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Chemistry AS > AS Chemistry Term 3 > Flashcards

AS Chemistry Term 3 Flashcards

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1
Q

Define a reversible reaction

A

A reaction in which the products can react to re-form the original reactants.

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2
Q

Define an equilibrium reaction and state the characteristics of an equilibrium reaction

A

A reversible reaction in which the products react together to re-form reactants at the same time as the reactants are forming products.

Characteristics:

  1. It is dynamic: reactants and products are continuously reacting.
  2. At equilibrium the rate of forward reaction equals the rate of the backward reaction.
  3. The concentrations of reactants and products remain constant at equilibrium.
  4. Requires a closed system, a system from which none of the reactants or products can escape from the reaction mixture.
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3
Q

State Le Chatelier’s Principle

A

If one or more factors are changed in the equilibrium reaction, the position of equilibrium shifts in the direction that opposes the change.

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4
Q

Describe the effect of concentration on the position of equilibrium

A

If concentration of products is increased:
- position of equilibrium shifts towards reactants to decrease the concentration of the products in order to oppose the change.
- More reactants are formed.
If concentration of reactants is increased:
- position of equilibrium shifts towards products to decrease concentration of the reactants
- More products are formed.

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5
Q

Describe the effect of pressure on the position of equilibrium

A
  • Only affects reactions where gases are reactants or products

Increase in pressure:

  • Molecules are closer together
  • Position of equilibrium shifts in direction of fewer gas molecules as this is the direction that opposes the increase in pressure.

Decrease in pressure:

  • Molecules are further apart
  • Position of equilibrium shifts in direction of more gas molecules.
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6
Q

Describe the effect of temperature on the position of equilibrium

A

Increase in temperature:

  • Increases energy of surroundings
  • reaction will go in direction that opposes increase in energy, in which energy is absorbed.
  • If endothermic, will favour forward reaction.
  • If exothermic, will favour reverse reaction

Decrease in temperature:

  • Decreases energy of surroundings
  • To oppose change, energy is released
  • If endothermic, will favour reverse reaction
  • If exothermic, will favour forward reaction
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7
Q

Describe the effect of a catalyst on the position of equilibrium

A
  • Increases rate of forward and backward reaction equally.
  • This means it takes less time to reach the equilibrium.
  • Has no effect on the position of equilibrium
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8
Q

Describe the effect of changing conditions on the equilibrium constant

A

Concentration: Kc does not change if concentration is changed.

Pressure: Kc does not change if pressure is changed.

Temperature: Kc does change. Depends on the type of reaction, whether temperature is increased or decreased, and if more products or reactants are formed.

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9
Q

Describe and explain the conditions used in the Haber process

A
  • It is an exothermic reaction meaning decreasing the temperature would increase yield.
  • However a temperature of about 450°C is used as this increases the rate of reaction meaning more reactions can occur in the same time producing a higher overall yield.
  • A pressure of 350 atm is used as a high pressure favours the forward reaction.
  • An iron catalyst is used
  • Unreacted gases are passed over the catalyst bed again.
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10
Q

Describe and explain the conditions used in the Contact Process

A
  • Exothermic reaction so yield is decreased by increasing temperature.
  • Temperature of 450 °C is used to increase rate of reaction.
  • Vanadium Pentoxide catalyst is used
  • Pressure is just above atmospheric pressure, even though high pressure would increase yield, as Kp is very high meaning it is unnecessary and expensive.
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11
Q

Define acids, bases, conjugate acids and conjugate bases

A

Acids: A proton donor

Bases: A proton acceptor

Conjugate Acid: The acid which is formed from its base accepting a proton, in an equilibrium reaction.

Conjugate Base: The base which is formed from its acid donating a proton, in an equilibrium reaction.

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12
Q

Define strong and weak acids and bases

A
  • Acids that disassociate almost completely in solution are strong acids.
  • Acids that partially disassociate in solution are weak acids
  • Bases that disassociated almost completely in solution are strong bases
  • Bases that disassociate partially in solution are weak bases.
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13
Q

Define the term rate of reaction

A

The change in the amount of reactants or products over time.

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14
Q

Describe the effect of concentration on the rate of reaction

A

The more concentrated a solution, the greater the number of particles in unit volume. This increases the frequency of collisions between reacting particles, increasing the rate of reaction.

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15
Q

Define activation energy and discuss its relevance on the Boltzmann distribution

A

The Boltzmann distribution shows the distribution of energies at a given temperature. At a certain point on the energy axis is the activation energy which is the minimum energy required for colliding particles to react.

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16
Q

Describe the effect of temperature on the rate of reaction

A

Raising temperature increases the average kinetic energy of the particles. As the particles move more quickly, the frequency of collisions increase. The proportion of successful collisions also increase as more particles have the required activation energy. This is shown in the Boltzmann distribution by a flattening and shifting to the right of the curve.

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17
Q

Explain the effect of catalysts on the rate of reaction

A

Catalysts are substances that increase the rate of reaction but remain chemically unchanged by the end of the reaction. They provide an alternative reaction pathway which has a lower activation energy. This means a higher proportion of successful collisions occur, increasing the rate of reaction. This is seen in the Boltzmann distribution by the shifting of the activation energy line to the left, meaning a higher proportion of particles have the required activation energy.

