Module 3 Flashcards
(144 cards)
1
Q
why do elements in the same group have similar chemical properties
A
- same outer shell electron configuration
2
Q
how are elements classified as s,p or d blocks
A
according to which orbitals the highest energy electrons are in
3
Q
order of blocks on the periodic table
A
s,d,p,f
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
5
Q
periodicity definition
A
repeating pattern across different periods
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
7
Q
first ionisation energy definition
A
the energy needed to remove one mole of electrons from one mole of gaseous atoms
8
Q
equation example for first ionisation energy
A
H(g) = H+(g) + e-
9
Q
factors affecting ionisation energy
A
- attraction of the nucleus
- more protons = greater attraction - distance of the electrons from the nucleus
- bigger atom means outer electrons are further from the nucleus, meaning weaker attraction to the nucleus - 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
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
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
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
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
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
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
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
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
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.
19
Q
metallic bonding definition
A
the electrostatic force of attraction between
the positive metal ions and the delocalised electrons
20
Q
what are the three main factors that affect the strength of metallic bonding
A
- Number of protons/ Strength of nuclear attraction.
(The more protons the stronger the bond) - Number of delocalised electrons per atom - the outer shell electrons are delocalised
(The more delocalised electrons the stronger the bond) - Size of ion.
(The smaller the ion, the stronger the bond)
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
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
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
24
Q
what is graphene
A
a single layer of graphite
25
structure of silicon oxide
- giant covalent lattice/ macromolecular
- tetrahedral
- each silicon is shared by four oxygens and each oxygen is shared by two silicons
- empirical formula = SiO2
26
which structures have metallic bonding and give two examples
GIANT METALLIC LATTICES
- magnesium and sodium
27
which structures have covalent bonding and give three examples
MACROMOLECULAR
- diamond, graphite, silicon dioxide
28
properties of macromolecular substances
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
29
properties of giant metallic substances
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.
30
explain the general trend of mp and bp across period three
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
31
what happens to atomic radius down group 2
- increases
- atoms have more shells of electrons, making the atom bigger
32
melting point down group 2
- 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
33
what happens to first and second IE's down group 2
- 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
34
second ionisation energy definition
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
35
example equation for 2nd IE
Ti+ (g) = Ti2+(g) + e-
36
reactivity down group 2
- 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
37
group 2 metals reactions with oxygen
- give equation
- what colour flame?
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
38
Mg + steam
- produces?
- what colour flame?
produces magnesium oxide and hydrogen
the Mg would burn with a bright white flame
39
Mg + warm water
- produces?
- flame and speed of reaction?
produces magnesium hydroxide and hydrogen
Mg + 2H20 → Mg(OH)2 + H2
this is a slower reaction than with steam and produces no flame
40
group 2 metals and cold water
- 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
41
what are the observations for group 2 metals reacting with cold water
- 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)
42
group 2 metals reacting with acid
give an example equation
- react with acids with increasing vigour down the group to form a salt and hydrogen
for example:
Ca + 2HCl(aq) = CaCl2(aq) + H2 (g)
43
why does barium react slowly with H2SO4
give the equation
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
44
group 2 oxides + water
- forms?
- forms hydroxides
45
why are group 2 oxides basic
- the oxide ions accept H+ ions to become hydroxide ions
46
what is magnesium hydroxide used for
- soluble?
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
47
what is calcium hydroxide used for
- is is soluble?
- 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
48
what is limewater?
what is it used for?
give equation
1. aqueous solution of calcium hydroxide
2. used as a test for co2
3. Ca(OH)2 + CO2 → CaCO3 + H2O
49
appearance of the halogens
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.
50
mp and bp down the halogens
- explain why
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
51
reactivity of the halogens going down the group
- 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
52
displacement reactions of halogens
look at chemsheets
53
disproportionation reaction definition
name of a reaction where an element simultaneously oxidises and reduces
54
chlorine + water equation
how is it disproportionation?
what will happen is UI is added?
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.
55
what are the benefits and negatives of chlorine being used in water treatment?
kills bacteria =good
toxic and irritates the respiratory system= bad
forms chlorinated hydrocarbons = bad
56
chlorine + cold dilute NaOH
- reaction
- colour change
- what are the products used for
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
57
chlorine + hot NaOH
3Cl2 + 6NaOH → NaClO3 + 5NaCl + 3H2O
sodium chlorate (V) is formed
- disproportionation reaction
58
reaction of halide ions with silver nitrate
- used for?
