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chemistry > topic nine > Flashcards

topic nine Flashcards

1
Q

define oxidation

A

gain of oxygen
loss of hydrogen
loss of electrons
increase in oxidation state

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2
Q

define reduction

A

loss of oxygen
gain of hydrogen
gain of electrons
decrease in oxidation state

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3
Q

two half equations should add together to create

A

the original ionic equation

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4
Q

do oxidising or reducing agents accept electrons

A

oxidising

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5
Q

do oxidising or reducing agents donate electrons

A

reducing

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6
Q

If an atom loses control over electrons, it has been waht

A

oxidised

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7
Q

If an atom gains control over electrons, it has been what

A

reduced

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8
Q

which part comes first in oxidations tate

A

symbol then number

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9
Q

what does oxidation state indicate

A

the gain or loss of electron control of an atom during a reaction

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10
Q

oxidation states of oxygen, chlorine, nitrogen

A

free elements like these are zero

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11
Q

the sum of the oxidation states of all the atoms in a compound mustbe equal to the

A

net charge on the compound

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12
Q

alkali metals oxidation states

A

always +1

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13
Q

fluorine oxidation states

A

always -1

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14
Q

alkaline earth metals oxidation states

A

+2

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15
Q

hydrogen oxidation state

A

always +1 unless with certain metal hydrides where it is -1

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16
Q

oxygen oxidation states

A

-2 except with perioxides where it is -1, or with fluroine where it is +2

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17
Q

ch;loride oxidation state

A

-1 unless with oxygen or fluorine

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18
Q

charge on a metal ion is its

A

oxidation state

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19
Q

average oxidation state

A

the average of the oxidation states of the same atom in a compound.

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20
Q

how are roman numerals used to show the oxidation number

A

Roman numerals are used to show the oxidation number when there is more than one possible oxidation state for an atom in a compound. For example, in copper(I) oxide, the copper has an oxidation number of +1, whereas in copper(II) oxide, it is +2.

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21
Q

when does disproportion occur

A

when the same species is oxidised and reduced simultaneously during a reaction to form two different products

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22
Q

activity series

A
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23
Q

explain the low reactiviy of aluminium

A

the formation of an extremely thin protective layer of aluminium oxide on the surface of the metal prevents the metal underneath from oxidising further.

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24
Q

Any metal above hydrogen will react with an acid to produce

A

a salt and hydrogen gas

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25
Q

any metals below hydrogen will

A

not react with dilute acids

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26
Q

when can metals displace hydrogen ions from solution to produce hydrogen gas

A

when the metals are above hydrogen in the reactivity series

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27
Q

how does difficulty of extraction correlate with reactivity

A

the more reactive metals are much harder to extract

Metals above carbon in the activity series, for example, magnesium and aluminium, cannot be extracted from their ores by reduction with carbon. Conversely, metals below carbon in the series, such as iron and zinc, can be extracted by reduction with carbon. The reaction of zinc oxide with carbon is shown below

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28
Q

how does the reactivity series correlate with reactions with dilute acid

A

the more reactive metals produce H2 with decreasing vigour, and the less reactive dont react with dilute acids

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29
Q

how does the reactivity series correlate with reactions with air/oxygen

A

the more reactive metals burn very brightly and vigorously. burn to form an oxide with decreasing vigour.

the middlely reactive elements react slowly to form the oxide

the low reactive elements dont react

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30
Q

how does the reactivity series correlate with reactions with water

A

produce hydrogen with decreasing vigour with cold water. react with steam with decreasing vigour

less reactive elements dont react with cold water or steam

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31
Q

displacement reaction

A

the more reactive metal displaces the ions of the less reactie metal from solution. they are redox reactions which involve the transfer of electrons as metal atoms are oxidised and aqueous metal ions are reduced

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32
Q

spectator ion

A

an ion appearing on both sides of the equation and does not undergo oxidation or reduction

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33
Q

how does iron change when rusting

A

The precipitate of iron(II) hydroxide formed is rapidly oxidised under basic conditions to form red-brown hydrated iron(III) oxide.