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18
Q

Define homogenous and heterogenous catalysts

A

Homogenous: the catalyst and reactants are in the same phase

Heterogenous: Catalyst is in a different phase to the reactants

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19
Q

Describe an enzyme

A
  • biological catalyst
  • large protein molecules
  • are much more efficient than inorganic catalysts
  • are very specific, usually only catalyse one reaction
  • do not produce byproducts
  • work under mild conditions (temp: 35, pH:7, atmospheric pressure)
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20
Q

Explain the effect of surface area on the rate of reaction

A

Increasing SA increases number of contact points on the particles for the collision to occur meaning frequency of collisions increase and so the rate of reaction increases.

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21
Q

Define exothermic reactions and endothermic reactions, and give examples of both

A

Exothermic: A reaction in which energy is released to the surroundings, causing the temperature of the surroundings to increase.

e. g.
- combustion of fuels
- oxidation of carbohydrates in plants and animals i.e. respiration
- reaction of water with quicklime

Endothermic: A reaction in which energy is absorbed from surroundings, causing the temperature of the surroundings to decrease.

e. g.
- thermal decomposition reactions
- photosynthesis

22
Q

Define enthalpy change of reaction

A

The energy exchange between a chemical reaction and its surroundings at constant pressure. Enthalpy itself is the energy present in the reactants but cannot be measured. Enthalpy change is measured by exchange of heat energy.

23
Q

Define standard conditions

A

The conditions which are used when comparing enthalphy changes, in order to ensure a fair comparison.
Conditions:
- approximately atmospheric pressure (10^5 Pa)
- Temperature: 298K
- Substance in its normal physical state (gas, liquid, solid) at 298K

24
Q

Define the standard enthalpy change of reaction

A

The enthalpy change when the amount of reactants react to give products under standard conditions. Can be exothermic or endothermic.

25
Q

Define the standard enthalpy change of combustion

A

The enthalpy change when one mole of a substance is burnt in excess oxygen under standard conditions. It is always exothermic.

26
Q

Define the standard enthalpy change of neutralisation

A

The enthalpy change when one mole of water is formed by the reaction of an acid and an alkali under standard conditions.

27
Q

Define the standard enthalpy change of solution

A

The enthalpy change when one mole of solute is dissolved in a solvent to form an infinitely dilute solution (no enthalpy change occurs when further solute is added) under standard conditions.

28
Q

Define the standard enthalpy change of atomisation

A

The enthalpy change when one mole of gaseous atoms is formed from its elements under standard conditions

29
Q

Define the standard enthalpy change of formation

A

The enthalpy change when one mole of a compound is formed from its elements under standard conditions.

30
Q

Define the standard enthalpy change of hydration

A

The enthalpy change when one mole of a hydrated salt is formed from one mole of the anhydrous salt under standard conditions.

31
Q

Explain the relevance of bond breaking and forming in enthalpy changes

A
  • Breaking bonds requires energy as forces between atoms have to be overcome.
  • Energy is released when new bonds are formed.
  • This means if energy required to break bonds is greater than energy released by forming bonds, it is an endothermic reaction and vice versa.
32
Q

Define bond energy and explain why an average bond energy is used

A

Bond energy: The amount of energy required to break a specific covalent bond.

  • Average is used as a specific bond found in one molecule will not have the same energy as the same bond in another, different molecule.
33
Q

State Hess’s Law

A

The total enthalpy change in a chemical reaction is independent of the route by which the chemical reaction takes place as long as the initial and final conditions are the same.

34
Q

Describe the reactions of group 2 elements with oxygen

A

React with oxygen to form the oxides of the metal

e.g. 2Mg + O2 → 2MgO

35
Q

Describe the reactions of group 2 metals with water

A
  • They all react with water to form the metal hydroxide and hydrogen
    e.g. Ca + 2H2O → Ca(OH)2 + H2
  • However, hot magnesium reacts with steam forming magnesium oxide:
    Mg + H2O → MgO + H2
36
Q

Describe the reactions of group 2 metals with dilute acids

A

With HCl:
- All produce metal chloride and hydrogen gas

With H2SO4:

  • Beryllium and Magnesium form metal sulfate and hydrogen gas
  • Rest of metals form insoluble sulfate that covers metal and prevents further reaction. Means hydrogen gas cannot be produced. Calcium produces small amount but this eventually stops.

With HNO3:

  • With very dilute: Metal nitrate and hydrogen gas produced
  • With moderately dilute: Metal nitrate, nitric oxide and water are produced.
37
Q

Describe the reactions of oxides, hydroxides and carbonates of group 2 metals in water

A

Oxides: Produce hydroxide
e.g. CaO + H2O → Ca(OH)2

Hydroxides: Dissolves to form weakly alkaline solution
e.g. Ca(OH)2 → Ca2+ + 2OH-

Carbonates: Are all insoluble in water

38
Q

Describe the reactions of oxides, hydroxides and carbonates of group 2 metals in dilute acids

A

Oxides: Corresponding salt and water as metal oxides are basic so neutralisation occurs

Hydroxides: Strong base so neutralised by dilute hydrochloric acid

Carbonates: Forms corresponding salt, water and carbon dioxide

39
Q

Describe the thermal decomposition of group 2 nitrates and carbonates

A

Carbonates: Form metal oxide and carbon dioxide gas

Nitrates: Forms metal oxide, nitrogen dioxide - a brown, toxic gas - and oxygen.

