Module 3 Flashcards

1
Q

why do elements in the same group have similar chemical properties

A
  • same outer shell electron configuration
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2
Q

how are elements classified as s,p or d blocks

A

according to which orbitals the highest energy electrons are in

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3
Q

order of blocks on the periodic table

A

s,d,p,f

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4
Q

what does it mean if an element is in the s,d, p or f block

A

s block: only has s electrons in its outer shell
p block:at least one p-electron in the outer shell
d block: those with at least one d-electron and at least one s-electron but no f or p electrons in the outer shell (up to 5d)
f block: are all those with at least one f-electron and at least one s-electron but no d or p electrons in the outer shell

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5
Q

periodicity definition

A

repeating pattern across different periods

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6
Q

trend in atomic radius as you move across a period

A
  • atomic radius decreases
  • this is because the increased number of protons creates more positive charge attraction for electrons which are in the same shell
  • similar shielding
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7
Q

first ionisation energy definition

A

the energy needed to remove one mole of electrons from one mole of gaseous atoms

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8
Q

equation example for first ionisation energy

A

H(g) = H+(g) + e-

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9
Q

factors affecting ionisation energy

A
  1. attraction of the nucleus
    - more protons = greater attraction
  2. distance of the electrons from the nucleus
    - bigger atom means outer electrons are further from the nucleus, meaning weaker attraction to the nucleus
  3. shielding of the attraction of the nucleus
    - an electron in the outer shell is repelled by electrons in complete inner shells, weakening attraction of the nucleus
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10
Q

trend in first ionisation energy down a group

A
  • nuclear charge increases but ionisation energy decreases
    due to:
  • atomic radius increases
  • shielding increases
  • therefore, the attraction between the nucleus and outer electrons decreases
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11
Q

trend in first ionisation energy across a period

A
  • increases
    due to:
  • nuclear charge increases
  • distance between the nucleus and outer electron remains relatively constant
  • shielding remains the same
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12
Q

why are successive ionisation energies larger

A

the ion formed is smaller than the atom
proton to electron ratio in the 2+ ion is greater than in the 1+ ion
attraction between electron and nucleus is therefore stronger

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13
Q

what does it mean if there is a big jump between 2nd and 3rd ionisation energies?

A
  • the element must be in group 2
  • this is because the 3rd electron is removed from an electron shell closer to the nucleus with less shielding and so has a larger ionisation energy
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14
Q

why does helium have the largest first ionisation energy in group one

A
  • its first electron is in the first shell closest to the nucleus and has no shielding effects from inner shells
  • He has a bigger first ionisation energy than hydrogen as it has one more proton
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15
Q

why do 1st IE decrease down a group

A
  • outer electrons are found in shells further from nucleus
  • more shielded
  • attraction to nucleus decreases
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16
Q

why does 1st IE increase across a period

A
  • electrons added to the same shell which has the same distance from the nucleus and same shielding effect
  • number of protons increases making attraction greater
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17
Q

why is there a small drop in IE from group 2 to 3 eg from Mg to Al

A
  • group 2 has its outer electrons in the 3s sub shell, whereas group 3 is starting to fill a 3p sub shell
  • The electrons in the 3p subshell are slightly easier to remove because
    the 3p electrons are higher in energy and are also slightly shielded by the 3s electrons
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18
Q

why is there a small drop in IE from group 5 to 6 eg from P to S

A

in group 6 r there are 4 electrons in the 3p sub shell and the 4th is starting to doubly
fill the first 3p orbital.

When the second electron is added to a 3p orbital there is a slight repulsion between
the two negatively charged electrons which makes the second electron easier to
remove.

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19
Q

metallic bonding definition

A

the electrostatic force of attraction between
the positive metal ions and the delocalised electrons

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20
Q

what are the three main factors that affect the strength of metallic bonding

A
  1. Number of protons/ Strength of nuclear attraction.
    (The more protons the stronger the bond)
  2. Number of delocalised electrons per atom - the outer shell electrons are delocalised
    (The more delocalised electrons the stronger the bond)
  3. Size of ion.
    (The smaller the ion, the stronger the bond)
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21
Q

why has Mg got stronger metallic bonding than Na

A

Mg has stronger metallic bonding than Na and hence a higher melting point.

