3.1 Physical Flashcards
(192 cards)
1
Q
What is ionic bonding?
A
The electrostatic force of attraction between oppositely charged ions in a lattice.
2
Q
What is the formula for a sulfate ion?
A
SO4^2-
3
Q
What is the formula for a hydroxide ion?
A
OH-
4
Q
What is the formula for a nitrate ion
A
NO3^-
5
Q
What is the formula of a carbonate ion?
A
CO3^2-
6
Q
What is the formula for an ammonium ion?
A
NH4^+
7
Q
What are the 4 crystal structures?
A
Ionic
Metallic
Molecular
Macromolecular
8
Q
What is the definition of a single covalent bond?
A
A shared pair of electrons.
(Multiple bonds contain multiple pairs of electrons)
9
Q
What is the definition of a co-ordinate bond?
A
It contains a shared pair of electrons with both electrons supplied by one atom.
10
Q
What is the definition of metallic bonding?
A
The electrostatic attraction between delocalised electrons and positive ions arranged in a lattice.
11
Q
What are some examples of crystal structures?
A
Diamond
Graphite
Ice
Iodine
Magnesium
Sodium chloride
12
Q
What are the physical properties of ionic structures?
A
-They conduct electricity when molten or dissolved- the ions in the liquid are free to move.
-High melting points- they have a lattice structure so there are strong electrostatic forces which require lots of energy to overcome.
-Tend to dissolve in water- water molecules are polar, the charged particles pull ions away from the lattice.
13
Q
What is the definition of a macromolecular structure?
A
A huge network of covalently bonded atoms.
14
Q
What are the physical properties of graphite?
A
-Lubricant- there are weak bonds between layers which are broken easily so, the sheets can slide over each other.
-Electrical conductor- there are delocalised electrons which are free to move through the sheet.
-Low density- the layers are far apart so, it’s used for strong lightweight sports equipment.
-High melting point- strong covalent bonds.
-Insoluble- covalent bonds are too strong.
15
Q
What are the physical properties of diamond?
A
-Tetrahedral shape
-High melting point
-Hard
-Good thermal conductor- vibrations travel easily through the lattice
-Bad electrical conductor- Outer electrons are in localised bonds
-Insoluble
-Used for gemstones as it refracts light a lot
16
Q
What are the physical properties of metallic structures?
A
-Good conductors- sea of delocalised electrons to carry flow of charge.
-Malleable- the layer of positive Ions can slide over each other. The delocalised electrons prevent fragmentation as they move around the lattice.
-High melting point/solid at room temp- strong electrostatic attraction.
17
Q
What are the physical properties of molecular structures?
A
-Consist of covalently bonded molecules held together by weak Ven der Waals forces.
-Low melting and boiling points- Not much energy required to overcome van der waals forces.
-Poor conductors- no charged particles.
18
Q
Which factors affect the strength of a metallic bond?
A
Charge on the metal ion
Ionic radius
Number of delocalised electrons
19
Q
What factors affect the strength of ionic bonding?
A
Ionic radius
Charges on ions
20
Q
Why are specific bond angles formed?
A
Pairs of electrons in the outer shell arrange themselves as far apart as possible to minimise repulsion. Bonding pairs and lone pairs are clouds of charge that repel each other.
21
Q
Which types of repulsion are the strongest?
A
Lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
22
Q
What is the effect of lone pair repulsion on the bond angle?
A
It reduces it by 2.5 degrees
23
Q
What is electronegativity?
A
The power on an atom to attract the bonding electron pair towards itself within a covalent bond.
24
Q
Which factors affect electronegativity?
A
Size
Nuclear charge
25
What happens to electronegativity across a period?
It increases because the atomic radius decreases, due to increasing nuclear charge and similar shielding.
26
What happens to electronegativity down a group?
It decreases as shielding increases.
27
What is a polar bond?
