measured but change in enthalpy can

47
Q

What is a chemical system?

A

The atoms, molecules or ions making up the chemicals

48
Q

Give the equation for enthalpy change

A

Enthalpy change = (enthalpy of products) - (enthalpy of reactants)

49
Q

What is the conservation of energy? What does this mean for enthalpy changes?

A

Energy cannot be created or destroyed.
When a reaction involving an enthalpy change occurs, heat energy is transferred between the system and surroundings

50
Q

Describe an exothermic reaction.

Generally describe their enthalpy profile diagrams

A

The chemical system releases energy (causing a negative enthalpy change). The surroundings gain that energy (causing a temperature increase).

The reactants have more energy than the products, with a curve connecting the two

51
Q

Describe an endothermic reaction.

Generally describe their enthalpy profile diagrams

A

The chemical system takes in heat energy (causing a positive enthalpy change). The surroundings lose this energy (causing a temperature decrease).

The reactants have less energy than the products, with a curve connecting the two

52
Q

What is activation energy (Ea)?

What does this mean for reactions?

A

The minimum energy required for a reaction to occur.

In a reactions, bonds in reactants must be broken by an input energy so new bonds can form between the products. The energy input acts as an energy barrier (= Ea).

53
Q

How does the activation energy of a reaction affect rate?

A

Reactions with small activation energies occur rapidly as the energy needed to break bonds is readily available from the surroundings

Reactions with large activation energies occur very slowly or sometimes not at all, as the energy needed to break bonds is not present in the surroundings

54
Q

What is standard enthalpy change?

A

An enthalpy change under standard conditions:
- Standard pressure = 101kPa (1atm)
- Standard temp = 25 degrees C (298K)
- Standard concentration = 1mol/dm3
- Standard states of all the reactants at standard conditions

55
Q

What is the standard enthalpy change of reaction (ΔrH⦵)?

A

The enthalpy change of a reaction in the molar quantities shown in a chemical equation under standard conditions in standard states

56
Q

What is the standard enthalpy change of formation (ΔfH⦵)?

A

The enthalpy change when one mole of a compound is formed from its elements under standard conditions with all reactants and products in their standard states.

The enthalpy change of formation of an element is 0, as there is no change

57
Q

What is the standard enthalpy change of combustion (ΔcH⦵)?

A

The enthalpy change when one mole of a substance reacts completely with oxygen under standard conditions in their standard states

58
Q

What is the standard enthalpy change of neutralisation (ΔneutH⦵)?

A

The enthalpy change of a reaction of an acid and a base to form one mole of water under standard conditions in standard states.

It is the same for all neutralisation reactions (-57kJ/mol) as all have the ionic equation:
H+ + OH- –> H2O

59
Q

What is the equation for enthalpy change of reaction using bond enthalpy values?

A

ΔrH⦵ = (sum of energy of bonds broken) - (sum of energy from bonds forming)

60
Q

What is the equation for enthalpy change of reaction using standard enthalpy of formation values?

A

ΔrH⦵ = (sum of ΔfH⦵ of products) - (sum of ΔfH⦵ of reactants)

61
Q

What is the equation for enthalpy change of reaction using standard enthalpy of combustion values?

A

ΔrH⦵ = (sum of ΔcH⦵ of reactants) - (sum of ΔcH⦵ of products)

62
Q

How would you calculate ΔrH⦵ using experimental values?

A
  • Use q = mcΔT
    (Heat energy(J) = mass of solution(g) x specific heat capacity (4.18 for water) x temp change (same in C and K))
  • Divide by 1000 as q is in J but ΔrH⦵ is in kJ
  • Use ΔrH⦵ = q/n, using the smallest number of moles for n
63
Q

How would you determine the enthalpy change of combustion in an experiment?

A
  • Put 150cm3 of water into a beaker above a spirit burner
  • Record the initial temp of the water to 0.5 degrees
  • Add methanol (or another fuel) to the spirit burner and weigh it
  • Light the burner and stir the water with the thermometer for 3 mins
  • Extinguish the flame, immediately measure the temperature and re-weigh the spirit burner
64
Q

Why will the value recorded of enthalpy change of combustion be lower from an experiment than the true value?