The Metallic bonding gets stronger because in Mg
there are more electrons in the outer shell that are released to the sea of electrons.

The Mg ion is also smaller and has one more proton. There is therefore a stronger electrostatic attraction
between the positive metal ions and the delocalised electrons and higher energy is needed to break bonds

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22
Q

explain the structure of diamond

A

MACROMOLECULAR
- Tetrahedral arrangement of
carbon atoms.
- 4 covalent bonds per atom
- high melting points because of strong covalent forces in the giant structure. It takes a lot of energy to break the many strong covalent bonds

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23
Q

explain the structure of graphite

A

MACROMOLECULAR
- Planar arrangement of carbon atoms in layers.
- 3 covalent bonds
per atom in each layer. 4th outer electron per atom is delocalised.
- Delocalised electrons between layers.
- high melting points because of strong covalent forces in the giant structure. It takes a lot of energy to break the many strong covalent bonds

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24
Q

what is graphene

A

a single layer of graphite

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25
Q

structure of silicon oxide

A
  • giant covalent lattice/ macromolecular
  • tetrahedral
  • each silicon is shared by four oxygens and each oxygen is shared by two silicons
  • empirical formula = SiO2
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26
Q

which structures have metallic bonding and give two examples

A

GIANT METALLIC LATTICES
- magnesium and sodium

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27
Q

which structures have covalent bonding and give three examples

A

MACROMOLECULAR
- diamond, graphite, silicon dioxide

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28
Q

properties of macromolecular substances

A

bp and mp: high due to many covalent bonds which take a lot of energy to break

solubility: insoluble

conductivity when solid: diamond and sand are poor because electrons cant move, whereas graphite is good as free electrons between layers

conductivity when molten: poor

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29
Q

properties of giant metallic substances

A

mp and bp points: high due to strong electrostatic forces between positive ions and sea of deloc electrons

solubility: insoluble

conductivity: good when solid and molten - deloc electrons can move through structure

general: malleable as the positive ions in the lattice are all identical, so the planes of ions can slide easily over one another.

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30
Q

explain the general trend of mp and bp across period three

A

or Na, Mg, Al- Metallic bonding : strong bonding – gets stronger the more electrons there are in the outer shell that
are released to the sea of electrons. A smaller positive
centre also makes the bonding stronger. High energy is needed to break bonds.

Si is Macromolecular: many strong covalent bonds
between atoms high energy needed to break covalent
bonds– very high mp +bp

Cl2 (g), S8 (s), P4 (S)- simple Molecular : weak London forces between molecules, so little energy is needed to break them – low mp+ bp

S8 has a higher mp than P4 because it has more electrons
(S8 =128)(P4=60) so has stronger London forces

Ar is monoatomic weak London forces between atoms

  • 2 is similar - boron is covalent
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31
Q

what happens to atomic radius down group 2

A
  • increases
  • atoms have more shells of electrons, making the atom bigger
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32
Q

melting point down group 2

A
  • decreases
  • metallic bonding weakens as the atomic size increases
  • distance between the + ions and deloc electrons increases
  • therefore, electrostatic
    attractive forces between the positive ions and the
    delocalized electrons weaken
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33
Q

what happens to first and second IE’s down group 2

A
  • decrease
  • The outermost electrons are held more weakly because they are successively further from the nucleus in additional
    shells
  • In addition, the outer shell electrons become more shielded from the attraction of the nucleus by the repulsive force of inner shell electron
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34
Q

second ionisation energy definition

A

the enthalpy change when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge

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35
Q

example equation for 2nd IE

A

Ti+ (g) = Ti2+(g) + e-

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36
Q

reactivity down group 2

A
  • increases
  • As the atomic radii increase there is more shielding.
  • The nuclear attraction decreases and it is easier to remove outer electrons.
  • Cations form more
    easily
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37
Q

group 2 metals reactions with oxygen
- give equation
- what colour flame?