A bond formed between 2 atoms with different electronegativities. (Electron distribution is unsymmetrical, and a permanent dipole may form.)
28
Why do some molecules with polar bonds not have a dipole?
The charge distribution is symmetrical so, charges cancel out
29
What causes a permanent dipole?
A difference in charge causing a difference in electron density.
30
What is an induced dipole?
Forms when the electron orbitals around a molecule are influenced by another charged particle.
31
What is a quick method to figure out if a bond is polar or not?
If it has different terminal atoms or if the central atom has a lone pair, the molecule is likely to be polar.
32
What determines molecule shape?
The number of electron pair
If they’re bonding or lone pairs.
33
What types of electron pairs and bond angle is present in a V- shaped molecule?
2 bonding pairs
2 lone pairs
104.5
34
What types of electron pairs and bond angle is present in a trigonal planar molecule?
3 bonding pairs
120
35
What types of electron pairs and bond angle is present in a trigonal pyramidal molecule?
3 bonding pairs
1 lone pair
107
36
What types of electron pairs and bond angle is present in a tetrahedral molecule?
4 bonding pairs
109.5
37
What types of electron pairs and bond angles are present in a trigonal bipyramid molecule?
5 bonding pairs
90 and 120
38
What types of electron pairs and bond angle is present in a octahedral molecule?
6 bonding pairs
90
39
What types of electron pairs and bond angle is present in a seesaw molecule?
4 bonding pairs
1 lone pair
90 and 120
40
What types of electron pairs and bond angle is present in a T- shaped molecule?
3 bonding pairs
2 lone pairs
90 and 180
41
What types of electron pairs and bond angle is present in a square planar molecule?
4 bonding pairs
2 lone pairs
90
42
What are the 3 types of intermolecular forces?
Van der Waals
Permanent dipole-dipole
Hydrogen bonds
43
Which properties are influenced by intermolecular forces?
Melting and boiling points
44
What is a van der waals force?
Induced dipole attraction
45
How does atomic radius affect Van der Waals forces?
The larger the molecules, the more electrons and mass it has. So, the forces are stronger.
46
What is a permanent dipole-dipole force?
The force of electrostatic attraction between two polar molecules.
47
How are hydrogen bonds formed?
H is really small and becomes very positive when bonded to F, O or N as they have high electronegativities. The bonds are always linear.
48
Which type of intermolecular force is the strongest?
Hydrogen bond > permanent dipole-dipole > Van der Waals
But Van der Waals can be stronger depending on the size of the molecule.
49
How does shape affect Van der Waals forces?
The more long or straight a molecule is, the closer the two molecules can get increasing the strength of the bond.
50
Why does water expand as it turns into ice?
As liquid water cools to form ice, the molecules make more hydrogen bonds and arrange themselves into a regular lattice structure.
In this regular structure the molecules are further apart on average than the molecules in liquid water - so ice is less dense than liquid water.
51
What is the first ionisation energy?
The energy required to remove one electron from each atom of one mole of gaseous atoms.
52
What is the equation for the first ionisation energy?
X(g) → X+(g) + e-
53
Which factors affect ionisation energy?
-Nuclear charge
-Shielding
-Atomic radius
54
Why does the first ionisation energy generally increase across a period?
-Increasing nuclear charge
-Shielding remains constant
-Therefore there is a stronger attraction of electrons towards the nucleus.
55
Why does aluminium have a lower first ionisation energy than magnesium?
Even though it has more protons,
-The outer electron is shielded by the full 3s orbital
-The 3p orbital is also more further away from the nucleus
-Therefore attraction between the outer electron and the nucleus decreases
56
Why does sulfur have a lower first ionisation energy than phosphorus?
-Sulfur has 4 electrons in the P orbital, so, there's an electron pair creating electron-electron repulsion.
-Phosphorus has no electron pairs and they have the same shielding and atomic radius.
57
Why do first ionisation energies decrease down a group?