How could you improve this?

A
  • Heat is lost to the surroundings
  • Incomplete combustion of methanol (fuel)
  • Evaporation of methanol (fuel) from the wick
  • Non-standard conditions in the lab
  • Insulate beaker
  • Have an oxygen supply to the spirit burner
  • Control conditions in the lab
65
Q

How would you find the enthalpy change of reaction or neutralisation from an experiment?

A
  • Add the reactants to a polystyrene cup (as it offers some insulation)
  • Measure the volume (as you can work out mass using the density of water - 1g/cm3)
  • Measure the change in temperature
  • The specific heat capacity will be that of water, 4.18J/g/K
66
Q

What is mean bond enthalpy?

Why is an average used?

A

The mean energy required to break one mole of a specified covalent bond in the gaseous state. It is always positive because it takes energy to break bonds.

The actual bond enthalpy can vary depending on the chemical environment of the bond

67
Q

Describe the breaking and formation of bonds in terms of exo/endo thermic proceses

A
  • Breaking bonds is endothermic (positive enthalpy change)
  • Making bonds is exothermic (negative enthalpy change)
  • The difference between energy required for bond breaking and energy released by bond making determines whether the overall reaction is exothermic or endothermic
68
Q

What is Hess’ law?

A

If 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

69
Q

What does rate of reaction measure?

Give the equation

A

How fast a reactant is being used up / a product being formed, measured in mol/dm^3/s

Rate = change in concentration / time

70
Q

Explain how the rate of a reaction changes as it progresses

A
  • Rate is fastest at the start as each reactant is at its highest concentration
  • Rate slows down as reaction proceeds as reactants are used up and their concentrations decrease
  • Once one reactant is completely used up, concentrations stop changing and the rate of reaction is zero
71
Q

Give some factors that affect the rate of reaction

A
  • Concentration / pressure of gases
  • Surface area of solid reactants
  • Temperature
  • Catalysts
71
Q

What are the two types of catalysts?

A
  • Homogenous catalyst
  • Heterogenous catalyst
71
Q

What is collision theory?

A

Two reacting particles must collide effectively for a reaction to occur

Only a small proportion of collisions are effective because particles must:
- collide with the correct orientation
- have sufficient energy to overcome the activation energy barrier of reaction

71
Q

Describe how:
- concentration/pressure of gases
- surface area of solid reactants
- temperature

affect rate of reaction

A

All increase rate because:
- More particles in a given volume so effective collisions are more frequent

  • A greater area for other reactants to collide with so effective collisions are more frequent
  • More particles have the activation energy so more frequent effective collisions. Particles also have more kinetic energy so move faster, increasing frequency of all collisions
72
Q

What are catalysts?
Describe how they increase rate of reaction

A

They are substances that change the rate of reaction without undergoing any permanent change.

They provide an alternative reaction pathway with a lower activation energy.

Could react with a reactant to produce an intermediate (and is regenerated at the end) or provide a surface on which the reaction can occur.

73
Q

Describe homogenous catalysts, including example reactions

A

Have the same physical state as the reactants

They react with the reactants to form an intermediate, which breaks down, giving the products and regenerating the catalyst

e.g making esters with a sulfuric acid catalyst, ozone depletion with Cl* radical catalyst

74
Q

Describe heterogenous catalysts, including example reactions

A

Have a different physical state to the reactants

Reactant molecules are adsorbed (weakly bonded) to the surface, where the reaction occurs. Product molecules then leave the surface by desorption.

e.g Haber process with an iron catalyst, hydrogenation of alkenes with a nickel catalyst

75
Q

What are the benefits of using catalysts in industry?

A

Economic - lowers energy requirements (lower activation energy), cutting fuel costs and increasing profit

Environmental - lower fuel demands means less fossil fuels are burnt, reducing carbon dioxide emissions

76
Q

What is autocatalysis?

