Chem Flashcards
1
Q
How can random uncertainties be reduced?
A
Taking repeated readings and calculating the average value
2
Q
How can systematic errors be reduced?
A
Subtracting the error or replacing faulty equipment
3
Q
Random uncertainties lead to a loss of ___ and are reflected in ___
A
- Precision
- Percentage uncertainty
4
Q
Systematic errors lead to a loss of ___ and are reflected in ___
A
- Accuracy
- Percentage error
5
Q
Define precision and accuracy
A
- Precision: How close repeated readings are to one another
- Accuracy: How close readings are to true value
6
Q
If calculations involve addition and subtraction, final answer should be quoted to same ___ as piece of original data with fewest ___
A
Decimal places
7
Q
If calculations involve multiplication and division, final answer should be quoted to same ___ as piece of original data with fewest ___
A
Significant figures
8
Q
Define an element
A
Pure compound containing only 1 type of atom
9
Q
Define an atom
A
Smallest part of an element that can still be recognised at that element
10
Q
Define a compound
A
Pure substance formed when 2 or more elements combine chemically in a fixed ratio
11
Q
Define a mixture
A
2 or more substances physically mixed together in an unfixed ratio
12
Q
Give the formula for number of molecules
A
Moles x 6.02 x 10^23
13
Q
Give properties of a mixture
A
- Retain individual properties
- Can be separated by physical means
14
Q
Give the formula for number of atoms
A
Moles x No. of atoms x 6.02 x 10^23
15
Q
Give the formula for mass of 1 molecule
A
Molar mass / 6.02 x 10^23
16
Q
Give the formula for volume of gas (dm^3) at STP
A
Amt of gas (mol) x 22.7
17
Q
Give the formula for the relationship between the concentration, volume and moles of two substances
A
C1V1/N1 = C2V2/N2
18
Q
Give the formula for % composition
A
(no. of atoms x relative atomic mass)/relative molecular mass * 100
19
Q
Define empirical and molecular formula
A
- Empirical formula: Simplest ratio of different types of element in a compound
- Molecular formula: Actual number of each elements present in a molecule of the compound
20
Q
Give the formula for % yield
A
Theoretical yield/Actual yield * 100
21
Q
Give two assumptions in an ideal gas
A
- Volume of gas is negligible compared to volume of gas container
- Negligible intermolecular forces of attraction between gas particles
22
Q
Give two conditions for achieving an ideal gas
A
- Low pressure –> Gas particles are far apart, resulting in insignificant volume of gas compared to volume of gas container
- High temperature –> Gas particles possess high kinetic energy to overcome intermolecular forces of attraction between them
23
Q
Give the ideal gas law
A
pV = nRT
- P = Pressure (Pa)
- V = Volume (m^3)
- n = Amt (mol)
- R = 8.31 J k^-1 mol^-1
- T = Kelvin
24
Q
Give the combined gas law
A
P1V1/T1 = P2V2/T2
- Units for P and V: Any as long as they are consistent
- Units for T: Kelvin
25
Give the formula for concentration in ppm
- (Mass of solute * 10^6)/Mass of solution
- (Volume of gas * 10^6)/Volume of air
26
Define hydrated and anhydrous salt
- Hydrated: Salt with water
- Anhydrous: Salt without water
27
Give the formula for total concentration of ions when ionic substances dissolve in water
Original concentration * No. of ions
28
Define mass number
Number of protons and neutrons in the nucleus
29
Define atomic number
- Number of protons in the nucleus and thus the number of electrons
- Unique to an element
30
Define an isotope
Different atoms of the same element with the same number of protons but different number of neutrons in the nucleus
- Same chemical properties as chemical properties depend on the number and arrangement of electrons
- Different physical properties because their mass is different e.g. mass, speed of atoms
31
Define relative atomic mass
Weighted mean mass of all naturally-occurring atoms of an element, relative to 1/12 of the mass of a carbon-12 atom
- (% abundance * mass) + ... + (% abundance * mass)/100
32
Why are isotopes radioactive
- Nuclei break down spontaneously
- When they break down, they emit radiation which is dangerous to living things
33
Energy levels/Shells ___ in energy and are ___ in distance as they get further from the nucleus.