-The outer electron will occupy orbitals further away from the nucleus
-There's increased shielding from inner electrons
-The effect of the nuclear charge decreases
58
Why does helium have the highest first ionisation energy of all the elements?
-It has the configuration 1s2 and has no shielding.
-So, the electron experiences a very strong electrostatic force of attraction towards the 2 protons in the nucleus.
59
Why is the second ionisation energy of an atom always greater than the first?
The remaining electrons will experience a greater effect of nuclear charge pulling on each electron.
60
Why does atomic size decrease across a period?
-The atomic radius and shielding remains stable
-The nuclear charge increases so, there's more effective nuclear charge and the electrons are pulled closer.
61
What is ion drift?
Where ions enter a region with no electric field so they just drift through this region.
Lighter ions drift faster as their velocity will be higher whereas heavier ions drift slower as their velocity will be lower. This is because every particles kinetic energy within the mass spectrometer is constant.
62
What does a mass spectrometer tell you?
Relative atomic mass
Relative molecular mass
Relative isotopic abundance
63
What are the 4 different phases of a mass spectrometer?
Electron spray ionisation
Acceleration
Ion drift
Detection
64
Describe elctrospray ionisation.
A sample is dissolved in a volatile liquid
Its forced through a needle connected to a positively charged terminal with a high voltage
Each particle gains a H+ ion
The solvent evaporates
65
Describe electron impact ionisation.
The sample is vaporised
Fired at by high energy electrons
1 electron is knocked off
66
Describe the process of ion acceleration.
They accelerate towards a negatively charged plate as they're attracted to it.
Lighter ions have a higher acceleration
All ions have the same kinetic energy
67
Describe the process of ion drift.
The ions pass through a hole in the plate and form a beam
They stop accelerating as there's no electrical field
Lighter ions drift at a faster velocity
68
Describe the process of ion detection.
Lighter ions arrive at the detector and they gain an electron, so there's a flow of current.
The time taken to reach the detector = mass of the
isotope
The size of current = abundance of isotopes
69
Why is the mass spectrometer kept under a vacuum?
To prevent ions colliding with molecules of air.
70
Why are positive ions formed in the mass spectrometer?
To accelerate them to the detector plate
Ions pass through hole forming a beam
So that they can be detected
71
Why is Phosphorus's sixth ionisation energy much larger than its fifth?
The electron is being removed from the second energy level, which is closer to the nucleus.
72
For bigger molecules, which ionisation technique is used?
Electro spray as bombardment would fragment the molecule.
73
What are the uses of mass spectrometry?
To identify elements
Detecting illegal drugs
Forensic science
Space exploration
74
Explain, in detail, how the relative atomic mass of this element can be calculated from data obtained from
the mass spectrum of an element.
Spectrum gives (relative) abundance (1)
And m/z (1)
Multiply m/z by relative abundance for each isotope (1)
Sum these values (1)
Divide by the sum of the relative abundances (1)
75
State how you would collect hydrogen. State the measurements that you would make in order to calculate the number of moles of hydrogen produced.
Hydrogen collection
Using a gas syringe or measuring cylinder/ graduated vessel over water
Measurements
(i) P
(ii) T
(iii) V
Use ideal gas equation to calculate mol hydrogen or mass/Mr
Mol H2 = mol Mg
76
In terms of structure and bonding explain why the boiling point of bromine is different from that of magnesium. Suggest why magnesium is a liquid over a much greater temperature range compared to bromine.
-Structures
Bromine is (simple) molecular / simple molecules
Magnesium is metallic / consists of (positive) ions in a (sea) of delocalised electrons
-Strength
Br2 has weak (van der Waals) forces between the molecules / weak IMFs
So, more energy is needed to overcome the stronger (metallic) bonds or converse.
-Liquid range
Mg has a much greater liquid range because forces of attraction in liquid / molten metal are strong(er)
77
Name the strongest type of intermolecular force between hydrogen fluoride molecules and draw a diagram to illustrate how two molecules of HF are attracted to each other.