A

When a reaction product acts as a catalyst for that reaction

77
Q

What are two ways to follow reaction progress?

How do you adapt these for reactions involving gases?

A

By monitoring decrease in concentration of a reactant or increase in concentration of a product

You can measure the volume of gas produced / the mass lost, as these are proportional to the change in concentration of a reactant/product

78
Q

Give two ways to measure the rate of a reaction involving gases

A

1) Monitor gas produced:
- Put reactants in a conical flask and immediately replace the bung
- Measure the volume of gas in the gas syringe at regular intervals until no more gas is produced

2) Monitor mass lost:
- Add the reactants to an open conical flask on a mass balance
- Record initial mass and repeat measurements at regular time intervals until there is no change

At the end of both plot a graph of mass/volume against time and draw tangents to find rates

79
Q

What is the Boltzmann distribution?

Give some features of it

A

The spread of molecular energies in gases. Collisions of particles are elastic (no energy is lost as a result)

  • The curve starts at the origin as no molecules have zero energy
  • The area under the curve is the total number of molecules
  • Curve never reaches the x axis as there is no maximum energy for molecules
  • Only a small proportion of molecules have the activation energy
80
Q

How does the Boltzmann distribution change as:
- Temperature increases
- A catalyst is added

A
  • The peak is lower and shifted to the right, more molecules have the activation energy
  • The distribution of molecules is the same, there is a lower activation energy so more molecules have it
81
Q

What is a reversible reaction?

A

When products are being turned into reactants at the same time as reactants are being turned into products

82
Q

Give some characteristics of a reversible reaction at dynamic equilibrium

A
  • The rate of forward + reverse reactions are equal
  • Reactant + product concentrations are constant
  • Requires a closed system (isolated from its surroundings)
  • Dynamic as both reactions are still occurring even when the concentrations remain constant
83
Q

What is le Chatelier’s principle?

A

When a system in equilibrium is subjected to an external change, the system readjusts itself to minimise the effect of that change

84
Q

What does the position of equilibrium indicate?

A

The extent of the reaction

More products formed = position of equilibrium is more to the right
More reactants formed = position of equilibrium is more to the left

85
Q

What does increasing the concentration of a reactant in a reversible reaction at equilibrium do?

A

Position of equilibrium shifts to the right, producing more products (to decrease the concentration of the reactant), and vice versa in reverse conditions

86
Q

What does increasing temperature of a reversible reaction at equilibrium do?

A

If one reaction is exothermic, the reverse is equally endothermic.

This shifts the position of equilibrium to favour the endothermic reaction, and vice versa in reverse conditions

87
Q

What does increasing the pressure of a reversible reaction at equilibrium do?

A

Position of equilibrium shifts to favour the reaction producing fewer moles of gas, and vice versa in reverse conditions.

If there is no gas, or both sides produce equal moles of gas, pressure has no effect

88
Q

What effect do catalysts have on a reversible reaction at equilibrium?

A

Does not affect the position of equilibrium, as it speeds up the rate of both reactions equally

89
Q

What is the Haber process?

What kind of conditions are used for it and why?

A

The reversible reaction producing ammonia:
N2 + 3H2 <=> 2NH3, where the forward reaction is exothermic

Le Chatelier’s principle means a low temperature and high pressures maximises ammonia yield, but a low temperature makes the rate too slow and high pressures are too expensive and dangerous, so compromise conditions are used.

90
Q

For the general reversible reaction:
aA + bB <=> cC + dD,

Give the equation for the equilibrium constant

A

Kc = [C]^c[D]^d
[A]^a[B]^b

Where square brackets denote concentration

91
Q

What does the value of Kc indicate?

A

Position of equilibrium:

Kc = 1 means P.O.E is halfway between reactants + products
Kc > 1 means P.O.E is towards the products
Kc < 1 means P.O.E is towards the reactants

Chemistry (4 decks)

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