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Chemistry > Elements from the sea > Flashcards

Elements from the sea Flashcards

1
Q

What is electrolysis ?

A

If you pass an electric current through an ionic substance that molten or in solution, it breaks down a substance into the elements its made of.

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2
Q

What is the electrolyte ?

A

A liquid to conduct electricity.

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3
Q

Do electrolytes contain free ions ?

A

Yes, the ions are usually molten or dissolved ionic substances.

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4
Q

Do free electrons conduct electricity ?

A

Yes

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5
Q

What are anions ?

A

Negative ions that move to the positive electrode (anode) and lose electrons.

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6
Q

What are cations ?

A

Positive ions that move to the negative electrode (cathode) and gain electrons.

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7
Q

How do you carry out electrolysis in an aqueous solution ?

A

1) Use wires and clips to connect each electrode to the power supply. The electrode connected to the positive pole will be the anode, and the electrode that is connected to the negative pole will be the cathode.
2) You usually inert electrodes (such as platinum or carbon electrodes) so that they dont start reacting and interfering with the electrolysis.
3) Place the electrodes into a beaker containing the electrode, making sure that the electrodes aren’t touching each other.
4) Turn the power supply on
5) Depending on what electrodes you’re using, the products will form as metals (as a thin layer on the surface of the cathode - known as plating), or as gases (as bubbles at the cathode or the anode).

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8
Q

What do half-equation shows ?

A

The movement of electrons during a reaction.

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9
Q

What will the half equation at the anode show ?

A

Negative ions losing electrons to form atoms

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10
Q

What will the half equation at the cathode show ?

A

Positive ions gaining electrons to form atoms

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11
Q

What are the half equations for the electrolysis of molten zinc chloride ?

A

Anode : 2Cl-(aq) –> Cl2(g) + 2e-
Cathode : Zn2+(l) + 2e- –> Zn(s)

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12
Q

Electrolysis is carried out using lead bromide. Write half-equations to show what will form at each electrode.

A
  • Lead ions are positive so will move to the cathode, to form lead.
    Pb2+(l) + 2e- –> Pb(s)
  • Bromide ions are negative so will move to the anode, to form bromine gas.
    2Br-(l) –> Br2(g) + 2e-
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13
Q

In aqueous solutions what do you have because of the water ?

A

H+ and OH- ions from the water as well as the ions from the ionic compound.

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14
Q

What do the products of the electrode depend on ?

A
  • The reactivity of the ions
  • Concentration of the salt solution
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15
Q

What is the rule at the cathode ?

A
  • If the metal is less reactive than hydrogen (e.g. silver or copper), then the metal will be formed.
  • If the metal is more reactive than hydrogen (e.g. all group 1 and 2 metals and aluminium) hydrogen gas will be formed (from hydrogen in the water)
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16
Q

What is the rule a the anode ?

A
  • If the solution doesn’t contain a halide, oxygen will be formed (from hydroxide ions in the water). Here is the half-equation for this reaction:
    4OH-(aq) –> O2(g) +2H2O(l) + 4e-
  • If the solution is concentrated and contains a halide then halogen will be formed.
  • If solution contains a halide but is dilute, oxygen will be formed (from hydroxide ions in the water).
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17
Q

What metals are less reactive than
hydrogen ?

A

Gold
Silver
Mercury
Copper
Lead
Tin
Ferum

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18
Q

What metals are more reactive than
hydrogen ?

A

All of group 1 and 2 metals, and aluminium.

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19
Q

In the purification of copper, the anode is made from impure copper and the cathode is made from pure copper. How are copper ions made ?

A

At the anode, copper atoms lose electrons and become copper ions, which enter the solution. These ions are then attracted to the cathode where they gain electrons to become copper atoms again and plate the pure copper cathode:
Anode: Cu(s) –> Cu2+(aq) + 2e-
Cathode: Cu2+ + 2e- –> Cu(s)

The pure cathode increases in mass.
The impure anode shrinks.

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20
Q

Why does an ionic substance need to be molten or in solution before you can electrolyse it ?

A

To conduct electricity it requires a liquid, called an electrolyte. The electrolyte contains free ions, these ions are usually the molten or dissolved ionic substance. The ions conduct the electricity.

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21
Q

Describe how you would set up the equipment for the electrolysis of a solution of aqueous copper sulfate ?

A
  • Use wires and clips to connect each electrode to the power supply. The electrode that is connected to the positive pole is the anode, and the electrode that is connected to the negative pole is the cathode.
  • You usually use the inert electrodes so that they don’t start reacting and interfering with the electrolysis.
  • Place the electrodes into a beaker containing the electrolyte, making sure that the electrodes aren’t touching each other.
  • Turn the power supply on
  • Depending on what electrode your using the products will form as metals, or gases.
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22
Q

Explain what determines the product formed at the cathode when you electrolyse an aqueous solution ?

A
  • The reactivity of the ions
  • The concentration of the salt solution.
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23
Q

What is formed at the anode when a concentrated halide solution is
electrolysed ?

A

A halogen will be formed

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24
Q

Two carbon electrodes are placed into molten copper chloride and the power supply is turned on.

What would you see happening at the anode ? Explain your answer.