In your diagram show all lone pairs of electrons and any partial charges. Explain the origin of these charges.
Suggest why this strong intermolecular force is not present between HI molecules.
Hydrogen bonding
Draw diagram
Dipole results from electronegativity difference
Fluorine more/very electronegative
HI dipole weaker or bonding e- more equally shared
78
Crystals of sodium chloride and of diamond both have giant structures. Their melting points are 1074 K and
3827 K, respectively. State the type of structure present in each case and explain why the melting point of diamond is so high.
NaCl is an ionic lattice
Diamond is macromolecular
Many covalent C-C bonds need to be broken
79
Describe the bonding in, and the structure of, sodium chloride and ice. In each case draw a diagram showing how each structure can be represented. Explain, by reference to the types of bonding present, why the melting point of these two compounds is very different.
NaCl is ionic cubic lattice
ions placed correctly
Electrostatic attraction between ions
Covalent bonds between atoms in water
Hydrogen bonding between water molecules
Tetrahedral representation showing two covalent and two hydrogen bonds
2 hydrogen bonds per molecule
Attraction between ions in sodium chloride is very strong
Covalent bonds in ice are very strong
Hydrogen bonds between water molecules in ice are much weaker
80
Explain how the concept of bonding and non-bonding electron pairs can be used to predict the shape of, and bond angles in, a molecule of sulfur tetrafluoride, SF4
Illustrate your answer with a diagram of the structure.
4 bonding electron pairs and one lone pair repel as far apart as possible
lone pair - bond pair repulsion > bp-bp
Pushes S-F bonds closer together
shape is trigonal bipyramidal with lone pair
angles <90
and < 120
81
What is relative molecular mass (Mr)?
The mass of one molecule relative to 1/12 the mass of one atom of carbon-12.
82
What is relative atomic mass (Ar)?
The average mass of one of its atoms relative to 1/12 the mass of one atom of carbon-12.
83
What is empirical formula?
The simplest whole number ratio of atoms of each element in a compound.
84
How do you figure out which parts of the equation are a part of the ionic equation?
The parts which change state or are a part of a new molecule.
85
Do isotopes have different physical or chemical properties to the element?
Physical properties as electrons change chemical properties.
86
What is atom economy?
The proportion of the reactant's atoms that end up in the useful products.
87
Why can the yield sometimes be above 100%?
The solid is wet.
88
What is relative formula mass?
It can refer to Mr or Ar. It is the general term.
89
Why do we want a high atom economy?
-Minimises waste of non-renewable reactants.
-Makes as much useful product as possible.
-Reduces pollution from waste products.
90
Why can the yield be less than 100%?
-Side reactions
-The reaction may be reversible.
-The reaction is incomplete.
-Product is lost in transfer losses
-Gas product may be lost.
91
How many particles are there in one mole?
Avagadro's number of particles.
92
How can you figure out the number of hydrogen atoms in NH3?
Using the moles and Avagadro's constant, calculate the number of NH3 molecules and then multiply by 3.
93
What is the concentration of a substance measured in?
moldm-3
94
How do you convert the empirical formula to the molecular formula?
Divide the Mr of the molecular formula by the Mr of the empirical formula to find the ratio.
95
What is molecular formula?
The actual number of atoms of each element in a compound.
96
What is the ideal gas equation?
pV=nRT
97
What is Avogadro's law?
Equal volumes of gases, at the same temperature and pressure, contain the same number of molecules.
98
What are titrations used for?
To work out concentrations of unknown substances.
99
Which indicators are used for an acid- base titration?
Phenolpthalein or methyl orange
100
How can you decrease the uncertainty of the measurement for a titration?
-Decrease the concentration of the solution in the burett
-Increase the size of the aliquat used.
101
How can you improve the overall technique of a making a standard solution?