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34
Q

what is the formula of rust

A

Fe2O3.xH2O

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34
Q

why does the rust form flakes that flake off the surface of the iron

A

it has a lower density than iron

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34
Q

sacrifical protection

A

blocks of magnesium or zinc are used to protect iron pipes, ships and storage tanks

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35
Q

galvanisation

A

a layer of zinc is coated over all the iron surface to give complete coverage.

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36
Q

drawback of tin plating

A

it is easily scratched to expose the iron surface. In this case the can rusts very rapidly. Once the iron is brought into contact with moist air, it gives sacrificial protection to the less reactive tin.

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37
Q

winkler method is used to

A

determine the biochemical oxygen demand or bod

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38
Q

what is bod

A

the amount of dissolved oxygen required to biologically decompose the organic matter in a water sample over a set time period (usually five days). Polluted water with a high BOD without the means of replenishing oxygen will not be able to sustain aquatic life. This can lead to the growth of green algae that feed on excessive nutrients in the water, known as eutrophication

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39
Q

when does eutrophication occur

A

when a body of water beocmes enriched in dissolved nutrients that lead to excessive growth of aquatic platn life such as algae

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40
Q

COMPLETE WINKLER METHOD

A
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41
Q

what is a redox titration

A

one that involves a redox reaction between the analyte and titrant

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42
Q

what usually happens in a redox titration

A

A common experiment using a redox titration is to determine the mass of iron in an iron tablet (iron tablets are prescribed to people who have low iron levels in their blood). In this experiment, an iron tablet is first dissolved in an acidic solution and the iron in the tablet is converted to the iron(II) ion, Fe2+. Next, the solution is titrated with an oxidising agent which is usually acidified potassium manganate(VII) or potassium dichromate(VI).

the iron(II) ion is oxidised to the iron(III) ion:
Fe2+ (aq) → Fe3+ (aq) + e–

the manganate(VII) ion, MnO4– is reduced to the manganese(II) ion:
MnO4– (aq) + 8H+ (aq) + 5e– → Mn2+(aq) + 4H2O (l)

The balanced equation in acidic solution is as follows:

5Fe2+ (aq) + MnO4– (aq) + 8H+ (aq) → Mn2+(aq) + 5Fe3+ (aq) + 4H2O (l)

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43
Q

what are voltaic cells aka

A

galvanic cells or batteries

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44
Q

what do voltaic or galvanic cell suse to produced an ee;ctric current

A

a redox reaction

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45
Q

what is a danielll cell

A

an early example of a simple voltaic cell

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46
Q

what does a daniell cell consist of

A

two half-cells: a zinc electrode in a solution of zinc sulfate, and a copper electrode in a solution of copper(II) sulfate. The two electrodes are connected by an external circuit with a voltmeter to measure the voltage. The voltage produced by a voltaic cell depends on the difference in reactivity between the two metals in the half-cells. The circuit is completed by a salt bridge that allows ions to flow in order to complete the circuit and prevent build-up of electric charge. Electrons flow through the wires of the external circuit in a spontaneous process; no external energy source is required.

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47
Q

what does a half cell consist of

A

an electrode, usually a metal, in a solution of its own ions. For example, a zinc half-cell is composed of a piece of solid zinc metal in a solution of aqueous zinc (Zn2+) ions.

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48
Q

how to make a salt bridge

A

a strip of filter paper soaked in saturated potassium nitrate (KNO3), although more complex salt bridges can be made using a glass tube filled with agar gel.

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49
Q

why is potassium nitrate chosen to make a salt bridge with

A

its constituent ions do not react with the other ions present in the two half-cells, or with the electrodes

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50
Q

what happens in the salt beidge

A

anions (negative ions) migrate from the salt bridge to the anode and cations (positive ions) migrate from the salt bridge to the cathode. This migration of ions prevents the build-up of electric charge in the two half-cells.

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51
Q

in a voltaic cell, which mental is oxideised and which is reduced

A

The more reactive metal is oxidised, and the less reactive metal is reduced

52
Q

what creates a potential difference

A

The difference in the tendency to either undergo oxidation or reduction creates a potential difference between the two half-cells, also known as the electromotive force (EMF) or cell potential (E⦵cell).