- Increase
- Closer
34
Give the formula for the number of electrons in each energy level/shell
2n^2
35
Give the order of the electromagnetic spectrum
RMIVUXG
- For Visible region: ROYGBIV
36
Longer wavelength, ___ frequency, ___ energy
Lower
37
Explain the emission spectra
- When energy is supplied to an atom, electrons are excited from their ground state to an excited state
- (Electrons can only exist in certain fixed energy levels) --> Electrons will be unable to exist in the higher level and fall to a lower energy level
- When electrons drop to a lower energy level they emit energy which corresponds to a particular wavelength and shows up as a line in the spectrum
38
When electrons return to __, they show up as a line in which region?
- n = 1
- n = 2
- n = 3
- n = 1: Ultraviolet region (largest energy change)
- n = 2: Visible region
- n = 3: Infrared region
39
Why do lines in the electromagnetic spectrum converge?
Because energy levels themselves converge. Beyond the continuum, electrons can have any energy and are free from the influence of the nucleus
40
Why do transition metals show different coloured flames when burnt?
- Heat causes electrons to move to a higher energy level
- Electrons are unstable in the higher energy level and fall back down
- When they fall from a higher to lower energy level, energy is emitted in the form of visible light energy (with the wavelength of the observed energy/colour)
41
Give the difference between a line spectrum and continuous spectrum
- Line spectrum: Only has specific wavelengths
- Continuous spectrum: Has all wavelengths
42
Give the number of orbitals in the s, p, d, f subshells
s:1, p:3, d: 5, f:7
43
Is the electronic configuration of the chromium ion (transition metal) [Ar]4s13d5 or [Ar]4s23d6
[Ar]4s13d5
44
When transition metals form ions, the electrons are removed from the ___ orbital first.
S orbital
45
Give the formula for effective nuclear charge
Number of protons - Shielding/inner electrons
46
Describe and explain the change in atomic radius across a period
Decreases because
- Electrons are added to the same valence shell
- But number of protons increase so effective nuclear charge increases
- Valence shells are more strongly attracted to the nucleus
47
Describe and explain the change in atomic radius down the group
Increases because
- Outermost electron is in a valence shell further from the nucleus
- Attraction of nucleus for electrons decreases
48
Describe the difference in radius between a parent atom and its cation
The radius of the cation will be smaller because
- The cation has less electrons but the same number of protons
- Valence electrons are more strongly attracted to the nucleus
- Number of electron shells also decrease
49
Describe the difference in radius between a parent atom and its anion
The radius of the anion will be larger because
- The anion has more electrons but the same number of protons
- Valence electrons are less strongly attracted to the nucleus
50
Given that they have the same number of electrons, rank them from the smallest to largest radius and explain why: Noble gas atom, cation, anion
Cation < Noble gas < Anion
- Same number of electrons but an increasing number of protons
- Effective nuclear charge increases
- Valence electrons are more strongly attracted to the nucleus
51
Define first ionisation energy and give a chemical equation representing it
- Energy required to remove an electron from an atom in a gaseous state (kJ mol^-1)
- X (g) --> X+ (g) + e-
52
Describe and explain the change in first ionisation energy across a period
Increases because
- Number of protons increase but number of inner electrons remain the same
- Effective nuclear charge increases
- Valence shell is more strongly attracted to the nucleus
53
Describe and explain the change in first ionisation energy down the group
Decreases because
- Electron to be removed is in a valence shell further from the nucleus
- Attraction of nucleus for electrons decreases
54
Explain why boron (1s2,2s2,2p1) has a lower ionisation energy than beryllium (1s2,2s2) even though boron has a higher effective nuclear charge
- Electron to be removed from boron is in the 2p-sublevel while electron to be removed from beryllium is in the 2s-sublevel
- The p-sublevel is higher in energy than the s-sublevel
- Less energy will be required to remove an electron from boron
55
Explain why oxygen (1s2,2s2,2p4) has a lower ionisation energy than nitrogen (1s2,2s2,2p3) even though oxygen has a higher effective nuclear charge
- Oxygen has 2 electrons paired up in the same p-orbital but nitrogen does not
- It will be easier to remove one of the paired electron from oxygen due to interelectronic repulsion
56
Define electron affinity and give a chemical equation representing it
- Energy change/enthalpy change when one electron is added to each atom in gaseous state
- X (g) + e- --> X- (g)
57
Describe and explain the change in electron affinity across a period
Increases because
- Increase in effective nuclear charge and decrease in atomic radius
- Electron will be more strongly attracted when brought into the valence shell
58
Describe and explain the change in electron affinity down Group 17
Generally decreases because