Wash the solid into the beaker
Wash the wet rod into the beaker
Wash wet beaker into the flask
Wash the filter funnel into the flask
Use a teat pipette to make up to the mark on the volumetric flask
Ensure the bottom of the meniscus is on the graduation mark
Shake the final solution
102
What is an endothermic reaction?
Reactions that absorb energy. ΔH is positive.
103
What is an exothermic reaction?
Reactions that give out energy. ΔH is negative.
104
What is enthalpy change, ΔH?
The heat energy transferred in a reaction at constant pressure. The units of ΔH are kj mol-1.
105
What is standard enthalpy?
Enthalpy under standard conditions:
100kPa
298K
106
What is Hess's law?
The total enthalpy change of a reaction is independent of the route taken.
107
Is breaking and making bonds endothermic or exothermic?
You need energy to break bonds, so bond breaking is endothermic. Stronger bonds take more energy to break.
Energy is released when bonds are formed, so bond making is exothermic.
Stronger bonds release more energy when they form.
108
What is the effect of different environments on bond enthalpies?
The energy needed to break a bond depends on the environment it's in. So you can only find the mean bond enthalpy.
109
What are the sources of inaccuracy in the calorimetry experiment?
Heat loss to surroundings. To reduce effects, use a lid.
Heat loss to equipment
Incomplete reactions
110
Define mean bond enthalpy.
The amount of energy needed to break a covalent bond averaged over different compounds.
111
Define bond enthalpy.
The amount of energy needed to make or break a bond.
112
What is required for reactions to occur?
Collisions between particles taking place with sufficient energy.
113
Define activation energy.
The minimum amount of kinetic energy that particles need to react.
114
Define the term rate of reaction.
The change in concentration of a reactant or product over time.
115
Describe the effect of increasing temperature on the rate of reaction.
A small increase in temperature leads to a large increase in rate of reaction.
The kinetic energy of the particles increases. So, their speed increases and the frequency of collisions increases. More particles collide with the required activation energy so, this also increases the rate.
116
Describe the effect of increasing concentration on the rate of
reaction.
There are more particles in a given volume. So, the distance between particles decreases so, there's more frequent and successful collisions.
117
Describe the effect of increasing pressure on the rate of reaction.
There's the same number of particles in a smaller volume so the distance between the particles decreases and there's more frequent collisions.
118
Define a catalyst.
A substance that increases the rate of a chemical reaction without being changed in chemical composition or amount.
119
How do catalysts work?
They provide an alternative reaction route of lower activation energy.
120
What is the definition of dynamic equilibrium?
The rates of the forward and reverse reactions are the
same.
The net concentrations of the reactants and products of the reaction mixture remain constant.
In a closed system.
121
What is Le Chatelier's principle?
If a factor affecting the position of equilibrium is altered, the position of the equilibrium shifts to oppose the effect of change.
122
What is Le Chatelier's principle used for?
Le Chatelier's principle can be used to predict the effects of changes in temperature, pressure and concentration on the position of equilibrium in homogeneous reactions.
123
What is the effect of increasing the temperature of a reversible reaction?
It favours the endothermic side.
124
What is the effect of increasing the pressure of a reversible reaction?
It favours the side with the least gaseous moles.
125
What is the effect of increasing the concentration of one side of a reversible reaction?
It favours the other side.
126
What happens if you increase the temperature of a reversible reaction that's endothermic in the forward direction? [3 marks]
(Increasing) the (temperature) shifts the equilibrium to the (RHS)
This is because the equilibrium shifts to oppose the (increase) in (temperature)
Increasing the temperature favours the endothermic reaction.
127
What is the effect of adding a catalyst to a reversible reaction?
It has no effect it just increases the speed at which the reaction reaches equilibrium. The rate of both the forwards and the backward reaction are increased by the same amount.
128
What is Kc?
The ratio between the concentrations of reactants and products in a reversible reaction.