A

At the anode, you would see bubbles of gas appearing because chloride ions are losing electrons to form chlorine.

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25
Q

Two carbon electrodes are placed into molten copper chloride and the power supply is turned on.

What would you see happening at the cathode ? Explain your answer.

A

At the cathode, you would see the electrode being plated because copper ions are gaining electrons to form copper metal.

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26
Q

A student carries out an electrolysis of a dilute solution of magnesium chloride, MgCl2(aq). Write a half-equation for the reaction that occurs at:

1) the cathode

2) The anode

A

1) 2H+ (aq) + 2e- –> H2(g)

2) 4OH-(aq) –> O2(g) + H2O (l) + 4e-

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27
Q

What can be extracted by the electrolysis of halide solutions ?

A

Most halogens

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28
Q

When you electrolyse concentrated aqueous solutions, which contain halide ions, what is produced at the anode ?

A

The halogen is released

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29
Q

How do you form molecules in electrolysis of halide solutions ?

A

The halide ions lose electrons to the electrode and are oxidised to atoms, which combine to form molecules.

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30
Q

Write half-equations to show what will form at each electrode when a concentrated solution of sodium bromide is electrolysed ?

A

Bromine forms at the anode.
Anode: Br-(aq) –> 2e- + Br2(aq)

Hydrogen ions - from the water - form hydrogen gas at the cathode.
Cathode: H+(aq) + 2e- –> H2(g)

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31
Q

What can be extracted from brine in electrolysis ?

A

Chlorine

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32
Q

What is brine ?

A

A solution of water with a high concentration of salts - mainly sodium chloride, but also bromine and iodine salts.

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33
Q

Where can you find brine ?

A
  • Naturally in salt lakes
  • Seawater
  • Dissolving rock salt in water
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34
Q

Industrially, chlorine is made from the electrolysis of brine. Explain the process of the electrolysis of brine ?

A
  • At the cathode. two hydrogen ions accept two electrons to become one hydrogen molecule:
    2H+(aq) + 2e- –> H2(g)
  • At the anode, two chlorine (Cl-) ions lose their electrons and become one chlorine molecule:
    2Cl-(aq) –> Cl2(g) + 2e-
  • The sodium ions stay in solution because there more reactive than hydrogen.
  • Sodium ions and hydroxide ions (from water) are left behind while hydrogen and chlorine are removed, o sodium hydroxide (NaOH) is left in the solution.
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35
Q

Are the electrodes inert in the electrolysis of brine ?

A

Yes

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36
Q

Is chlorine collected as a gas ?

A

Yes

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37
Q

Is the sodium chloride solution (used in the electrolysis for brine) concentrated or
dilute ?

A

Concentrated

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38
Q

Why can’t you use a dilute sodium chloride solution, when doing the electrolysis of brine ?

A

The chloride (Cl-) ions aren’t discharged - they hang on to their electrons. The OH- ions lose electrons instead and the products at the anode at oxygen and water, not chlorine.

4OH-(aq) –> 4e- + 2H2O(l) + O2(g)

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39
Q

Is bromine and iodine displaced by chlorine in a displacement reaction, in the electrolysis of brine ?

A

Yes

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40
Q

How is bromine formed in the electrolysis reaction of brine ?

A
  • Brine contains bromide ions
  • Chlorine is more reactive than bromine. So when you bubble chlorine has though brine, the chlorine will displace the bromine.
  • Here’s the equation for this reaction:
    2Br-(aq) + Cl2(g) –> Br2(g) + 2Cl-(aq)
  • The bromine that is produced is then collected, condensed into a liquid and purified.
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41
Q

Describe how iodine is displaced by chlorine in the electrolysis reaction of brine ?

A
  • Brine contains iodine ions
  • Chlorine is more reactive than iodine. So, when you bubble chlorine gas through brine, the chlorine will displace the iodine.
  • Here is the equation for this reaction:
    2I-(aq) + Cl2(aq) –> I2(aq) + 2Cl-(aq)
  • The iodine is collected, purified and condensed into a grey solid.
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42
Q

What is brine ?

A

A solution of water with a high concentration of salt - mainly sodium chloride, but also some bromine and iodine.

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43
Q

Which electrode does chlorine form at during the electrolysis of brine ?

A

At the anode
2Cl-(aq) –> Cl2(g) + 2e-

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44
Q

Describe how bromine is produced from brine ?

A
  • Brine contains some bromine ions
  • Chlorine is more reactive than bromine. So, when you bubble chlorine gas through brine, the chlorine will displace the bromine.
  • Here is the equation for this reaction:
    2Br-(aq) + Cl2(g) –> Br2(g) + 2Cl-(aq)
  • The bromine that is produced is then collected, condensed into a liquid and purified.
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45
Q

Chlorine gas is produced industrially from the electrolysis of sodium chloride solution.

Name one natural source of sodium chloride

A

Sea water
Rock salt
Brine

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46
Q

Chlorine gas is produced industrially from the electrolysis of sodium chloride solution.

A student carries out the electrolysis of concentrated sodium chloride. Once the electrolysis is complete, they add a few drops of universal indicator to the remaining solution. What would you expect the student to see ? Explain your answer.

A

The universal indicator would turn the solution blue/purple. The solution would be alkaline, because of the hydroxide ions left in the solution.