- Electron is brought into the valence shell
- As size of atom gets bigger, there is weaker attraction between added electron and nucleus as it is brought into a position further from the nucleus
59
Define electronegativity
Measure of the attraction an atom has for a shared pair of electrons when it is covalently bonded to another atom
60
Describe and explain the change in electronegativity across a period
Increases because
- Number of protons increase but number of inner electrons remain the same
- Effective nuclear charge increases
- Stronger attraction of bonding electrons to nucleus
61
Describe and explain the change in electronegativity down the group
Decreases because
- Increase in number of valence shells
- Bonding electrons are further from the nucleus
- Attraction of nucleus for bonding electrons decreases
62
Explain why metals have low electronegativity
Metals have few valence electrons and increase their stability by losing electrons
63
Describe and explain the change in melting and boiling points down Group 1
Decreases because
- Ions get larger
- Weaker electrostatic forces of attraction between positive ions in lattice and delocalised electrons
- Less energy required to overcome electrostatic forces of attraction
64
Describe and explain the change in reactivity down Group 1
Reactivity increases/ Reactions become more vigorous as
- Size of atom increases
- Ionisation energy decreases
- Outer electrons are lost more easily
65
Describe and give the equation for the reaction between a Group 1 element and oxygen
- React vigorously with oxygen and tarnish rapidly in air
- 4M (s) + O2 (g) --> 2M2O (s)
66
Give the equation for the reaction between a Group 1 element and water
2M (s) + H2O (l) --> 2MOH (aq) + H2 (g)
67
Give some observations for the reaction between a Group 1 element and water
- React rapidly
- Produces strong base
- Effervescence
- Solution gets hot
- Solution becomes basic
- Metal hydroxide/hydrogen gas produced
68
Describe the change in melting and boiling points down Group 17
Increases because
- Molecular mass increases
- London dispersion forces of attraction get stronger
- More energy required to overcome stronger intermolecular forces of attraction
69
Describe the change in reactivity down Group 17
Decreases because electron affinity decreases
70
When a reactive halogen reacts with a halide ion of a less reactive halogen, what happens?
More reactive halogen displaces the halide ion
71
Describe change in metallic properties across Period 3 oxides (sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine, and argon)
- Sodium, Magnesium, Aluminium: Metal
- Silicon: Metalloid
- Phosphorus, Sulfur, Chlorine, Argon: Non-metal
72
Describe the solubility of the Period 3 oxides (sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine, and argon)
- Sodium, Magnesium: Soluble
- Aluminium, Silicon: Insoluble
- Phosphorus, Sulfur: Soluble
73
Describe change in acid-base properties across Period 3 oxides (sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine, and argon)
- Sodium, Magnesium: Basic
- Aluminium: Amphoteric
- Silicon, Phosphorus, Sulfur: Acidic
74
Describe change in melting points across Period 3 (sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine, and argon)
Sodium - Aluminium: Increase
- Metallic bonding
- Increase in number of valence electrons and decrease in size
Silicon: High melting point
- Giant covalent structure
Phosphorus - Chlorine: S8 > P4 > Cl2
- P4, S8, Cl2
- Higher molecular mass
- Stronger intermolecular forces of attraction
Argon: Low melting point
- Weak intermolecular forces of attraction
75
Define metallic bonding
Electrostatic forces of attraction between a lattice of positive metal ions and sea of delocalised electrons
76
Describe structure of metallic bonding
Giant metallic structure consisting of lattice of positive metal ions and sea of delocalised electrons
77
Describe physical properties of metallic bonding
- Lustrous
- Good conductor of electricity because of presence of delocalised electrons
- Good conductors of heat
- Ductile
- Malleable
78
Describe and explain the melting and boiling points of metallic bonding
High as large amount of energy is required to overcome strong electrostatic forces of attraction between positive ions and sea of delocalised electrons
- Increases with size of cation, charge of cation, number of delocalised electrons
79
Describe the solubility of metallic bonding
Insoluble in both polar and non-polar solvent
80
Define alloys and explain why alloys are stronger and stiffer than pure metals
- Homogenous mixtures of two or more metals, or a metal and non-metal
- Introduction of a larger atom makes it harder for planes of atoms/ions to slide over one another
81
Define ionic bonding
- Electrostatic forces of attraction between oppositely charged ions
- Formed when electrons are transferred from one atom to another to form ions with noble gas configuration
82
Describe the structure of ionic bonding
- Giant crystal lattice
- Consists of oppositely charged ions held together by ionic bonds. Ions are arranged such that attraction between oppositely charged ions is maximum and repulsion between similarly charged ions is minimum
83
Describe physical properties of ionic bonding