A

The group 2 metals will burn in oxygen.
Mg burns with a bright white flame
2Mg + O2 → 2MgO

  • Mg will also react slowly with oxygen without a flame.
  • Mg ribbon will often have a thin layer of magnesium oxide on it formed by reaction with oxygen
  • this is cleaned off by emery paper before doing reactions with Mg ribbon as Mg and MgO react at different rates
  • if testing for reaction rates with Mg and acid, an un-cleaned Mg ribbon would give a false result as both the Mg and MgO would react but at different rates
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38
Q

Mg + steam
- produces?
- what colour flame?

A

produces magnesium oxide and hydrogen

the Mg would burn with a bright white flame

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39
Q

Mg + warm water
- produces?
- flame and speed of reaction?

A

produces magnesium hydroxide and hydrogen

Mg + 2H20 → Mg(OH)2 + H2
this is a slower reaction than with steam and produces no flame

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40
Q

group 2 metals and cold water

A
  • react with increasing vigour down the group to form hydroxides

eg : Ca + 2 H2O (l) → Ca(OH)2
(aq) + H2(g)
- hydroxides produced make the water alkaline

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41
Q

what are the observations for group 2 metals reacting with cold water

A
  • fizzing (more vigorous down the group)
  • the metal dissolving (faster down group)
  • the solution heating up (more down group)
  • with calcium, a white precipitate is formed (less precipitate forms down the group)
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42
Q

group 2 metals reacting with acid
give an example equation

A
  • react with acids with increasing vigour down the group to form a salt and hydrogen
    for example:
    Ca + 2HCl(aq) = CaCl2(aq) + H2 (g)
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43
Q

why does barium react slowly with H2SO4
give the equation

A

the insoluble barium sulfate produced will cover the surface of the metal and act as a barrier to further attack.

Ba + H2SO4 → BaSO4 + H2

  • the same effect will happen to a lesser extent with metals going up the group as solubility increases
  • does not happen with any other acids eg HCl
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44
Q

group 2 oxides + water
- forms?

A
  • forms hydroxides
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45
Q

why are group 2 oxides basic

A
  • the oxide ions accept H+ ions to become hydroxide ions
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46
Q

what is magnesium hydroxide used for
- soluble?

A

used in medicine (in suspension as milk of magnesia) to neutralise excess acid in the stomach so treats indigestion

Mg(OH)2 + 2HCl → MgCl2 + 2H2O

It is safe to use as it is weakly alkaline.

  • partially soluble in water
  • calcium carbonate can also be used to treat indigestion
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47
Q

what is calcium hydroxide used for
- is is soluble?

A
  • It is used in agriculture to neutralise acidic soils.
  • If too much calcium hydroxide is added to the soil, excess will result in soils becoming too alkaline to sustain crop growth
  • soluble
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48
Q

what is limewater?
what is it used for?
give equation

A
  1. aqueous solution of calcium hydroxide
  2. used as a test for co2
  3. Ca(OH)2 + CO2 → CaCO3 + H2O
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49
Q

appearance of the halogens

A

Fluorine (F2): very pale yellow gas. It is highly reactive

Chlorine : (Cl2) greenish, reactive gas, poisonous in high concentrations

Bromine (Br2) : red liquid, that gives off dense brown/orange poisonous fumes

Iodine (I2) : shiny grey solid sublimes to purple gas.

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50
Q

mp and bp down the halogens
- explain why

A

increases down the group

As the molecules become larger they have more electrons and so have larger induced dipole-dipole forces
(London forces) between the molecules. As the intermolecular forces get larger more energy has to be put into break these intermolecular forces. This increases the melting and boiling points

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51
Q

reactivity of the halogens going down the group

A
  • decreases
  • as the atoms get bigger with more shielding so they less easily attract and accept electrons. They therefore form -1 ions less easily down the group
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52
Q

displacement reactions of halogens

A

look at chemsheets

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53
Q

disproportionation reaction definition

A

name of a reaction where an element simultaneously oxidises and reduces

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54
Q

chlorine + water equation
how is it disproportionation?
what will happen is UI is added?