53
Q

what is cell potential

A

The cell potential (E⦵cell), also known as the electromotive force (EMF), is the difference in the electrode potentials of the half-cells in a voltaic cell. The cell potential of a voltaic cell is measured in volts (V).

54
Q

in a voltaic cell, what happens at the anode and cathode

A

the anode is where oxidation takes place and is negatively charged. The cathode is where reduction takes place and is positively charged.

55
Q

what does teh voltage of a cell in a voltaic cell epend on

A

The voltage of a cell depends on the nature of the electrodes and the ions involved, the temperature and the ion concentrations of the solutions.

56
Q

the more reactive metal forms the

A

anode

57
Q

cell diagram conventiosn have

A

A phase boundary between a solid and an aqueous solution is represented by a single vertical line.
A salt bridge is represented by a double vertical line.
The anode is usually placed on the left and the cathode on the right.
Two types of ions in the same half-cell are separated by a comma.
Spectator ions are not included in the diagram.

58
Q

explain the electrical conductivity of ionic compounds when molten or dissolved in aqueous solutions

A

mobile ions move in a particular direction in an electric field

59
Q

what is electrolysis

A

the process by which a compound is broken down into its constituent elements using electricity.

60
Q

what is an electrolytic cell composed of

A

a molten or aqueous electrolyte
a battery and two electrodes, the anode and cathode.

61
Q

what are electrodes usually made of

A

graphite (carbon) or an unreactive metal such as platinum

62
Q

which terminal is the anode conneced to and which terminal is the cathode connected to

A

The anode is connected to the positive terminal of the battery and the cathode is connected to the negative terminal.

63
Q

in an electrolytic cell, where does oxidation occur

A

at the positive anode

64
Q

in an electrolytic cell, where does reduction occur

A

negative cathode

65
Q

what type of reactiona t a voltaic cell

A

Involves an exothermic spontaneous redox reaction

66
Q

what type of reaction at an electrolyic cell

A

endothermic non-spontaneous redox reaction

67
Q

what does a voltaic cell convert into what

A

Cell converts chemical energy into electrical energy

68
Q

what does an electrolytic cell convert into what

A

Cell converts electrical energy into chemical energy

69
Q

which electrode is pos or neg in the voltaic cell

A

The cathode is positive and the anode is negative during discharge

70
Q

which electrode is pos or neg in the electrolytic cell

A

The cathode is negative and the anode is positive during electrolysis

71
Q

what does a voltaic cell use

A

two separate aqueous solutions connected by a salt bridge and an external circuit

72
Q

what does an electrolytic ccell electrolye consist of

A

The electrolyte is a molten liquid (or an aqueous solution)

73
Q

how is electric current conducted in a voltaic cell

A

Electric current is conducted by the electrons in the external circuit and the movement of ions in the salt bridge

74
Q

how is electric current conducted in an electrolytic cell

A

Electric current is conducted by the electrons in the external circuit and the movement of ions in the electrolyte

75
Q

When the two half-cells are connected together, the electrons flow from… to…

A

from the half-cell with the more negative standard electrode potential (E⦵) value to the half-cell with the more positive standard electrode potential value

76
Q

The voltage generated by a voltaic cell depends on

A

the difference between the values of the standard electrode potentials of the two half-cells.

77
Q

what is the standard electrode potential E⦵

A

the voltage produced when a half-cell is connected to a standard hydrogen electrode under standard conditions.

78
Q

how is standard electrode potential measured

A

by connecting it to a standard hydrogen electrode (SHE)

79
Q

conditions of the SHE:

A

hydrogen gas, H2 (g), at a pressure of 100 kPa and temperature of 298 K
acidic solution with [H+] of 1.0 mol dm–3
inert platinum electrode.

80
Q

A standard hydrogen electrode (SHE).

