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Chemistry OCR A > Module 3: Periodic Table and Energy > Flashcards

Module 3: Periodic Table and Energy Flashcards

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1
Q

Define first ionisation energy.

A

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

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2
Q

Why is the second ionisation energy of an element greater than the first?

A

Once an electron has been removed, there is a proton : electron ratio, increasing nuclear /electrostatic force of attraction, decreasing atomic radius, and this means more energy is required to remove electrons.

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3
Q

What are the four factors you should always mention when answering a question about ionisation energy?

A
  1. Number of shells/shielding (these are separate points)
  2. Nuclear charge
  3. Nuclear attraction
  4. Atomic radius
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4
Q

Define periodicity.

A

Repeating trend in physical or chemical properties across periods in the periodic table.

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5
Q

Ionisation energy generally increases across period 2.

Why do drops occur?

A
  1. There is a fall between Be and B. This is because Be has an electron configuration of 1s² 2s² while B has an electron configuration of 1s² 2s² 2p¹. B has the first electron to be removed in the P sub-shell, which is further away and requires less energy to remove.
  2. There is a fall between N and O. This is because the first electron to be removed in N is in an orbital by itself, while the first electron to be removed in O is in an orbital with 2 electrons, which repel each other, which means less energy is required to remove it.
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6
Q

Why is graphite able to conduct electricity despite being a giant covalent structure?

A

Only 3 of its 4 valence electrons are in a covalent bond, the remaining electrons become part of a pool of delocalised electrons.

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7
Q

What are the three types of reactions that group 2 metals partake in?

A
  1. Metal + Oxygen → Metal oxide
  2. Metal + Water → Metal hydroxide + hydrogen
  3. Metal + Acid → Salt + hydrogen
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8
Q

What are the uses of group 2 metals?

A
  1. Neutralise acidic soils
  2. Treat acid indigestion by neutralising hydrochloric acid in the stomach
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9
Q

What two observations are made when magnesium reacts with steam?

A
  • Bright white light
  • White solid is formed
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10
Q

How does the reactivity of group 2 elements change down the group?

A

It increases, this is because the energy required to lose electrons decreases as increase in shielding outweighs increase in nuclear charge.

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11
Q

How does the solubility of group 2 sulfates in water change down the group?

A

Decreases.

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12
Q

What is the role of water in reactions with group 2 metals?

A

Oxidising agent.

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13
Q

How can you test to distinguish between group 2 metals?

A

Add NaOH because as you go down the group, the solubility of metal hydroxides increases (higher up metals form precipitates).

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14
Q

How does reactivity of group 7/17 halogens change as you go down the group?

A

It decreases, this is because the energy required to attract electrons increases as the increase in shielding outweighs the increase in nuclear charge.

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15
Q

How can you test to compare the reactivity of halogens?

A

A solution of one halogen is added to another aqueous solution another halide. If the halogen added is more reactive than the halide, a colour change will take place.

It’s difficult to distinguish between bromine and iodine so adding cyclohexane and shaking it will turn them orange and purple respectively.

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16
Q

What are the two disproportionation reactions that chlorine undergoes?

A
  1. Cl₂ + H₂O → HClO + HCl
    Chloric acid
  2. Cl₂ + 2NaOH → NaClO + NaCl + H₂O
    Sodium chlorate
    Must be cold and dilute
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17
Q

Outline the uses of chlorine and its risks.

A

Chlorine can disinfect drinking water, HClO can be used as a disinfectant and NaClO is a component of household bleach.

The risks of chlorine use include:
- Chlorinated hydrocarbons being carcinogenic
- Chlorine is a toxic gas that is a respiratory irritant in small concentrations and fatal in large concentrations

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18
Q

Describe the carbonate test.

A
  • Add HNO₃ to solution to be tested
  • If effervescence is produced, bubble it through limewater
  • If it goes cloudy, a CO₂ is present, which means that a carbonate was present in the original solution
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19
Q

Describe the sulfate test.

A
  • Add BaCl₂ to solution to be tested
  • BaSO₄ is insoluble in water, a white precipitate will form is a sulfate ion was present in the original solution
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20
Q

Describe the test for halides.

A
  • Add AgNO₃ to solution to be tested
  • If chlorine is present a white precipitate will form
  • If bromine is present a cream precipitate will form
  • If iodine is present a yellow precipitate will form
    If the colours are difficult to distinguish, add aqueous ammonia to test solubility, solubility decreases down the group
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21
Q

Why are fluorides not tested for in the halide test?

A

AgFl is water soluble and won’t form a precipitate.

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22
Q

Outline the test for ammonium ions.

A
  • Add NaOH to the solution to be tested and heat
  • If ammonium ions are present, ammonia gas will form, but there is no effervescence because ammonia is water soluble
  • Heat for a second time and hold damp litmus paper over solution
  • If it turns blue, ammonia is present
23
Q

What is enthalpy?

A

Energy stored in bonds.

24
Q

How is enthalpy change calculated?

A

ΔH = H(reactants) - H(products)

25
Q

What are the standard conditions?