A

Cl2(g) + H2O(l) →HClO(aq) + HCl (aq)

Chlorine is both simultaneously reducing and oxidising. It changes from 0 in Cl2 to -1 in HCl and +1 in HClO

If some universal indicator is added to the solution it will first turn red due to the acidity of both reaction products. It will then turn
colourless as the HClO bleaches the colour.

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55
Q

what are the benefits and negatives of chlorine being used in water treatment?

A

kills bacteria =good
toxic and irritates the respiratory system= bad
forms chlorinated hydrocarbons = bad

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56
Q

chlorine + cold dilute NaOH
- reaction
- colour change
- what are the products used for

A

Cl2(aq) + 2NaOH(aq) → NaCl (aq) + NaClO (aq) + H2O(l)

  • disproportionation reaction
  • the colour of the halogen solution will fade to colourless
  • The mixture of NaCl and NaClO (sodium chlorate (I)) is used as Bleach and to disinfect/ kill bacteria
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57
Q

chlorine + hot NaOH

A

3Cl2 + 6NaOH → NaClO3 + 5NaCl + 3H2O

sodium chlorate (V) is formed
- disproportionation reaction

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58
Q

reaction of halide ions with silver nitrate
- used for?
- why is nitric acid added?
- colours?
- how can ammonia be used to help differentiate

A

used as a test to identify which halide ion is present. The test solution is made acidic with nitric acid, and then Silver nitrate solution is added drop wise

The role of nitric acid is to react with any carbonates
present to prevent formation of the precipitate
Ag2CO3 - This would mask the desired observations

Fluorides produce no precipitate
Chlorides produce a white precipitate
Ag+(aq) + Cl- (aq) →AgCl(s)
Bromides produce a cream precipitate
Ag+ (aq) + Br- (aq) → AgBr(s)
Iodides produce a pale yellow precipitate
Ag+ (aq) + I- (aq) →AgI(s)
The silver halide precipitates can be treated with ammonia
solution to help differentiate between them if the colours look similar:

Silver chloride dissolves in dilute ammonia to form a
complex ion
AgCl(s) + 2NH3(aq) → [Ag(NH3)2]+(aq) + Cl- (aq)

Silver bromide dissolves in concentrated ammonia to form a complex ion
AgBr(s) + 2NH3(aq) →[Ag(NH3)2]+(aq) + Br - (aq)

Silver iodide does not react with ammonia – it is too insoluble

59
Q

testing for the presence of a carbonate
- equation

A
  • add any dilute acid and observe effervescence
  • fizzing due to co2 would be observed if carbonate present
  • bubble gas through limewater to test for co2- will turn limewater cloudy
    2HCl + Na2CO3 → 2NaCl + H20 + CO2
60
Q

testing for presence of a sulfate
- why can sulfuric acid not be used
- equation

A
  • An acidified BaCl2 solution is used
  • white precipitate forms
  • acid is used as needed to react with carbonate impurities that are often found in salts which would form a white barium carbonate precipitate and give a false result
  • sulfuric acid cannot be used as it contains sulfate ions which would form a precipitate
  • Ba2+(aq) + SO42- (aq) = BaSO4 (s)
61
Q

how to test for an ammonium ion (positive ions/cations)
- how can ammonia gas be identified

A

react with NaOH(Aq), forming NH3 gas
- ammonia gas can be identified by its pungent smell or by turning red litmus paper blue

62
Q

in what order should you test for ions

A
  1. carbonate
  2. sulfate
  3. halide
    - this prevents false results
63
Q

is △H positive or negative in exothermic and endothermic reactions?

A

In an exothermic reaction the
∆H is negative

In an endothermic reaction
the ∆H is positive

64
Q

activation energy definition

A

the minimum energy which particles need to collide to start a reaction

65
Q

standard conditions

A
  • 100 kPa pressure
  • 298 K (room temperature or 25oC)
  • Solutions at 1mol dm-3
  • all substances should have their normal state at 298K
66
Q

standard enthalpy change of formation definition

A

the enthalpy change when 1 mole of the compound is formed from its elements under standard conditions (298K and 100kpa), all
reactants and products being in their standard states

67
Q

what is the enthalpy of formation of an element

68
Q

standard enthalpy change of combustion definition

A

the enthalpy change that occurs when one mole of a substance is combusted completely in oxygen under standard conditions (298K and 100kPa), all reactants and products being in their
standard states

69
Q

enthalpy change of neutralisation definition

A

the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water.