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47
Q

Chlorine gas is produced industrially from the electrolysis of sodium chloride solution.

Another student carries out electrolysis of dilute sodium chloride. They also add a few drops of universal indicator to the remaining solution after electrolysis is complete. How would their observations differ from those of the student before ? Explain your answer.

A

The universal indicator would turn the solution green because hydrogen ions and hydroxide ions have been discharged, leaving a neutral solution of NaCl.

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48
Q

Chlorine gas is produced industrially from the electrolysis of sodium chloride solution.

Chlorine gas can be used to displace iodine from brine. State why chlorine displaces iodine from a solution of iodine ions.

A

Chlorine is more reactive than iodine.

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49
Q

What do oxidation states show ?

A

How many electrons it has donated or accepted to form an ion or a bond.

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50
Q

What are the rules of oxidation states ?

A

1) All uncombined elements have an oxidation state of 0. Elements that are bonded to identical atoms also have an oxidation state of 0.

2) The oxidation state of a simple, monatomic (that’s an ion consisting of just one atom) is the same as its charge.

3) For molecular ions, the overall oxidation state is the same as the overall charge on the ion. Each of the individual atoms in the compound ion will have an oxidation state of it’s own, and the sum of their oxidation states will equal the overall oxidation state.

4) For a neutral compound, the overall oxidation state is 0.

5) The common states of oxygen and hydrogen
Oxygen: -2
HYdrogen: +1

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51
Q

What is the exception, when the oxidation state of oxygen is not -2 ?

A

Peroxides = -1
Flourides
O2

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52
Q

What tells you the oxidation state ?

A

Roman Numerals.

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53
Q

What are two examples where you can tell the oxidation state by the roman numerals ?

A

In iron (II) sulfate = oxidation no. = +2
In iron (III) sulfate = oxidation no. = +3

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54
Q

Name the ion;
- Nitrate or nitrate(V)
- Sulfate or sulfate(VI)
- Carbonate
- Manganese(VII)
- Hydroxide
- Ammonium
- Hydrogencarbonate
- Sulfide

A
  • NO3-
  • SO42-
  • CO32-
  • MnO4-
  • OH-
  • NH4+
  • HCO3-
  • S2-
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55
Q

When naming an ionic compound do you put the positive ion in front of the negative ion ? Example.

A

Put the positive ion in front of the negative ion
H2S = Hydrogen sulfide
LiOH = lithium hydroxide

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56
Q

If the ion ends in an -ate, what does the ion contain ? Example.

A

Oxygen
Another element
Example - sulfate = sulfur + oxygen
nitrate = nitrogen + oxygen

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57
Q

What is the oxidation state of sulfur in the molecule SO2-3

A

+4

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58
Q

What is the oxidation state of sulfur in the molecule So2-4

A

+6

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59
Q

What is the formula of the compound iron(III) sulfate(IV) ?

A
  • From the name you can tell that the iron has an oxidation state of +3.
  • The formula of the sulfate(VI) ion is SO2-4 - this ion has an overall oxidation state of -2.
  • The compound is neutral, so you need to find a ratio of Fe3+ : SO2-4 that will give the overall compound an oxidation state of 0.
    (+3 x 2) + (-2 x 3) = 6 + -6 = 0
  • So the formula is Fe2(SO4)3.
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60
Q

What is the name of the compound ion ClO2- ?

A
  • ClO2- contains chlorine and oxygen, so it is a chlorate.
  • The overall oxidation state of this compound is -1.
  • Oxygen has an oxidation state of -2. The total charge from the 2 oxygen atoms is
    -2 x 2 = -4.
  • The chlorine atom must have an oxidation state of +3 since -4 + 3 = -1.
    So the systematic name of the ion is chlorate(III).
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61
Q

What is the usual oxidation state of oxygen when its combined with another element ?

A

-2

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62
Q

What is the oxidation state of hydrogen in H2 ?

A

0

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63
Q

Work out the oxidation state of nitrogen in the ammonium ion ?

A

NH4+
The overall charge = +1
X + (+1 x 4) = +1
X + 4 = +1
X = -3

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64
Q

Lead(II) Sulfate can be made by reacting Lead(II) Oxide with sulfuric acid. State the oxidation state of:

a) Lead in Lead(II) Oxide

B) Sulfur in Lead(II) Sulfate(VI)

A

a) +2 (The roman numerals after the lead ion tells you)

b) +6

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65
Q

Chromium (Cr) can exist in a range of oxidation states. Give the full name of the compound Cr2O3 ?

A
  • Oxygen has an oxidation state of -2.
  • Total charge from the oxygen = -2 x 3 = -6
  • For the overall oxidation state of the compound to be O, the total charge from the chromium must be +6. +6 / 2 = +3, so the chromium has an oxidation charge of +3.
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66
Q

What is the formula of iron(II) nitrate(V)?

A

Oxidation state of iron = +2
Overall state of nitrate ion = +5 + (-2 x 3) = -1
So you need the ration of 1:2 of iron:nitrate to make the compound.
Fe(NO3)2.

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67
Q

What does oxidation mean ?

A

A loss of electrons

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68
Q

What does reduction mean ?

A

A gain of electrons

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69
Q

When will the oxidation state increase by 1 ?