A
  • Pressure: 100 kPa
  • Temperature: 298 K
  • Concentration: 1 moldm⁻³
  • State: Physical state of a substance under standard conditions
26
Q

Define standard enthalpy change of a reaction?

A

The enthalpy change that accompanies a reaction in the given molar quantities shown, under standard conditions with all reactants and products in their standard states.

27
Q

Define standard enthalpy change of formation.

A

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

28
Q

Define standard enthalpy change of combustion.

A

The enthalpy change that takes place when mole of a compound reacts completely with oxygen under standard conditions with all reactants and products in their standard states.

29
Q

Define standard enthalpy change of neutralisation.

A

The enthalpy change that accompanies the reaction between an acid and a base to form one mole of H₂O under standard conditions with all reactants and products in their standard states.

30
Q

What two equations are used to calculate standard enthalpy change?

A
  1. q = mcΔT
  2. ΔH = q/moles
31
Q

How and why would the value of enthalpy change of combustion obtained experimentally be different from the data book value?

A

How: it would be less exothermic (release less energy)
Why: Heat loss to surroundings other than water being heated, incomplete combustion of fuel, evaporation of fuel from wick, experiment taking place in non-standard conditions.

32
Q

Define average bond enthalpy.

A

Energy required to break one mole of a specific type of bond in a gaseous molecule. Bond enthalpies are always endothermic.

33
Q

Is bond breaking endothermic or exothermic?

A

Endothermic

34
Q

Is bond making endothermic or exothermic?

A

Exothermic

35
Q

State Hess’s 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.

36
Q

How do you include enthalpy changes of vaporisation etc when calculating bond enthalpies?

A

C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(l) → 3CO₂(g) + 4H₂O(g)

The first part of the reaction is all the reactants and products in their standard states.

The second part is the enthalpy of of vaporisation.

37
Q

What are the conditions required for dynamic equilibrium?

A
  • The forward and reverse reactions occur simultaneously at the same rate
  • Concentrations of products and reactant don’t change.
38
Q

State 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.

39
Q

What is the general expression for Kc?

A

[products]/[reactants]

40
Q

What is heterogenous equilibrium and how will it affect a Kc expression?

A

An equilibrium which contains species that aren’t all in the same state, and as a result, only aqueous and gaseous species in the reaction are included in the Kc expression and solid and liquid species are omitted.

41
Q

What is homogenous equilibrium and how will it affect a Kc expression?

A

An equilibrium which contains species that are all in the same state, and as a result, all species in the reaction are included in the Kc expression.

42
Q

What is the mole fraction of a gas?

A

Proportion of moles of a particular gas out of all the moles of all the gases in the gas mixture.

43
Q

Define partial pressure.

A

The contribution that a gas makes to the total pressure of a gas mixture.

44
Q

What is the general expression for Kp?

A

(products)/(reactants)
only includes gases

45
Q

What is the only change that can result in a change in the value of K and why?

A

Change in temperature. This because if the change in temperature favours the forward reaction, yield of products will increase, and since K = products/reactants, K must change as well. The opposite is also true

However, changes in concentration or pressure means a reaction mixture is no longer in equilibrium, and K has changed. As a result the system readjusts to restore the original value of K

46
Q

Describe and explain the Boltzmann distribution graph.

A
  • Energy on the x-axis and number of particles on the y-axis
  • Starts from origin because no particle has zero energy
  • Peak is the most probable energy of the particles
  • Never touches the x-axis after origin
  • Area under the graph is the number of particles
47
Q

How does increasing temperature affect the Boltzmann distribution?

A
  • Most probable energy shifts to the right
  • More particles have sufficient energy
  • Area under the graph is the same as number of particles hasn’t changed
48
Q

How does using a catalyst affect the Boltzmann distribution?

A

Shape doesn’t change but activation energy shifts to the left so more particles have sufficient energy.

49
Q

How does a catalyst work?

A

It increases the rate of reaction by lowering the activation energy without being used up. It does this by providing an alternative pathway for the reaction to take place.

50
Q

What is a homogenous catalyst and how does it work?

A

A catalyst that is the same physical state as the reactants.

The catalyst reacts to form an intermediate, and at the end of the reaction it is reformed.

e.g. H₂SO₄ in the esterification of carboxylic acids and chloride radicals in ozone depletion.

51
Q

What is a heterogenous catalyst and how does it work?

A

A catalyst with a different physical state to the reactants (usually a solid).

The catalyst provides a surface for the reaction to take place. The reactants stick to the surface of the catalyst (adsorption) and once the product forms, they leave the surface of the catalyst (desorption).

e.g Iron in the Haber process and nickel in the hydrogenation of alkenes.

52
Q

Explain the economic effects of catalysts.

A
  • Less energy is required, which saves money.
  • Solid catalysts are made from precious metals, which are expensive to buy and difficult to clean.
53
Q

Explain the environmental effects of catalysts.

A
  • Because less energy is needed for reactions, less fuels are burnt, reducing emissions.
  • The solid catalysts are usually made from precious metals, and mining them may cause environmental damage.
  • It’s difficult to dispose of precious metals once they can no longer be used.

Chemistry OCR A (5 decks)

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