- High melting and boiling point (Strong electrostatic forces of attraction)
- Hard and brittle (Ionic bonds are strong and non-directional; Ordered arrangement of ions can be easily dislocated when sheer stress is applied)
- Non-electrical conductor in solid state but good conductor when molten or dislocated
- Soluble in polar solvent but insoluble in non-polar solvent
84
Explain why ionic compounds are good conductors when molten or dissolved in aqueous solution but are non-conductors in solid state
- In solid state, there is the absence of mobile charge carriers as ions are held in fixed positions by strong electrostatic forces of attraction in giant ionic lattice
- In molten or aqueous state, there are mobile ions as lattice is broken down and ions are free to move
85
Define covalent bonding
- Electrostatic forces of attraction between the nuclei and the shared pair/s of electrons between them
- Formed when atoms share one or more pairs of electrons so each atom can achieve a noble gas configuration
86
Describe a simple molecular structure
- Atoms are held by strong covalent bonds but molecules are held together by weak intermolecular forces of attraction
87
Describe a giant molecular structure
- Atoms are bonded together by strong covalent bonds to give a giant three dimensional structure
88
Describe the melting and boiling points of covalent compounds
- Simple molecular structure: Low as low amount of energy is required to overcome weak intermolecular forces
- Giant molecular structure: High as large amount of energy is required to overcome strong intermolecular forces
89
Describe the solubility of covalent compounds
- Simple molecular structure: Polar molecules soluble in polar solvent, non-polar molecules soluble in non-polar solvent
- Giant molecular structure: Insoluble in both polar and non-polar solvents
90
Describe the electrical conductivity of covalent compounds
Simple molecular structure:
- Generally do not conduct electricity in any physical state
- Absence of mobile charge carriers as valence electrons are localised in covalent bonds
Giant molecular structure:
- Insoluble in both polar and non-polar solvents
- Exception: Graphite due to presence of delocalised electrons within layers
91
Describe the structure and bonding of silicone dioxide (SiO2)
- Tetrahedral
- Each silicon atom is covalently bonded to 4 oxygen atoms and each oxygen atom is covalently bonded to 2 silicon atoms
92
Describe the structure and bonding of diamond (C)
- Tetrahedral
- Each carbon atom is covalently bonded to 4 other carbon atoms
93
Describe the structure and bonding of silicon (C)
- Tetrahedral
- Each carbon atom is covalently bonded to 4 other carbon atoms
94
Describe the structure and bonding of graphite
- Trigonal planar
- Each carbon atom is bonded to 3 carbon atoms
- Honeycomb structure within each layer
- Continuous layers of hexagons
95
Explain why graphite is soft and slippery
Two-dimensional layers are held by weak intermolecular forces of attraction and layers can slide over one another when force is applied
96
Describe the physical properties of graphite
- Soft and slippery
- Conductor of electricity
97
Explain why graphite is a good conductor of electricity
Each carbon bonds to 3 other atoms. One valence electron per carbon atom can delocalise within the layer. The delocalised electrons can act as charge carriers.
98
Describe the bonding of graphene
- Each carbon atom is covalently bonded to 3 other carbon atoms
99
Describe the structure and bonding of buckminsterfullerene
- Each carbon atom is bonded covalently to 3 others
- Weak intermolecular forces of attraction between molecules that require little energy to overcome
- Slippery
100
Describe the direction of the dipole moment in a polar molecule
From less electronegative to more electronegative
101
Explain why diatomic molecules are non-polar
Both atoms have the same electronegativity value. Both atoms will exert identical attraction thus the electron pair is shared equally.
102
Give the bond angles from
- Linear
- Trigonal planar
- Bent
- Tetrahedral
- Trigonal pyramidal
- Bent
- 180
- 120
- 117.5
- 109.5
- 107
- 105
103
Explain London dispersion forces
- Electrons in an atom are in constant motion and at any one time will not be symmetrically distributed about nucleus
- This results in a temporary instantaneous dipole in the atom which will induce an opposite dipole in a neighbouring atom
- Dipoles attract each other so that there is an attractive force between atoms
104
Explain dipole-dipole forces
- Electrostatic forces of attraction between polar molecules
105
Explain hydrogen bonding
- Occurs when hydrogen is bonded directly to a small highly electronegative element (e.g. O, N, F)
- As the electron pair is drawn away from hydrogen atom by the electronegative element, all that remains is the proton in the nucleus as there are no inner electrons
- Proton attracts a non-bonding pair of electrons from O, N or F resulting in a much stronger dipole-dipole attraction
106
Explain why overall force between atoms are always attractive if dipoles constantly disappear and reappear
A dipole always induce an opposite one
107
In an exothermic reaction, heat is __ surroundings.