A

When an electron is lost

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70
Q

When will the oxidation state decrease by
1 ?

A

When an electron is gained.

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71
Q

What do half equations show ?

A

Whats been reduced and whats been oxidised

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72
Q

Magnesium burns in oxygen to form magnesium oxide (MgO). What is the half
equation ?

A

Magnesium is oxidised:
Mg –> Mg2+ + 2e-

Oxygen is reduced:
0.5O2 + 2e- –> O2-

Combine the half equations:
Mg + 0.5O2 –> MgO

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73
Q

What do dilute ions contain ?

A

Hydrogen ions, H+ in solution

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74
Q

What will be produced when magnesium metal reacts with dilute acid ?

A

Hydrogen gas

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75
Q

In the reaction between magnesium metal and dilute hydrochloric acid, what is the reducing agent and oxidising agent ?

A
  • Dilute acids contain hydrogen ions, H+, in solution.
  • Magnesium metal will react with dilute HCl to produce hydrogen gas.
  • The metal ions are oxidised, losing electrons and forming soluble magnesium ions. The hydrogen ions in solution are reduced, gaining electrons and forming hydrogen molecules.
  • Mg(s) + 2HCl(aq) –> MgCl2(aq) + H2(g)
  • Magnesium is oxidised:
    Mg(s) –> Mg2+(aq) + 2e-
  • Hydrogen ions are reduced:
    2H+(aq) + 2e- –> H2(g)
  • In this reaction, the reducing agent is magnesium > it has reduced hydrogen and has been oxidised itself.
  • Hydrogen is the oxidising agent > it has oxidised magnesium while being reduced itself.
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76
Q

Does the reducing agent accept or donate electrons ?

A

Donates electrons

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77
Q

Does the oxidising agent accept or donate electrons ?

A

Accepts electrons

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78
Q

What does balancing redox reactions require ?

A

Balancing the charges

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79
Q

Balance this redox equation:
Au3+ + I- –> Au + I2

A
  • First balance the atoms on each side of the equation: Au3+ + 2I- –> Au + I2
  • Now, check to see if the charges are balanced. The left hand side has a charge of (3+ + (2 x -1)) = 1, and the right-hand side has a charge of 0. So the charges aren’t balanced.
  • Find the change in oxidation state of both elements in the equation:
    Au > +3(Initial Ox.no) > 0(Final Ox.no) > -3
    I > -1 (Initial Ox. no) > 0(Final OX. no) > +2
  • To balance the charges, all you need to do is to change in oxidation states for the two elements. To do this you have to work out what numbers you would need to multiply the top and bottom rows of the table by, so that the changes in the oxidation state will cancel each other out.
  • In this case it x 2 for gold and x3 for iodine (since -3 x 2 = -6 and +2 x 3 = +6)
  • Now all there is left to do is put these numbers into the equation. To do this, you need to: multiply all the Au species in the equation by 2 and all the I species by 3:
    (2 x Au3+) + (3 x 2I-) –> (2 x Au) + (3 x I2)
  • So the balanced equation is:
    2Au3+ + 6I- –> 2Au + 3I2
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79
Q

Balance this redox equation:
Au3+ + I- –> Au + I2

A
  • First balance the atoms on each side of the equation: Au3+ + 2I- –> Au + I2
  • Now, check to see if the charges are balanced. The left hand side has a charge of (3+ + (2 x -1)) = 1, and the right-hand side has a charge of 0. So the charges aren’t balanced.
  • Find the change in oxidation state of both elements in the equation:
    Au > +3(Initial Ox.no) > 0(Final Ox.no) > -3
    I > -1 (Initial Ox. no) > 0(Final OX. no) > +2
  • To balance the charges, all you need to do is to change in oxidation states for the two elements. To do this you have to work out what numbers you would need to multiply the top and bottom rows of the table by, so that the changes in the oxidation state will cancel each other out.
  • In this case it x 2 for gold and x3 for iodine (since -3 x 2 = -6 and +2 x 3 = +6)
  • Now all there is left to do is put these numbers into the equation. To do this, you need to: multiply all the Au species in the equation by 2 and all the I species by 3:
    (2 x Au3+) + (3 x 2I-) –> (2 x Au) + (3 x I2)
  • So the balanced equation is:
    2Au3+ + 6I- –> 2Au + 3I2
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80
Q

What is oxidation ?

A

A loss of electrons

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81
Q

What happens to the oxidation state of a substance when its reduced ?

A

Increase by 1

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82
Q

Complete the sentence: during the redox reaction, an oxidising agent … electrons and is …. .

A

During the redox reaction, an oxidising agent accepts electrons and is reduced.

83
Q

What is a reducing agent ?

A

Donates electrons

84
Q

The equation below shows the reaction of chlorine has with Iron(II) in solution:
Cl2(g) + 2Fe2+(aq) –> 2Cl-(aq) + 2Fe3+(aq)

a) Write a half equation to show the conversion of chlorine, Cl2, into chloride ions, Cl- ?

b) Write half equation to shows the conversion if Iron(II) ions, Fe2+, into iron(III) ions, Fe3+ ?

c) In his reaction, which species is the reducing agent ? Explain why.