A
81
Q

The half-equation for the reaction that takes place at the surface of the platinum electrode is at an SHE is

A

2H+ (aq) + 2e– → H2 (g) or H2 (g) → 2H+ (aq) + 2e–

82
Q

When E⦵ > 0, this half-cell will undergo

A

reduction when connected to a SHE. The electron flow will be from the SHE to the half-cell in question, which is the cathode of the cell.

83
Q

When E⦵ < 0, this half-cell will undergo

A

when connected to a SHE. The electron flow will be from the half-cell to the SHE. This half-cell will be the anode of the cell.

84
Q

which metals in the electrochemical series are stronger reducing agents

A

metals in the half cells at the top of the series with large negative electrode potential values have the greatest tendency to lose electrons and undergo oxidation and from positive ions in aqueous solution

85
Q

what causes the electromotive force (EMF) or cell potential (E⦵cell)

A

The difference in the tendency to either undergo oxidation or reduction creates a potential difference between the two half-cells, also known as the electromotive force (EMF) or cell potential (E⦵cell).

86
Q

SHE: When two half-cells are connected, the half-cell with the more negative standard electrode potential value undergoes

A

oxidation (is the anode of the voltaic cell)

87
Q

When two half-cells are connected, he half-cell with the more positive standard electrode potential value undergoes

A

reduction (is the cathode of the voltaic cell)

88
Q

what is the equation used to measure E⦵cell=

A

E⦵cell = E⦵half-cell where reduction occurs − E⦵half-cell where oxidation occurs

89
Q

ΔG⦵ =

A

– nFE⦵

90
Q

when E⦵ is positive, ΔG⦵ is negative and the reaction is

A

spontaneous

91
Q

when E⦵ is negative, ΔG⦵ is positive and the reaction is

A

non spontaneous

92
Q

when E⦵ = 0, ΔG⦵ = 0, the reaction is

A

at equilibrium

93
Q

what is electrolysis

A

a process in which an electric current is used to break apart a compound into its constituent elements

94
Q

Electrolysis involves a non-spontaneous reaction with what ecell values and what gvalues

A

negative E⦵cell value and a positive ΔG⦵ value.

95
Q

where can electrolysis take palce

A

in eithe rmolten liquids or aqueous solutions of ioni compou ds

96
Q

why is the electrolysis of aqueous solutions of ionic compounds is more complicated than that of molten ionic compounds

A

because the water itself can be electrolysed.

97
Q

equation for how water can be pxidised at the anode

A

Oxidation: 2H2O (l) → O2 (g) + 4H+ (aq) + 4e–

98
Q

equation for how water can be reduced at the cathode

A

Reduction: 2H2O (l) + 2e– → H2 (g) + 2OH– (aq)

99
Q

The process of selective discharge occurs when

A

one species is preferentially discharged over another at either the anode or cathode of an electrolytic cell.

100
Q

When there is more than one possible species that could be discharged at the electrodes, the choice of species that is actually discharged depends on three factors:

A

The standard electrode potential (E⦵) value of the half-equations for the reaction.
The relative concentrations of the solutions (such as for the halogens in aqueous solution).
The material that the electrodes are composed of (such as the use of graphite or copper electrodes in the electrolysis of copper(II) sulfate).

101
Q

at the anode, with a dilute ocnc of sodium chloride, what is oxidised

A

H2O (l) is oxidised to form oxygen gas:

2H2O (l) → O2 (g) + 4H+ (aq) + 4e–

102
Q

at the anode, with a saturated conc of sodium chloride, what is oxidised

A

chloride ions are oxidised to form chlorine gas:

2Cl– (aq) → Cl2 (g) + 2e–

103
Q

at the cathode, with sodium chloride solution, waht are the two possible reactions

A

The two possible reactions at the cathode are the reduction of sodium ions or the reduction of water:

Na+ (aq) + e– → Na (s) E⦵ = –2.71 V

2H2O (l) + 2e– → H2 (g) + 2OH– (aq) E⦵ = –0.83 V

H2O is a stronger oxidising agent than Na+ ions (H2O has a more positive E⦵ value), meaning that it is more easily reduced. Therefore, at the cathode, water undergoes reduction to form hydrogen gas.