Released to
108
In an endothermic reaction, heat is __ surroundings
Absorbed from
109
In an exothermic reaction, bonds in product are __ than bonds in reactant and the product is more __ than reactants
Stronger
Stable
110
Define enthalpy change
Amount of heat energy absorbed/evolved in a reaction
111
Define standard enthalpy change of reaction. Give the formula to find enthalpy change of reaction if the enthalpy change of formation of products and reactants are known
- Enthalpy change when molar amounts of reactants react together to form products under standard conditions
- Enthalpy change of formation of products - Enthalpy change of formation of reactants
112
Define standard enthalpy change of combustion
Enthalpy change when 1 mole of substance is completely burnt in oxygen under standard conditions
113
Define standard enthalpy change of formation
Enthalpy change when 1 mole of compound is formed from its elements under standard conditions
114
Define specific heat capacity (c)
Energy required to raise temperature of 1kg of substance by 1K
115
Define heat capacity (C)
Energy required to raise temperature by 1K
116
Give some sources of error in a calorimeter and explain how to minimise these errors
- Heat loss to surroundings. Insulate reaction vessel or use a lid
- Calorimeter could absorb some of the heat evolved. Determine the specific heat capacity of the calorimeter and add the heat absorbed by the calorimeter to heat absorbed by water to find total heat evolved by burning of fuel
- Incomplete combustion (formation of soot and carbon monoxide)
117
Give possible assumptions when calculating enthalpy change
- Density of all aqueous solutions is assumed to be 1 g cm-1
- Specific heat capacity is assumed to be the same as water
118
Define bond enthalpy and give the formula for calculating enthalpy change from the bond enthalpy values of the product and reactant
- Energy required to break one mole of a covalent bond in a gaseous molecule averaged over similar compounds
- BE (bonds broken) - BE (bonds formed)
119
Bond breaking is ___ (exothermic/endothermic).
Bond forming is ___ (exothermic/endothermic)
- Bond breaking is endothermic
- Bond forming is exothermic
120
Give the limitations of using bond enthalpy values
- Values are average values and are not specific to the bond broken for that reaction
- Values only apply to bonds in gaseous state
121
Define rate of reaction
Increase in concentration of one of the products per unit time or decrease in concentration of one of the reactants per unit time
122
Define the collision theory
For a reaction to occur between two particles, particles must
- Collide with the appropriate geometry so the reactive parts of the particle come into contact with each other
- Must collide with sufficient energy to bring about the reaction
123
Define activation energy
Minimum amount of energy required for a reaction to take place
124
Describe the effect of increasing temperature on rate of reaction
- Increase in rate
- As temperature increases, kinetic energy of particles increase
- Greater number of particles with energy greater than or equal to the activation energy
- Frequency of effective collisions increase
125
Describe the effect of increasing surface area on rate of reaction
- Increase in rate
- Larger exposed surface area for reaction
- Higher frequency of effective collisions
126
Describe the effect of increasing concentration on rate of reaction
- Increase in rate
- Number of reacting particles per unit volume increases
- Frequency of effective collisions increases
127
Describe the effect of increasing pressure on rate of reaction
- Increase in rate
- Reactant particles are closer
- Frequency of effective collisions increases
128
Describe the effect of a catalyst on a reaction
- Provides an alternative pathway with lower activation energy
- Greater proportion of reacting particles colliding with energy greater than or equal to activation energy
- Frequency of effective collisions increases
- Greater rate of reaction
129
Define dynamic equilibrium
Equilibrium where
- The concentration of all the reactants and products are constant
- Rate of forward reaction is equal to the rate of the reverse reaction
130
Define closed system
One where neither matter nor energy can be lost or gained
131
If equilibrium constant is a large number, reaction will proceed to the ___
Right
132
If the reaction quotient is larger than the equilibrium constant, reaction will proceed to the ___
Left
133
Equilibrium constant can be affected by ___
Temperature
- If forward reaction is favoured, equilibrium constant increases