It is not affected by changes in concentration or the addition of a catalyst.
It can only be affected by temperature.
129
What is oxidation?
The process of electron loss
130
What is an oxidising agent?
The electron acceptors
131
What is reduction?
The process of electron gain
132
What are reducing agents?
The electron donors
133
What are the rules of oxidation states?
The oxidation number of an element or neutral compound is 0
Overall oxidation number of an ion is equal to the charge of the ion.
134
What is the order of priority for oxidation states?
Group 1: +1
Group 2: +2
Fluorine: -1
Hydrogen: +1
Oxygen: -2
Group 7: -1
Group 6: -2
Group 3: +3
Group 5: -3
135
What is the rate equation?
Rate= k[A]^m[B]^n
Where m and n are the orders of the reaction with respects to the reactants.
136
Which factors affect the value of k, the rate constant?
Temperature, they are directly proportional.
137
What is Arrhenius equation?
k = Ае^-(Ea/RT)
Where A is the Arrhenius constant, this varies between different reactions.
138
What is the rate determining step?
The slowest step in a multistep reaction.
139
What are the continuous methods for monitoring rate?
Calorimetry
Gas volume
Gas mass
140
What are the methods for measuring the initial rate of a fixed concentration change?
Disappearing cross
lodine clock
141
Describe the method for a continuous method for monitoring rates.
Have an excess of other reactants.
Monitor the change in [] of the reactant you're interested in.
Plot a graph for [] vs time
Determine the order with respect to the reactant.
Calculate the rate using the gradient
142
Describe the method for an initial rate method for monitoring rates.
Fix the [] of all the reactants except 1 and double this and determine the initial rate.
Then compare the impact of changing [] to change in rate, to determine order.
Repeat with the other reactants.
OR
Repeat 3 times for different []
Plot rate vs [] to determine order
143
When determining order with respect to a reactant, why do we add other reactants in excess?
So that they don't affect rate.
The concentrations of other reaction remains almost constant (as the change is negligible.)
So, it essentially becomes 0 order with respect to other reactants.
144
What is the definition of the enthalpy of combustion?
The enthalpy change when 1 mole of a substance undergoes complete combustion in oxygen with all substances in standard states.
145
What is the definition of the enthalpy of neutralisation?
The enthalpy change when 1 mole of water is formed in a reaction between an acid and alkali under standard conditions.
146
What is the definition of the enthalpy of atomisation?
The enthalpy change when 1 mole of gaseous atoms is produced from an element in its standard state.
147
What is the definition of first electron affinity?
The enthalpy change when each atom in 1 mole of gaseous atoms gains one electron to form 1 mole of gaseous 1- ions
148
What is the definition of the enthalpy of formation?
The enthalpy change when one mole of a substance is formed from its constituent elements with all substances in their standard states.
149
What is the definition of the lattice enthalpy of formation?
The enthalpy change when 1 mole of a solid ionic compound is formed from its constituent gaseous ions.
150
What is the definition of the lattice enthalpy of dissociation?
The enthalpy change when 1 mole of solid ionic compound is broken up into its constituent ions in the gas phase.
151
What is the definition of the enthalpy of solution?
The enthalpy change when 1 mole of a solid ionic substance dissolves in sufficient water to form an infinitely dilute solution.
(Forms aqueous ions)
152
What is the definition of the enthalpy of hydration?
The enthalpy change when 1mol of gaseous ions are dissolved in water to give 1mol of aqueous ions and a solution of infinite dilution.
153
What is a Brønsted-Lowry acid?
A proton donor
154
What is a Brønsted-Lowry base?
A proton acceptor
155
How do you calculate pH?
pH = -log [H+]
156
What is the Kw expression?
Kw = [H+] [OH-]
157
What affects the value of Kw?
Temperature
158
What is Ka and what is the expression for Ka?
The dissociation constant for a weak acid
Ka = ([H+] [A-]) / ([HA])
159
What is a buffer?