A

a) Cl2 + 2e- –> 2Cl-

b) Fe2+ –> e- + Fe3+

c) The reducing agent is Iron(II), Fe2+, because in the reaction they are oxidised/donate electrons.

85
Q

Use oxidation states to balance this redox equation ?
Ca(s) + Al3+(aq) –> Ca2+(aq) + Al(s)

A

Ca > 0(Initial Ox. no) > +2(Final Ox.no) > -2
Al > +3(Initial Ox. no) > 0(Final Ox. no) > +3

(So you need to multiply all the Ca species by 3 and all the Al species by 2)

3Ca(s) + 2Al3+(aq) –> 3Ca2+(aq) + 2Al(s)

86
Q

To find out the concentration of the oxidising agent, what process do you have to do ?

A

Iodine - sodium thiosulfate titration

87
Q

If there is a concentrated oxidising agent what does this suggest ?

A

The more ions will be oxidised by a certain volume of it.

88
Q

How do you find out the concentration of the oxidising agent, potassium iodate(V) (KIO3) ?

A

STAGE 1 - Use a sample of oxidizing agent to oxidise as much iodide as possible.
- Measure out a certain volume of potassium iodate(V), the oxidising agent > 25cm^3.
- Add this to excess of acidic potassium iodide solution. The iodate(V) ions in the potassium iodate(V) solution will oxidise some of the iodide ions to iodine.
IO3-(aq) + 5I-(aq) + 6H+(aq) –> 3I2(aq) + 3H2O(l)

STAGE 2 - Find out how many moles of iodine have been produced.
- You do this by titrating the resulting solution with sodium thiosulfate.
- The iodine in the solution reacts with thiosulfate ions like this:
I2(aq) + 2SO2O32-(aq) –> 2I-(aq) + S4O62-(aq)

STAGE 3 - Calculate the concentration of the oxidising agent.
- Now you look at your original equation:
IO3-(aq) + 5I-(aq) +6H+(aq) –> 3I-(aq) + 3H2O(l)
- The equation shows that one mole of iodate(V) ions produces three moles of iodine. 25cm3 of potassium iodate(V) solution produced 6.66 x 10^-4 moles of iodine. So there must have been
(6.66 x 10^-4 / 3) = 2.22 x 10^-4 moles of iodate(V) ions. This is the same as the number of moles of potassium iodate(V) in the solution. SO now it’s straightforward to find the concentration of the potassium iodate(V) solution, which is what you’re after.

Moles = conc. x vol.
2.22 x 10^-4 = (conc. x 25.0cm3) / 1000
Conc. = (2.22 x 10^-4 x 1000) / 25.0
Conc. = 8.88 x 10^-3 mol/dm

89
Q

The iodine in the solution produced in stage 1 reacted fully with 11.1 cm3 of 0.120 mol/dm thiosulfate solution, what are the moles of iodine ?
I2(aq) + 2S2O32-(aq) –> 2I-(aq) + S4O62-(aq)

A

No. of moles of thiosulfate
(conc. x vol.) / 1000
(0.120 x 11.1) / 1000
= 1.332 x 10^-3 moles
1 mole of iodine reacts with 2 moles of thiosulfate, so number of moles of iodine in the solution
(1.332 x 10^-3) / 2 = 6.66 x 10^-4 moles

90
Q

What don’t you do in a titrate that would make the results inaccurate ?

A
  • Using contaminated apparatus could make your results inaccurate > so make sure the burette is very clean, and rinse it out with sodium thiosulfate before you start.
  • It’s important to read the burette correctly
  • To reduce the effect of random errors, you should repeat the experiment until you get at least three results within 0.1cm3 of each other. Then take an average.
  • Wash the flask between repeat experiments or use a new, clean one.
91
Q

What don’t you do in a iodine-thiosulfate reaction that would make the results inaccurate ?

A
  • The solutions you’re using will react with oxygen in the air, so they should be made up as freshly possible.
  • If you add the starch solution too soon during the titration, the iodine will ‘stick’ to the starch and won’t react as expected with the thiosulfate, making the result unreliable. Only add the starch when solution is pale yellow.
92
Q

What is added during an iodine-thiosulfate titration to make the end point easier to
see ?

A

Add starch solution

93
Q

Describe the colour change at the end point of the iodine-sodium thiosulfate titration ?

A

When the iodine colour fades to a pale yellow, add starch > the colour will turn to a dark blue.

94
Q

10 cm3 of potassium iodate(V) solution was reacted with excess acidified potassium iodide solution. All of the resulting solution was titrated with 0.16 mol/dm3 sodium thiosulfate solution. It fully reacted with 20cm3 of the sodium thiosulfate solution.

Balance the following ionic equation, showing how iodine is formed in the reaction between iodate(V) ions and iodide ions in acidic solution:
IO-3(aq) + I-(aq) + H+(aq) –> I2(aq) + H2O(l)

A

IO-3(aq) + 5I-(aq) + 6H+(aq) –>
3I2(aq) + 3H2O(l)

95
Q

10 cm3 of potassium iodate(V) solution was reacted with excess acidified potassium iodide solution. All of the resulting solution was titrated with 0.16 mol/dm3 sodium thiosulfate solution. It fully reacted with 20cm3 of the sodium thiosulfate solution.