134
Define Le Chatelier's principle
- Provided temperature remains constant, value of equilibrium constant must remain constant
- If a system at equilibrium is subjected to a small change, equilibrium will shift to minimise the effect of the change
135
Describe the effect of temperature on the position of equilibrium
- Increase in temperature favours endothermic reactions
- Decrease in temperature favours exothermic reactions
136
Describe the effect of pressure on the position of equilibrium
- Increase in pressure will favour the reaction pathway leading to less volume (to minimise pressure)
- Reduction in pressure will favour the reaction pathway leading to greater volume
137
Describe the effect of concentration on the position of equilibrium
- Increasing the concentration of reactants will favour the forward reaction
138
Describe the effect of a catalyst on the position of equilibrium
- No effect on position of equilibrium as it lowers the activation energy for both reactions equally and increases the rate of both reactions equally
139
Define an acid
Substance that produces hydrogen ions, H+ (aq), in aqueous solution
140
Define a base
Substance that can neutralise an acid
141
Define an alkali
Base that is soluble in water
142
Define a Bronsted-Lowry acid
Substance that can donate a proton
143
Define a Bronsted-Lowry base
Substance that can accept a proton
144
Define a conjugate base
Species remaining after acid has lost a proton
145
Define a conjugate acid
Species remaining after base has accepted a proton
146
Define an amphoteric species
Substance that can donate or accept a proton
147
Give the equation for a reaction between acid and base
Acid + Base --> Salt + Water
148
Give the equation for a reaction between acid and reactive metal
Acid + Reactive Metal --> Salt + Hydrogen
* Reactive metal must be above copper
149
Give the equation for a reaction between acid and carbonate
Acid + Carbonate --> Salt + Carbon dioxide + Water
150
Give the equation for a reaction between acid and hydrogencarbonate
Acid + Hydrogencarbonate --> Salt + Carbon dioxide + Water
151
Define pH
- Measure of the concentration of H+ (aq) ions in a solution
- pH - log10 [H+ (aq)]
152
Define a strong acid and give examples
- Acid that completely dissociates into its ions in aqueous solution
- E.g. HCl, HNO3, H2SO4
153
Define a strong base and give examples
- Base that completely dissociates into its ions in aqueous solution
- E.g. NaOH, KOH, Ba(OH)2
154
Define a weak acid and give examples
- Acid that partially dissociates into its ions in aqueous solution
- Only 1 molecule in 100 dissociates
- E.g. CH3COOH, CO2 in H2CO3
155
Define a weak base and give examples
- Base that only partially dissociates into its ions in aqueous solution
- E.g. NH3, C2H5NH2
156
Give possible solutions to differentiate between a strong and weak acid/base
1. Measure pH using pH meter
2. Measure conductivity using a conductivity meter. Strong acids and bases will give higher readings because they contain more ions in solution
3. For acids: React with metals or carbonates. Strong acids will give a greater rate of reaction because the concentration of hydrogen ions is greater
157
Define VSEPR theory
Pairs of electrons arrange themselves around the central atom so they are as far apart as possible in order to minimise repulsion and obtain maximum stability
158
Give the order of repulsion for: Lone pair-lone pair, lone pair-bond pair, bond pair-bond pair
Lone pair-lone pair, lone pair-bond pair, bond pair-bond pair
159
Sulfur dioxide is produced from
- Volcanoes
- Combustion of sulfur-containing fossil fuels
- Smelting of sulfide ions
160
Give the reactions of sulfur that produce acid deposition
Formation of sulfur dioxide: S (s) + O2 (g) --> SO2 (g)
Oxidation of sulfur dioxide: SO2 (g) + 1/2 O2 (g) --> SO3
Reaction between sulfur dioxide and water in air:
- SO2 (g) + H2O (l) --> H2SO3 (aq) Sulfurous acid
- SO3 (g) + H2O (l) --> H2SO4 (aq) Sulfuric acid
161
Nitrogen oxides are produced from
- Electrical storms
- Internal combustion engine
- Jet engine
162
Give the reactions of nitrogen that produce acid deposition
Formation of nitrogen oxide: N2 (g) + O2 (g) --> 2NO (g)
Oxidation of nitrogen dioxide: 2NO (g) + O2 (g) --> 2NO2 (g)
Reaction between nitrogen dioxide and water: 2NO2 (g) + H2O (l) --> HNO3 (aq) + HNO2 (aq)
OR: Reaction between nitrogen dioxide and oxygen: 4NO2 (g) + O2 (g) + 2H2O (l) --> 4NO3 (aq)
163
Define acid deposition
Process by which acidic particles, gases and precipitation leave the atmosphere
164
Define acid rain
Rain with a pH less than 5.6
165