70
Q

what does incomplete combustion lead to

A

soot (carbon)
carbon monoxide
water
it will be LESS exothermic than complete combustion

71
Q

calorimetric method
- method

A

1.
- wash the equipment with the solutions to be used
- dry the cup after washing
- put polystyrene cup in a beaker for insulation and support
- measure out desired volumes of solutions with volumetric pipettes and transfer to insulated cup
- clamp thermometer into place making sure thermometer bulb is immersed in the solution
- measure the initial temperature of the solution, do this every minute for 2-3 minutes
- transfer second reagent to cup
- stir mixture
- record temp every minute after addition for 5 minutes

72
Q

errors in calorimetry experiments

A
  • energy transfer from surroundings (usually loss)
  • approximation in specific heat capacity of solution
  • neglecting specific heat capacity of calorimeter
  • reaction may be incomplete
  • density of solution is taken to be the same as water
73
Q

energy change equation

A

q(J) = m(g) x c x △T(K)

74
Q

errors in combustion experiments using calorimetry

A

Energy losses from calorimeter
* Incomplete combustion of fuel
* Incomplete transfer of energy
* Evaporation of fuel after weighing
* Heat capacity of calorimeter not included
* Measurements not carried out under standard conditions as H2O is gas, not liquid, in this experiment

75
Q

mean bond enthalpy definition

A

the enthalpy change when
one mole of bonds of (gaseous covalent) bonds is broken (averaged over different molecules)

76
Q

why are mean bond enthalpies positive

A

as energy is required to break a bond

77
Q

enthalpy change =

A

sum of bonds broken - sum of bonds made

78
Q

hess’ law definition

A

total enthalpy change for a reaction is independent of the route by which the chemical changes takes place

79
Q

hess’ law combustion equation

A

reactants - products

80
Q

hess’ law formation equation

A

products - reactants

81
Q

activation energy definition

A

the minimum energy which particles need to collide to start a reaction

82
Q

effect of increasing concentration and pressure on rate of reaction

A

more particles per unit volume
particles collide with a greater frequency
higher frequency of successful and effective collisions

83
Q

rate of reaction definition
units?

A

the change in concentration of a substance in unit time
units = moldm-3s-1

84
Q

on a graph of concentration vs time what
- does gradient show
- is initial rate
- how can reaction rates be calculated
- steeper gradient =

A
  • rate of reaction
  • the rate at the start of the reaction where it is the fastest
  • drawing a tangent and calculating the gradient of the tangent
  • faster rate
85
Q

catalyst definition

A

-increase reaction rates without getting used up
- they do this by providing an alternative route with a lower activation energy so more molecules gave energy above the AE

86
Q

heterogeneous catalyst

A
  • in a different phase from the reactants
  • reaction occurs at the surface of the catalyst
87
Q

homogeneous catalysts

A
  • in the same phase as the reactants
  • reaction proceeds through an intermediate species
88
Q

benefits of catalysts

A
  • speed up the rate of reaction. the use of a catalyst means lower temperatures and pressures can be used
  • this can save energy costs as there is reduced energy demand and less electrical pumping costs
  • this means fewer co2 emissions from burning of fossil fuels
  • catalysts can also enable different reactions to be used, with better atom economy and with reduced waste, of fewer undesired products and less use of hazardous substances

-Catalysts are often enzymes, generating very specific products, and operating
effectively close to room temperatures and pressures

89
Q

negatives of catalysts

A

some catalysts are toxic

90
Q

what does a maxwell boltzmann distribution graph show

A

the spread of energies that molecules of a gas or liquid have at a particular temperature

91
Q

How can a reaction go to completion if few particles have energy greater than EA?

A

particles can gain energy through collisions

92
Q

effect of increasing temp on Boltzmann distribution graph

A
  • at higher temps the energy of the particles increase. They collide more frequently and more often with energy greater than the activation energy. more collisions results in a reaction
  • As the temperature increases, the graph shows that a significantly bigger proportion of particles have energy greater than the activation energy, so the frequency of successful collisions
    increases
93
Q

effect of increasing surface area on rate of reaction

A

Increasing surface area will cause collisions to occur more frequently between the reactant particles and this increases the rate of the reaction.