  • The temperature required for thermal decomposition increases down group 2
40
Q

Describe the physical and chemical trends in group 2 metals as you go down the group

A

Metallic radius: Increases down the group as outer two electrons occupy a new principal quantum shell further from the nucleus.

Melting point: Generally decrease, however, magnesium is anomalous due to different types of packing of ions in metal structures.

Densities: General increase

Reactivity: Increases as ionisation energy decreases down the group meaning outer electrons are removed from outer shell more easily. This means they are oxidised easily and so are strong reducing agents.

41
Q

Describe the solubility of group 2 sulfates and hydroxides

A

Sulfates: Solubility decreases down the group

Hydroxides: Increases down the group, meaning metals down the group will have hydroxides that are more strongly alkaline.

42
Q

Describe the uses of calcium hydroxide and calcium carbonate

A

Calcium Hydroxide:
- Used by farmers to raise pH of acidic soil as it is basic and will neutralise acid.

Calcium Carbonate: (also known as limestone)

  • Used to make blocks for building
  • Used to make marble for expensive tiles
  • Used to make cement - limestone is roasted in lime kiln, limestone decomposes forming calcium oxide or lime. Lime is roasted with clay in lime kiln to form cement. Cement is mixed with sand and rock to make concrete.
43
Q

Explain why nitrogen is so unreactive and outline the reactions that it can undergo naturally

A
  • Due to its electron configuration, nitrogen atoms need to gain three electrons to achieve noble gas configuration meaning a triple bond forms.
  • This bond is extremely strong and will only react in extreme conditions:
  • During thunderstorms, lightning provides activation energy for NO to form from nitrogen and oxygen.
  • NO is oxidised from O2 in air giving NO2
  • Nitric acid forms when NO2 reacts with water in the rain and this dissolves into the soil and is absorbed by plants for growth.
44
Q

Describe why ammonia is basic and use this to show how ammonium is formed

A
  • Ammonia as a lone pair on the nitrogen atom that can be donated to a H+ ion forming a dative bond.
  • NH3 + H+ → NH4+
  • As ammonia accepts the H+ ion it is the base in this reaction.
45
Q

Describe how ammonia can be derived from its salts

A
  • If an ammonium salt is heated with a base, the ammonium ion produces ammonia gas.
    e. g. 2NH4Cl + Ca(OH)2 → CaCl2 + 2H2O + 2NH3
46
Q

Describe the uses of ammonia and ammonia compounds

A
  • Ammonia is made from Haber process and can be injected straight into soil as a fertiliser
  • Ammonium salts are usually made from ammonia by reacting ammonia with acids
  • Ammonia is also used to make nitric acid which is used in making one of the major ammonium salts, ammonium nitrate.
  • Also used to make explosives
  • Nitric acid used in detergents, paints and dyes.
47
Q

Describe the effect nitrate fertilisers have on the environment

A

Process called eutrophication:

  • Nitrates leach into groundwater by rain which move to lakes and rivers.
  • Nitrates promote the growth of water plants
  • When these plants and algae die, bacteria feed on them to decompose the dead material
  • Bacteria multiply rapidly as a result of food amount and use up dissolved oxygen in water
  • Aquatic life die as a result.
48
Q

Describe the effect of nitrogen oxides in the atmosphere

A
  • They cause acid rain and photochemical smog
  • Act as a catalyst in the oxidation of SO2 which reacts with rainwater forming sulfuric acid:
    1. SO2 + NO2 → SO3 + NO
    2. NO + 0.5O2 → NO2
49
Q

Explain how nitrogen oxides are made non-naturally and describe how this can be restricted

A
  • Air is compressed in engine cylinders and ignited by a spark forming nitrogen oxides which are released in exhaust fumes.
  • Catalytic converters are used to reduce pollutants: Reaction on surface of hot catalyst allows non-harmful gases to be produced
    2CO + 2NO → 2CO2 + N2
50
Q

Describe how sulfur dioxides are formed

A
  • From burning fossil fuels
  • They have sulfur compounds as impurities that are oxidised when fossil fuels are burnt.
  • Also released from volcanic activity
51
Q

Describe how acid rain is formed from sulfur dioxide

A
  • Sulfur dioxide which is oxidised by nitrogen dioxide to form sulfur trioxide reacts with water
    SO3 + H2O → H2SO4
52
Q

Describe the environmental effects of acid rain

A
  • Leaches nutrients in soil preventing healthy growth of plants.
  • Decrease pH of bodies of water killing many animals, affecting the entire food chain
  • Calcium carbonate and steel are corroded by acid rain

Chemistry AS (3 decks)

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