A buffer solution maintains an approximately constant pH, despite dilution or addition of small amounts of acid or base.
160
What is an acidic buffer?
Consists of a weak acid and its salt
161
What is a basic buffer?
Consists of a weak base and its salt
162
How is pH maintained when an acid is added to a buffer.
The buffer is a reservoir of HA and A-
The H+ reacts with the salt
But, the concentration of the salt stays constant
So, pH stays constant
163
How is pH maintained when a base is added to a buffer.
The OH- reacts with the H+
So, equilibrium shifts to oppose the change
[H+] remains constant
So, pH stays the same
164
What are the 2 methods of making a buffer?
A weak acid and its salt
Excess weak acid and strong base
165
What is equivalence point?
When the 2 substances are mixed in equal proportions.
166
How do you select an indicator for a reaction?
The pH range of the indicator / endpoint must fall within the rapid pH change section of the reaction.
167
What is the colour change, endpoint and use of Methyl orange?
Red -> yellow
pH 3.7 - orange
Weak base + strong acid
168
What is the colour change, endpoint and use of phenolphthalein?
Colourless -> pink/ purple
pH 9.3
Strong base + weak acid
169
What is the IUPAC convention for writing half-equations for electrode reactions?
The equation is always written as a reduction reaction.
170
Which electrode is the anode?
The -ve electrode
Oxidation
171
Which electrode is the cathode?
The +ve electrode
Reduction
172
What is the definition of an electrochemical cell?
A cell made from 2 different metals dipped in salt solutions of their own ions and connected by a wire.
173
What is an electrochemical series?
A list of standard electrode potentials in order of emf for different electrochemical half cells.
174
What is the function of the salt bridge?
It allows the movement of ions between electrodes to maintain the electric current.
175
What conditions must the salt bridge meet?
It must not react with the electrolytes.
176
What chemicals are typically used for the salt bridge?
KNO3
KCl (Not used if you have Ag+)
177
What are the standard conditions for the half cells?
298K
100kPa (only when there is a gaseous reactant)
Ion concentration of 1moldm^-3
178
How do you calculate E_cell values?
E_cell = E^0 (reduced) - E^0 (oxidised)
179
What is the conventional representation of cells?
Oxidised || Reduced
The standard hydrogen electrode is always on the left
Pt (s) | H_2 (g) | H+ (aq) ||
If the 2 substances are in the same state, put a comma between them, not a |
180
Write the conventional representation for the standard hydrogen half cell.
Pt (s) | H_2 (g) | H+ (aq) ||
181
How do you know if a specific cell reaction is thermodynamically feasible?
When E_cell > 0
182
How do you find the best oxidising agent in an electrochemical series?
The reduced species with the highest EMF
183
How do you find the best reducing agent in an electrochemical series?
The oxidised species with the lowest EMF
184
Why are there standard conditions for electrochemical cells?
Changing the conditions will affect the E0 value.
185
What are electrochemical cells used for?
A commercial source of electric energy
186
What are the 3 different types of cells?
Non-rechargeable
Rechargeable
Fuel cells
187
Give an example of a rechargeable battery.
Lithium batteries
188
What is the difference between a fuel cell and an electrochemical cells?
A fuel cell is an open system, reactants constantly flow in.
189
What is a fuel cell?
They are used to generate an electric current and do not need to be electrically recharged.
190
What are the advantages of a fuel cell over hydrogen combustion?
Fuel cells are more efficient
Fuel cells are quieter
191
What are the disadvantages of a fuel cell over hydrogen combustion?
Hydrogen is flammable and explosive - So its difficult to transport and store
Most hydrogen comes form methane and crude oil
Can produce hydrogen through electrolysis but energy comes from fossil fuels
192
What are the 2 methods through which hydrogen is produced?
Reactions of hydrocarbons with steam
Electrolysis of acidified water