94
Q

effect of catalysts on rate of reaction

A

If the activation energy is lower, more particles will have energy > EA, so there will be a higher
frequency of effective collisions. The reaction will be faster

95
Q

when does dynamic equilibrium occur and what does it mean

A
  • occurs when forward and backward reactions are occurring at a equal rate
  • the concentrations of reactants and products stay constant and the reaction is continuous
96
Q

effect of temperature on equilibrium

A
  • If temperature is increased the equilibrium will shift to oppose this and move in the endothermic direction to try to reduce the temperature by absorbing heat.
  • If temperature is decreased the equilibrium will shift to oppose this and move in the
    exothermic direction to try to increase the temperature by giving out heat.
97
Q

effect of pressure on equilibrium

A

Increasing pressure will cause the equilibrium to shift
towards the side with fewer moles of gas to oppose
the change and thereby reduce the pressure.
(and vice versa for decreasing pressure)

  • If the number of moles of gas is the same on both
    sides of the equation then changing pressure will have
    no effect on the position of equilibrium
98
Q

benefits and negatives with increasing pressure

A
  • gives a higher yield of product and produces a faster rate
  • high pressures are expensive to produce due to high energy costs for pumps and the equipment is expensive
99
Q

effect of catalysts on equilibrium

A

A catalyst has no effect on the position of equilibrium, but it will speed up the rate at which the equilibrium is achieved.

It does not effect the position of equilibrium because it speeds up the rates of the forward and backward reactions by the same amount.

100
Q

haber process
- equation

A

N2 + 3H2 =2NH3
-ve (exo)
low temp gives good yield but slow rate so compromise temp used
high pressure gives good yield and high rate: too high a pressure would lead to too high energy costs for pumps to produce the pressure

101
Q

kc =

A

equilibrium constant
only include gases

102
Q

what changes Kc

A

TEMPERATURE

103
Q

larger kc =

A

greater amount of products
if kc is small, equilibrium favours the reactants

104
Q

describe the solubility of magnesium and barium hydroxide

A
  • sparingly soluble
  • soluble
105
Q

solubility of group 2 metal sulfates down the group

A

solubility decreases down the group

106
Q

what happens to second ionisation energy down the group

A

it decreases

107
Q

why does oxidising power of the halogens decrease down the group

A

shielding increases and therefore it is more difficult to accept an extra electron

108
Q

are neutralisation reactions endothermic or exothermic

A

exothermic

109
Q

describe a similarity in terms of structure between copper and graphite

A

they both have layers of atoms that can slide over one another

110
Q

explain the difference in boiling point between carbon dioxide and diamond

A
  • diamond has a giant covalent structure
  • carbon dioxide has a simple molecular structure
  • diamond only contains strong covalent bonds, co2 contains weak IMF
  • covalent bonds are stronger than the IMF, hence require more energy to break
111
Q

suggest whether the mg2+ or na+ ion has a smaller ionic radius

A

mg2+
as it has a higher nuclear charge and same shielding

112
Q

state the meaning of a closed system

A

a system where none of the reactants or products can escape

113
Q

state le chatelier’s principle

A

if a system of equilibrium is disturbed, then the position of equilibrium will move to counteract this change

114
Q

give one key condition which must be satisfied for a reaction to reach dynamic equilibrium

A

must be a closed system

115
Q

how can a student check to make sure their reaction mixture has reached equilibrium

A
  • take several samples and test them over a period of time
  • to make sure they got the same results each time
116
Q

is the enthalpy sign -ve or +ve for these types of reaction
- formation
- combustion
- neutralisation

A
    • or -ve
  • negative
  • negative
117
Q

give the formula for calculating the standard enthalpy change of reaction, using bond energies

A

= sum of bonds broken - sum of bonds made

118
Q

why would H2 not have an enthalpy of formation value

A
  • it is an element and elements do not have enthalpy of formation values
119
Q

often when a bond enthalpy is calculated, it is a different value to the one which is quoted in a data book. why?