Give some effects of acid deposition
- Increased acidity in the soil can leech important nutrients such as Ca2+, Mg2+, K+
- Can kill fish
- Stone that contains calcium carbonate is eroded by acid rain and can dissolve away exposing a fresh surface to react with more acid
- Can irritate mucous membranes and increase the risk of respiratory diseases such as asthma
- Dissolve heavy metal compounds and release poisonous ions such as Cu2+ and Pb2+
166
Give some solutions for acid deposition
- Lower amounts of sulfur dioxides and nitrogen oxides by improving engine design
- Switch to alternate forms of energy such as wind and solar power
- Reduce amount of fuel by using public teansport
- Liming of lakes: Adding calcium oxide or calcium hydroxide to neutralise acidity
167
Define reducing agent
Species that reduces another while itself getting oxidised
168
Give some examples of reducing agents
- KI: Colourless I- --> Brown I2
- Iron (II) Salts: Pale green Fe2+ --> Yellowish-brown Fe3+
169
Define oxidising agent
Species that oxidises another while itself getting reduced
170
Give some examples of oxidising agents
- KMnO4: Purple MnO4- --> Colourless Mn2+
- K2CrO7: Orange Cr2O7^2- --> Green Cr3+
171
Explain the function of a salt bridge in a voltaic cell
Provides an electrical connection between half cells to complete the circuit and keep half-cells electrically neutral
172
Explain the movement of ions in a salt bridge in a voltaic cell
Anions move towards anode and cathodes move towards cathode
173
In a voltaic cell, the cathode is __ (negative/positive) and the anode is ___ (negative/positive)
Cathode is positive and the anode is negative
174
In a voltaic cell, at the cathode/positive cell, ___ occurs and ___ are produced
Reduction occurs and pure elements are produced
175
In a voltaic cell, at the anode/negative cell, __ occurs and ___ are produced
Oxidation occurs and cations are produced
176
Define electrolysis
Breaking down of a substance in molten state by the passage of electricity through it
177
Explain how electric current is produced in an electrolytic cell
Current is produced due to movement of electrons through wires and movement of ions in the electrolyte
178
In an electrolytic cell, the cathode is ___ (negative/positive) and the anode is ___ (negative/positive)
Cathode is negative and anode is positive
179
In an electrolytic cell, at the cathode/negative cell, ___ occurs and ___ are produced
Reduction occurs and pure elements are formed
180
In an electrolytic cell, at the anode/positive cell, ___ occurs and __ are produced
Oxidation occurs and pure elements are formed
181
Define a homologous series
- Series of compounds with the same functional group
- Each member has the same general formula
- Neighbouring members differ by -CH2-
182
Define a functional group
Atom/group of atoms in a molecule giving it its characteristic chemical properties
183
Give evidence for the resonance structure of benzene
- C-C bond lengths are all the same and have a value of 0.140nm which lies between the values for C-C and C=C
- Enthalpy of hydrogenation of cyclohexene is -120kJ mol-1. If benzene simply had the cyclohexa-1,3,5-triene structure with three double bonds the enthalpy change of hydrogenation of benzene would be 3 times the enthalpy change of hydrogenation of cyclohexene. The difference is the extra energy associated with the delocalisation
- Only one isomer exists for 1,2-disubstituted benzene compounds. If there were simply alternate bonds two isomers would exist
- If benzene had 3 normal double bonds it would be expected to readily undergo addition reactions. But it only undergoes addition reactions with difficulty and most commonly undergoes substitution reactions
184
Define structural isomers
Compounds with the same molecular formula but different structural formulas
185
Describe boiling points down a homologous series
- Down the homologous series, boiling points increase because carbon chain increases and London dispersion forces of attraction increase
- Increases steeply at first but as successive CH2 groups are added, the rate of increase decreases
186
Describe change in melting and boiling point when branching occurs
- Molecules become more spherical in shape and contact surface area between molecules decreases
- Melting point increases but boiling point decreases
187
What are the intermolecular forces of attraction in alkanes?
Only London dispersion forces of attraction
188
Explain the volatility of alkanes
Only have London dispersion forces of attraction
189
What are the conditions, reagents and products for a substitution reaction?