- why is nitric acid added?
- colours?
- how can ammonia be used to help differentiate
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
testing for the presence of a carbonate
- equation
- 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
testing for presence of a sulfate
- why can sulfuric acid not be used
- equation
- 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
how to test for an ammonium ion (positive ions/cations)
- how can ammonia gas be identified
react with NaOH(Aq), forming NH3 gas
- ammonia gas can be identified by its pungent smell or by turning red litmus paper blue
62
in what order should you test for ions
1. carbonate
2. sulfate
3. halide
- this prevents false results
63
is △H positive or negative in exothermic and endothermic reactions?
In an exothermic reaction the
∆H is negative
In an endothermic reaction
the ∆H is positive
64
activation energy definition
the minimum energy which particles need to collide to start a reaction
65
standard conditions
* 100 kPa pressure
* 298 K (room temperature or 25oC)
* Solutions at 1mol dm-3
* all substances should have their normal state at 298K
66
standard enthalpy change of formation definition
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
what is the enthalpy of formation of an element
zero
68
standard enthalpy change of combustion definition
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
enthalpy change of neutralisation definition
the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water.
70
what does incomplete combustion lead to
soot (carbon)
carbon monoxide
water
it will be LESS exothermic than complete combustion
71
calorimetric method
- method
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
errors in calorimetry experiments
- 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
energy change equation
q(J) = m(g) x c x △T(K)
74
errors in combustion experiments using calorimetry
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
mean bond enthalpy definition
the enthalpy change when
one mole of bonds of (gaseous covalent) bonds is broken (averaged over different molecules)
76
why are mean bond enthalpies positive
as energy is required to break a bond
77
enthalpy change =
sum of bonds broken - sum of bonds made
78
hess' law definition
total enthalpy change for a reaction is independent of the route by which the chemical changes takes place
79
hess' law combustion equation
reactants - products
80
hess' law formation equation
products - reactants
81
activation energy definition
the minimum energy which particles need to collide to start a reaction
82
effect of increasing concentration and pressure on rate of reaction
more particles per unit volume
particles collide with a greater frequency
higher frequency of successful and effective collisions
83
rate of reaction definition
units?
the change in concentration of a substance in unit time
units = moldm-3s-1
84
on a graph of concentration vs time what
- does gradient show
- is initial rate
- how can reaction rates be calculated
- steeper gradient =
- 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
catalyst definition
-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
heterogeneous catalyst
- in a different phase from the reactants
- reaction occurs at the surface of the catalyst
87
homogeneous catalysts
- in the same phase as the reactants
- reaction proceeds through an intermediate species
88
benefits of catalysts
- 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
negatives of catalysts
some catalysts are toxic
90
what does a maxwell boltzmann distribution graph show
the spread of energies that molecules of a gas or liquid have at a particular temperature
91
How can a reaction go to completion if few particles have energy greater than EA?
particles can gain energy through collisions
92
effect of increasing temp on Boltzmann distribution graph
- 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
effect of increasing surface area on rate of reaction
Increasing surface area will cause collisions to occur more frequently between the reactant particles and this increases the rate of the reaction.