A

the data book value is an average value derived from a number of different compounds

120
Q

explain in terms of bond breaking and forming, why a reaction is exothermic

A

bond breaking absorbs energy and bond forming releases energy
- more energy is released than absorbed

121
Q

describe how you could measure out a mass of water without using a balance

A
  • measuring a volume of 50cm3
  • because water has a known density of 1.0gcm-3
122
Q

explain why most collisions in the gas phase do not result in a reaction taking place

A

molecules do not have enough energy

123
Q

give a reason why a reaction may be slow at room temp

A

activation energy is too high

124
Q

why does a catalyst have no effect on the yield of the products

A

a catalyst increases the rate of the forward and reverse reactions to the same extent

125
Q

why does the gradient of the curve decrease as the time of the reaction progresses

A
  • the concentration of reactants decreases
  • the frequency of successful collisions decreases
126
Q

give a medical use of barium sulfate and state why it is safe to use, despite solutions containing Ba2+ ions being toxic

A
  • it is taken before a patient has an x-ray
  • safe because it is insoluble
127
Q

state the trends in solubility of the group 2 sulfates and the group two hydroxides as you descend the group

A

sulfates = solubility decreases
hydroxides = solubility increases

128
Q

the trend in solubility and strength as a base can also be observed going down group 2
- explain how these two trends are connected

A
  • solubility and strength of a base increases
  • increasing the solubility means there are more hydroxide ions in solution
  • the hydroxide ion makes the metal hydroxide acts as a base
129
Q

explain how the bonding in berylium chloride differs from the rest of the group 2 chlorides

A

predominately covalent
Be2+ has the highest mass-charge ratio
so the be ion polarises the chloride ion
causing the chloride ion to share its electrons

130
Q

describe the properties of magnesium hydroxide that make it suitable for use in the human body

A
  • it is sparingly soluble
  • weak alkali
131
Q

why is calcium hydroxide used rather than magnesium hydroxide to increase the pH of acidic soils

A
  • more soluble than magnesium hydroxide
  • higher concentration of OH- ions
132
Q

suggest why adding chlorine to the drinking water supply is safe to do so

A
  • it is only added in small amounts
  • therefore not harmful to humans in these levels
133
Q

reducing agent definition

A

a species that can donate electrons

134
Q

explain why an iodide ion is a better reducing agent than a bromide ion

A
  • an iodide ion is larger than a bromide ion
  • outermost electrons are more shielded by inner shells
  • outermost electrons are more easily lost
135
Q

how are elements arranged in the periodic table

A
  • in increasing atomic number
  • in periods showing repeating trends in physical and chemical properties
  • in groups having similar chemical properties
136
Q

outer shell electron configuration of group 2 and the halogens

A

group 2:s2
halogens: s2,p5

137
Q

state two ways that the use of catalysts helps chemical companies to make their processes more sustainable and less harmful to the environment.

A
  • reaction can be carried out at lower temperatures
  • less fossil burnt/less co2 emissions
138
Q

What is the shape around the carbon atoms in graphene?

A

trigonal planar

139
Q

Explain the differences in the melting points of sodium and magnesium, using the model of
metallic bonding.

A

magnesium has more outer electrons
magnesium ions have a greater positive charge
magnesium has a greater attraction between ions and delocalised electrons
therefore, Mg has a higher boiling point

140
Q

Give chemical explanations for the following statements.

Potassium is placed immediately after argon in the periodic table.

A

potassium atoms have one more proton than argon

141
Q

oxidation numbers in ch3cooh

A

CH₃ carbon = -3

COOH carbon = +3

Oxygen = -2

Hydrogen = +1

142
Q

State an example of a catalyst used by the chemical industry and write the equation for the
reaction that is catalysed.

A

Nickel

C2H4 + H2 → C2H6

143
Q

Suggest whether the enthalpy change of vaporisation of bromine is exothermic or
endothermic.

A

endothermic as energy is required to overcome induced dipole dipole forces

144
Q

Why do Br2 and I2 not exist in the gaseous state under standard conditions

A

because energy is needed to break IDD interactions