How many moles of thiosulfate ions were there in 20cm3 of the sodium thiosulfate solution ?

A

IO-3(aq) + 5I-(aq) + 6H+(aq) –>
3I2(aq) + 3H2O(l)

Moles = Conc. x Vol.

(0.16 x 20) / 1000

0.0032 moles of thiosulfate ions

96
Q

10 cm3 of potassium iodate(V) solution was reacted with excess acidified potassium iodide solution. All of the resulting solution was titrated with 0.16 mol/dm3 sodium thiosulfate solution. It fully reacted with 20cm3 of the sodium thiosulfate solution.

In the titration, iodine reacted with sodium thiosulfate according to this equation:

I2(aq) + 2Na2S2O3(aq) –>
2NaI(aq) + Na2S4O6(aq)

Calculate the number of moles of iodine that reacted with the sodium thiosulfate
solution ?

A

Moles = conc x vol
Moles of sodium thiosulfate
(0.16 x 20) / 1000 = 0.0032

Mole ratio I2 : Na2S2O3
1 : 2
0.0016 : 0.0032

Moles of iodine = 0.0016 moles

97
Q

What happens to volatility as you go down group 7 ?

A

Volatility goes decreases as you go down the group > due to increasing strength of the instantaneous dipole - induced dipole bond > these increase as the size and relative mass of the atoms increase.

98
Q

What is volatility ?

A

A measure of how easy it is to vaporise something.

99
Q

What is the appearance of fluorine at room temperature and pressure ?

A

Pale yellow gas

100
Q

What is the appearance of chlorine at room temperature and pressure ?

A

Yellow - green gas

101
Q

What is the appearance of bromine at room temperature and pressure ?

A

Red - brown liquid

102
Q

What is the appearance of iodine at room temperature and pressure ?

A

Shiny grey solid

103
Q

Where are halogens more soluble organic solvents or water ?

A

Organic solvents

104
Q

What do halogens in their natural state exist as ? Why ?

A

Covalent, diatomic molecules.

They are covalent and non-polar that have a low solubility in water, but do dissolve easily in organic compounds like hexane.

105
Q

What colour is chlorine in water ?

A

Virtually colourless

106
Q

What colour is chlorine in hexane ?

A

Virtually colourless

107
Q

What colour is bromine in water ?

A

Yellow / orange

108
Q

What colour is bromine in hexane ?

A

Orange / red

109
Q

What colour is iodine in water ?

A

Brown

110
Q

What colour is iodine in hexane ?

A

Pink / violet

111
Q

Why do halogens get less reactive as you go down the group ?

A

As you go down the group the atoms become larger, so their outer electrons are further from the nucleus. The outer electrons are also shielded more from the attraction of the positive nucleus, because there are more inner electrons. This makes it harder for larger atoms to attract the electron needed to form an ion. So larger atoms are less reactive and reactivity decreases down the group.

112
Q

How do halogens react ?

A

Gaining an electron in their outer p-shell. This means their reduced. As they’re reduced, they oxidise another substance - so they’re oxidising agents.

113
Q

Can halogens displace halide ions from solution ?

A

Yes

114
Q

What is a displacement reaction ?

A

Where one reactant replaces another reactant in a compound.

115
Q

What happens in the displacement reaction when a halogen replaces a halide ?

A

A more reactive halogen will replace a less reactive halide

116
Q

What ions will chlorine replace ?

A

Bromide and Iodide ions

117
Q

What ions will bromine replace ?

A

Iodide ions
NOT CHLORINE

118
Q

What ions will iodide replace ?

A

NOT DISPLACE CHLORIDE OR BROMIDE IONS

119
Q

What is a displacement reaction between a halogen and a halide ?

A

A redox reaction

120
Q

What are the half equations for the reaction of chlorine with potassium bromide ?

A

Cl2 + 2e- –> 2Cl-
Chlorine displaces bromine and is reduced.

2Br- –> Br2 + 2e-
Bromine is displaced by chlorine and gets oxidised.

121
Q

What happens if you add potassium chloride solution and chlorine water ?

A

No reaction

122
Q

What happens if you add potassium chloride solution and bromine water ?

A

No reaction

123
Q

What happens if you add potassium chloride solution an iodine solution ?

A

No reaction

124
Q

What happens if you add potassium bromide solution with chlorine water ?

A

Cl2(aq) + 2Br-(aq) –> 2Cl-(aq) + Br2(aq)

125
Q

What happens if you add potassium bromide solution with bromine water ?

A

No reaction

126
Q

What will happen if you add potassium bromide solution with iodine solution ?

A

No reaction

127
Q

What will happen if you add potassium iodide solution with chlorine water ?

A

Cl2(aq) + 2I-(aq) –> 2Cl-(aq) + I2(aq)

128
Q

What will happen if you add potassium iodide solution with bromine water ?

A

Br2(aq) + 2I-(aq) –> 2Br-(aq) + I2(aq)

129
Q

What will happen if you add potassium iodide solution with iodine solution ?

A

No reaction

130
Q

What colour change will happen if bromide is displaced and bromine is formed ?

A

The mixture will turn orange

131
Q

What colour change will happen if iodine is displaced and iodine is formed ?