Conditions: Presence of sunlight or ultraviolet light to provide extra energy required to overcome activation energy
Reagents: Alkane, Halogen
Products: Hydrogen halide, substituted alkane
190
Give equations for the steps in a substitution reaction between methane and chlorine
Initiation:
Covalent bond breaks and each of the two atoms forming the bond retains one of the shared electrons, resulting in the formation of two free radicals
Cl2 --> 2Cl radical
Propagation:
CH4 + Cl radical --> HCl + CH3 radical
CH3 radical + Cl2 --> CH3Cl + Cl radical
Termination:
Cl radical + Cl radical --> Cl2
CH3 radical + Cl radical --> CH3Cl
CH3 radical + CH3 radical --> C2H6
191
Arrange in order of bond strength and bond length: Single bond, Double bond, Triple Bond
Single bond: Weakest strength, longest bond length
Double bond: Medium, medium bond length
Triple bond: Strongest, shortest bond length
Stronger attraction between nuclei and bonding electrons
192
Define an unsaturated compound
Contains carbon double or triple bonds
193
Describe what happens when an alkane and an alkene react with bromine solution
- Alkane and bromine solution: No reaction
- Alkene and bromine solution: Bromine solution changes colour from orange to colourless
194
Give the conditions for a reaction between an alkene and halogen
Conditions: Room temperature
195
Give the conditions and products for a reaction between an alkene and hydrogen
Conditions: Heat, finely divided nickel (catalyst)
Products: Alkanes
196
Give the conditions and products for a reaction between an alkene and halogen halides
Conditions: Heat
197
Give the conditions and products for a reaction between an alkene and water
Conditions: Heat, concentrated sulfuric acid/phosphoric acid (catalyst)
Products: Alcohols
198
Explain the process by which primary alcohols are oxidised
- Oxidised with acidified K2Cr2O7
- First to aldehydes and then to carboxylic acids
- To stop the reaction at the aldehyde stage, the ethanal is distilled from the reaction mixture as soon as it is formed
- If complete oxidation to carboxylic acid is required the mixture can be heated under reflux so none of the ethanal can escape
199
How can primary alcohols be oxidised to aldehydes?
To stop the reaction at the aldehyde stage, the ethanal is distilled from the reaction mixture as soon as it is formed
200
How can primary alcohols be oxidised fully to carboxylic acids?
If complete oxidation to carboxylic acid is required the mixture can be heated under reflux so none of the ethanal can escape
201
Explain the process by which secondary alcohols are oxidised
Oxidised to ketones and cannot undergo further oxidation
202
Explain the process by which tertiary alcohols are oxidised
Cannot be oxidised as they have no hydrogen atoms attached directly to the carbon atom containing the -OH- group
203
What reactions do halogenoalkanes undergo?
Nucleophilic substitution
204
Give the products for the nucleophilic substitution reaction between ammonia and C2H5Br
NH3 + C2H5Br → C2H5NH2 + HBr
205
What reactions does benzene undergo?
Electrophilic substitution
- High electron density so it can undergo substitution reactions with electrophiles
206
Explain why benzene does not readily undergo addition reactions
Delocalisation of electrons in benzene ring give it extra stability
207
Give the products for the electrophilic substitution reaction between C6H6 and Cl2
C6H6 + Cl2 --> C6H6Cl + HCl
208
Give the reaction name, conditions and products for a reaction between an alcohol and carboxylic acid
Reaction: Condensation
Conditions: Presence of small amount of concentrated sulfuric acid catalyst
Products: Ester, water
209
Give the IHD values for:
- A double bond
- A triple bond
- A ring
- Double bond and ring: 1 IHD
- Trilpe bond: 2 IHD
210
Explain how infrared spectroscopy works
- When molecules absorb energy in the infrared region of the electromagnetic spectrum they vibrate
- When infrared radiation is passed through a sample the spectrum shows the characteristic absorption
- Confirms that a particular bond is present
211
Infrared spectroscopy only works for ___ compounds
Polar covalent compounds
212
What region of the electromagnetic spectrum is used in 1H NMR
Radio waves
213
How can 1H NMR help to identify compounds?
- Provides information on the chemical environment of all the hydrogen atoms in a molecule
- Peaks/signals indicates number of chemical environments
214
Determine the bond angle in benzene
109.5
215
The specific heat capacity of X is twice that of Y. If the same amount of heat is supplied to X and Y, the temperature change of Y will be ___ of X.
Four times that
216
In a voltaic cell, negative ions flow through the salt bridge from the __ to the ___.
Cathode/positive half-cell to the anode/negative half-cell
217
Differentiate between a homogenous and heterogenous mixture
- Homogenous: Uniform composition and can be separated by filtration
- Heterogenous: Non-uniform composition and can be separated by chromatography, distillation, crystallisation, etc.
218
In a voltaic cell, electrons flow through the wires from the ___ to the ___.
Anode to the cathode
219
What is the difference between an amphiprotic and amphoteric substance?
Amphiprotic: Can donate or accept a proton
Amphoteric: Can act as an acid or a base