104
Q

The electrolysis of brine (NaCl) produces

A

hydrogen and chlorine gases, but the product of greatest economic importance is the alkali, sodium hydroxide.

105
Q

where is sodium hydroxide produced in sodium chloride electrolysis

A

Note that sodium hydroxide is not produced at either the cathode or the anode of the cell, instead, it exists as the solution formed in the reaction.

106
Q

why is pure water a poor conductor of electricity

A

it lacks any of the mobile ions in solution

107
Q

what do you add to pure water to allow for its electrolyisis

A

To increase its conductivity and allow the solution to conduct electricity, a strong electrolyte such as sulfuric acid is added, effectively creating a dilute sulfuric acid solution. When very dilute sulfuric acid is electrolysed, one volume of oxygen gas is collected at the anode and two volumes of hydrogen gas are collected at the cathode

108
Q

electrolyis of dilute sulfuric acid at the anode:

A

At the anode, water is oxidised to produce oxygen gas
Anode: 2H2O (l) → O2 (g) + 4H+ (aq) + 4e–

109
Q

electrolyiss of dilute sulfuric acid at the cathode

A

at the cathode, water is reduced to produce hydrogen gas:
Cathode: 4H2O (l) + 4e– → 2H2 (g) + 4OH– (aq)

110
Q

The overall equation for the electrolysis of water is:

A

2H2O (l) → 2H2 (g) + O2 (g)

111
Q

describe the changes in pH for the electrolyis of waater

A

The changes in pH that occur at the electrodes are as follows: at the anode, the pH decreases as H+ ions are produced; at the cathode, the pH increases as OH– ions are produced.

112
Q

The electrolysis of copper(II) sulfate produces

A

pure copper at the cathode.

113
Q

At the anode of the electrolysis of copper(II) sulfate

A

water is oxidised to form oxygen gas:

2H2O (l) → O2 (g) + 4H+ (aq) + 4e–

114
Q

electrolysis of copper sulfate, If the graphite electrodes are replaced by copper electrodes, the reaction taking place at the cathode remains the same – copper(II) ions are reduced – but at the anode the copper electrode itself undergoes oxidation:
(give the equation)

A

Cu (s) → Cu2+ (aq) + 2e−

115
Q

why is the electrolysis of copper sulfate used

A

in the refinement of impure copper

116
Q

overall equation for the electrolyiss of copper sulfate

A

Cu (s) + Cu2+ (aq) → Cu2+ (aq) + Cu (s)

117
Q

what does electrplating invovle

A

the coating of a piece of metal with a thin layer of another metal. Metal objects are electroplated to prevent corrosion or to improve their appearance. Copper, chromium, silver and tin are the most commonly used metals for electroplating. Examples of electroplated objects include jewellery and cutlery

118
Q

silver electroplating at the cathode

A

silver ions undergo reduction to form silver atoms which coat the spoon:

Ag+ (aq) + e− → Ag (s)

119
Q

silver electroplating at the anode

A

the silver atoms that make up the piece of silver undergo oxidation:

Ag (s) → Ag+ (aq) + e−

120
Q

The object to be electroplated is

A

the cathode.

121
Q

The anode is composed of the metal to be

A

plated onto the obect

122
Q

electroplating: The electrolyte solution is a salt containing the ions of

A

the metal to be plated onto the object.

123
Q

Q =

A

It

where: Q = charge (in coulombs)

I = current (in amps)

t = time (in seconds)

124
Q

One mole of electrons has a charge of

A

96 500 coulombs.

125
Q

From the equation Q = It, we can see that charge (Q) is directly proportional to the current (I) and the time (t) for which the current is applied. For example:

if the current is doubled,

A

the charge is doubled

126
Q

From the equation Q = It, we can see that charge (Q) is directly proportional to the current (I) and the time (t) for which the current is applied. For example:

if the time is doubled

A

the charge is doubled

127
Q

From the equation Q = It, we can see that charge (Q) is directly proportional to the current (I) and the time (t) for which the current is applied. For example:

and if both the curretn and time are doubled,

A

the charge will be four times greater

128
Q
A

chemistry (13 decks)

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