A

The mixture will turn brwn

132
Q

How can you see the colour changes more easily ?

A

Mix with an organic solvent like hexane. The halogen present will dissolve in the organic solvent, which settles out as a distinct layer above the aqueous solution.

133
Q

Place the halogens F Cl Br I in order of increasing volatility ?

A

I
Br
Cl
F

134
Q

Place the halogens F Cl Br I in order of increasing reactivity ?

A

i
Br
Cl
F

135
Q

What colour solution is formed when iodine dissolves in hexane ?

A

Pink / Violet

136
Q

What would be seen when chlorine water is added to potassium bromide solution ?

A

Cl2(aq) + 2Br-(aq) –> 2Cl-(aq) + Br2(aq)

137
Q

Astatine is a halogen. it can be found below iodine in Group 7 of the periodic table.

Write an ionic equation for the reaction between iodine solution and sodium astatine ?

For the equation in a), state which substance is oxidised ?

A

I2(aq) + 2At-(aq) –> 2I-(aq) + At2(aq)

Astatine is oxidised

138
Q

A student has a sample of an aqueous potassium halide solution. She knows it contains either chloride, bromide or iodide ions. The student adds a few drops of aqueous bromine solution to the test tube and a reaction take place.

Which halide ion is present in the potassium halide solution ?

What colour will the aqueous solution in the test tube be after the reaction has finished ?

A

Iodide

Brown

139
Q

What are hydrogen halides made from ?

A

A Halogen + Hydrogen

140
Q

How many electrons does the halogens have in their outer shell ?

A

7 electrons

141
Q

How many electrons does hydrogen have in its outer shell ?

A

1 electron

142
Q

What do hydrogen and halides form ?

A

Hydrogen halides

143
Q

How can you make a hydrogen halide ?

A

Using a concentrated acid to a solid

144
Q

What can all hydrogen halides be made
from ?

A

Ionic halide and concentrated phosphoric acid

145
Q

What hydrogen halides can you not make with sulfuric acid ?

A

Hydrogen bromide and hydrogen iodide

146
Q

What other concentrated can you use to make a hydrogen halide ?

A

Sulfuric acid

147
Q

Why does it get more complicated to use sulfuric acid to make hydrogen halides ?

A

Sulfuric acid is an oxidising agent, so it can get involved in redox reactions.

148
Q

What happens when you react sulfuric acid with sodium bromide or sodium iodide ?

A

The bromide ions or iodide ions are oxidised to make bromine or iodine gas. This is because bromine and iodine are strong enough reducing agents to reduce sulfuric acid.

149
Q

What happens when you add sulfuric acid to sodium bromide ?

A

The bromide ions are oxidised to bromine gas and the sulfuric acid is reduced from sulfuric acid to sulfur dioxide.

150
Q

What happens when you add sulfuric acid to sodium iodide ?

A

Iodine is such a strong reducing agent it reduces all the way from sulfuric acid to hydrogen sulfide.

151
Q

Hydrogen fluoride and hydrogen chloride are stable when heated, how ?

A
  • When heated, hydrogen fluoride and hydrogen chloride are stable, and won’t split up into hydrogen and halide ions.
  • Hydrogen bromide will split slightly when heated, and hydrogen iodide even more so.
  • This is because of the strength of the hydrogen-halide bonds.
  • As you go down Group 7, the strength of the bond that the halide forms with the hydrogen gets weaker.
  • This is because the halogen atoms get bigger down the group, meaning the bonding electrons are further away from the nucleus and shielded by more inner electron shells.
152
Q

Are hydrogen halides acidic or alkaline ?

A

Acidic

153
Q

Hydrogen chloride, hydrogen bromide and hydrogen iodide all dissolve in water to to create … ?

A

Strong acids

154
Q

When a hydrogen halide is dissolved in water, what is it called ?

A

Dissociates

155
Q

What does dissociate mean ?

A

Two molecules split apart to form two ions

156
Q

What makes the solution acidic ?

A

The hydrogen ions

157
Q

Which hydrogen halide cannot fully dissociate in water ?

A

Hydrogen fluoride

158
Q

How is the ammonium ion formed ?

A

Ammonia is a base, so it can accept a proton

159
Q

How do you form an ammonium halide ?

A

The ammonium ion can bond with a negative halide ion

160
Q

Why cannot hydrogen fluoride and hydrogen chloride react with sulfuric acid ?

A

They aren’t strong enough reducing agents to reduce the sulfur.

161
Q

Write an equation to show how sodium chloride reacts with sulfuric acid ?

A

NaCl + H2SO4 –> NaHSO4 + HCl

162
Q

Which hydrogen halide is the most thermally stable ? Explain why ?

A

Hydrogen fluoride is the most thermally stable. The H-F bond is the strongest bond, as there is a greater attraction between the outer electron and the nucleus.

163
Q

What is the product when a hydrogen halide reacts with ammonia ?

A

Ammonium halide

164
Q

A student wants to make hydrogen iodide using sodium iodide. Which of these reactants should they use ?
A Hydrochloric acid
B Ammonia
C Phosphoric acid
D Sulfuric acid

A

C

165
Q

A sample of hydrogen bromide gas is dissolved in water to form aqueous solution.

a) State what you would see if universal indicator was added to the solution. Explain your answer.

b) If sulfuric acid is added to the solution, the hydrogen bromide ions reduce the sulfuric acid to sulfur dioxide. Write a full balanced equation for this reaction.