94
effect of catalysts on rate of reaction
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
when does dynamic equilibrium occur and what does it mean
- 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
effect of temperature on equilibrium
- 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
effect of pressure on equilibrium
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
benefits and negatives with increasing pressure
- 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
effect of catalysts on equilibrium
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
haber process
- equation
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
kc =
equilibrium constant
only include gases
102
what changes Kc
TEMPERATURE
103
larger kc =
greater amount of products
if kc is small, equilibrium favours the reactants
104
describe the solubility of magnesium and barium hydroxide
- sparingly soluble
- soluble
105
solubility of group 2 metal sulfates down the group
solubility decreases down the group
106
what happens to second ionisation energy down the group
it decreases
107
why does oxidising power of the halogens decrease down the group
shielding increases and therefore it is more difficult to accept an extra electron
108
are neutralisation reactions endothermic or exothermic
exothermic
109
describe a similarity in terms of structure between copper and graphite
they both have layers of atoms that can slide over one another
110
explain the difference in boiling point between carbon dioxide and diamond
- 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
suggest whether the mg2+ or na+ ion has a smaller ionic radius
mg2+
as it has a higher nuclear charge and same shielding
112
state the meaning of a closed system
a system where none of the reactants or products can escape
113
state le chatelier's principle
if a system of equilibrium is disturbed, then the position of equilibrium will move to counteract this change
114
give one key condition which must be satisfied for a reaction to reach dynamic equilibrium
must be a closed system
115
how can a student check to make sure their reaction mixture has reached equilibrium
- take several samples and test them over a period of time
- to make sure they got the same results each time
116
is the enthalpy sign -ve or +ve for these types of reaction
- formation
- combustion
- neutralisation
- + or -ve
- negative
- negative
117
give the formula for calculating the standard enthalpy change of reaction, using bond energies
= sum of bonds broken - sum of bonds made
118
why would H2 not have an enthalpy of formation value
- it is an element and elements do not have enthalpy of formation values
119
often when a bond enthalpy is calculated, it is a different value to the one which is quoted in a data book. why?
the data book value is an average value derived from a number of different compounds
120
explain in terms of bond breaking and forming, why a reaction is exothermic
bond breaking absorbs energy and bond forming releases energy
- more energy is released than absorbed
121
describe how you could measure out a mass of water without using a balance
- measuring a volume of 50cm3
- because water has a known density of 1.0gcm-3
122
explain why most collisions in the gas phase do not result in a reaction taking place
molecules do not have enough energy
123
give a reason why a reaction may be slow at room temp
activation energy is too high
124
why does a catalyst have no effect on the yield of the products
a catalyst increases the rate of the forward and reverse reactions to the same extent
125
why does the gradient of the curve decrease as the time of the reaction progresses
- the concentration of reactants decreases
- the frequency of successful collisions decreases
126
give a medical use of barium sulfate and state why it is safe to use, despite solutions containing Ba2+ ions being toxic
- it is taken before a patient has an x-ray
- safe because it is insoluble
127
state the trends in solubility of the group 2 sulfates and the group two hydroxides as you descend the group
sulfates = solubility decreases
hydroxides = solubility increases
128
the trend in solubility and strength as a base can also be observed going down group 2
- explain how these two trends are connected
- 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
explain how the bonding in berylium chloride differs from the rest of the group 2 chlorides
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
describe the properties of magnesium hydroxide that make it suitable for use in the human body
- it is sparingly soluble
- weak alkali
131
why is calcium hydroxide used rather than magnesium hydroxide to increase the pH of acidic soils
- more soluble than magnesium hydroxide
- higher concentration of OH- ions
132
suggest why adding chlorine to the drinking water supply is safe to do so
- it is only added in small amounts
- therefore not harmful to humans in these levels
133
reducing agent definition
a species that can donate electrons
134
explain why an iodide ion is a better reducing agent than a bromide ion
- an iodide ion is larger than a bromide ion
- outermost electrons are more shielded by inner shells
- outermost electrons are more easily lost
135
how are elements arranged in the periodic table
- in increasing atomic number
- in periods showing repeating trends in physical and chemical properties
- in groups having similar chemical properties
136
outer shell electron configuration of group 2 and the halogens
group 2:s2
halogens: s2,p5
137
state two ways that the use of catalysts helps chemical companies to make their processes more sustainable and less harmful to the environment.
- reaction can be carried out at lower temperatures
- less fossil burnt/less co2 emissions
138
What is the shape around the carbon atoms in graphene?
trigonal planar
139
Explain the differences in the melting points of sodium and magnesium, using the model of
metallic bonding.
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
Give chemical explanations for the following statements.
Potassium is placed immediately after argon in the periodic table.
potassium atoms have one more proton than argon
141
oxidation numbers in ch3cooh
CH₃ carbon = -3
COOH carbon = +3
Oxygen = -2
Hydrogen = +1
142
State an example of a catalyst used by the chemical industry and write the equation for the
reaction that is catalysed.
Nickel
C2H4 + H2 → C2H6
143
Suggest whether the enthalpy change of vaporisation of bromine is exothermic or
endothermic.
endothermic as energy is required to overcome induced dipole dipole forces
144
Why do Br2 and I2 not exist in the gaseous state under standard conditions
because energy is needed to break IDD interactions