A

a) The solution would turn red/pink because hydrogen bromide solution is strongly acidic.

b) 2HBr + H2SO4 –> Br2 + SO2 + 2H2O

166
Q

Silver ions react with halide ions to form what … ?

A

A precipitate

167
Q

What is a test for halides ?

A
  • Add dilute nitric acid to remove ions which might interfere with the reaction.
  • Then add silver nitrate solution.
168
Q

What precipitate does a fluoride ion
produce ?

A

No precipitate

169
Q

What precipitate does a chloride ion produce ?

A

White precipitate

170
Q

What precipitate does a bromide ions produce ?

A

Cream precipitate

171
Q

What precipitate does a iodide ion produce ?

A

Yellow precipitate

172
Q

What can you add to see the halide more easily ?

A

Ammonia solution

173
Q

AgCl, what is the observation ?

A

Precipitate dissolves in dilute ammonia solution to give a colourless solution

174
Q

AgBr, what is the observation ?

A

Precipitate remains unchanged if dilute ammonia solution is added, but will dissolve in concentrated ammonia solution to give a colorless solution.

175
Q

AgI, what is the observation ?

A

Precipitate does not dissolve, even in concentrated ammonia solution.

176
Q

Why is transporting chlorine a risky
business ?

A
  • Chlorine is a dangerous chemical
  • It is toxic and corrosive, so must be kept away from skin and eyes. It is harmful if it is breathed in, it irritates the respiratory system.
  • Chlorine must be kept away from the flammable materials.
  • Chlorine has to be transported and stored very carefully, it is kept as a liquid under pressure in small cylinders.
177
Q

Why is chlorine useful ?

A
  • Important part of water treatment. Adding chlorine too water sterilise it making it safe to drink or swim in.
  • It kills disease-causing microorganism, such as bacteria
  • If we didn’t treat our drinking water with chlorine in this way, we’d be at risk of getting all sorts of nasty and potentially dangerous infections.
  • Chlorine is also used to make bleach
178
Q

What is atom economy ?

A

A measure of the proportion of the reactant atoms that become part of the desired product in the balanced chemical equation.

179
Q

Calculate the atom economy of the reaction of iron oxide with carbon dioxide to make iron ?

Fe2O3 + 3CO –> 2Fe + 3CO2

A

% atom economy = (molecular mass of desired product) / (sum of molecular mass of all products) x 100
= (2 x 55.8) / (2 x 55.8 + 3 x (12.0 +(16.0 x 2)) x 100
= (111.6) /(111.6 + 132) x 100 = 45.8 %

180
Q

What is the equation for atom economy?

A

(Molecular mass of desired products) / (Sum of molecular mass of all products) x 100

181
Q

Why are reactions with higher atom economies better than reactions with lower atom economies ?

A

Produce less waste, making them cheaper and more environmentally friendly.

182
Q

Which is more cost effective a reaction with a high atom economy but a low percentage yield or a moderate atom economy and a moderate percentage yield ?

A

Moderate atom economy and a moderate percentage yield

183
Q

What conditions are expensive to maintain ?

A

High temperatures
High pressures

184
Q

Write an equation for the reaction between silver ions and halide ions ?

A

Ag+(aq) + X-(aq) –> AgX(s)

185
Q

How could you distinguish between a solution containing bromide ions and a solution containing iodide ions ?

A

Bromide ions create a cream precipitate
Iodide ions create a yellow precipitate

186
Q

What formula would you use to work out the atom economy of a reaction ?

A

Atom economy
(molecular mass of desired product) / (sum of molecular mass of all products) x 100

187
Q

Industrially, chlorine can be produced by electrolysing sodium chloride solution.

a) Describe a test you could perform to confirm that a sample of a solution contained chloride ions. Include any observations that you would make if the results of the test was positive ?

b) Explain why chlorine needs to be handled and stored extremely carefully ?

A

a) Add dilute nitric acid tp the sample. Then add silver nitrate solution. If the sample contains ions, a white precipitate will form.

b) Chlorine is toxic and corrosive, so must be kept away from skin and eyes.
Chlorine irritates the respiratory system, so must not be breathed in.

188
Q

Calculate the atom economy of this reaction. Methanol is the desired product.

CH3Br + NaOH –> CH3OH + NaBr

A

(molecular mass of desired product) / (sum of molecular mass of all products) x 100

(CH3OH) / (CH3OH + NaBr) x 100
(32.0) / (32 + 102.9) x 100 = 23.7 %

189
Q

Can reversible reactions reach dynamic equilibrium ?

A

Yes

190
Q

What system can dynamic equilibrium be
in ?

A

A closed system

191
Q

If the equilibrium moves to the left what happens to the reactants and products ?

A

There are more reactants than products

192
Q

If the equilibrium moves to the right what happens to the reactants and products ?

A

There are more products then reactants

193
Q

What is the general equation for Kc ?
aA + bB –> dD + eE

A

[D]^d x [E]^e / [A]^a x [B]^b

[products] / [reactants]

194
Q

Where is the equilibrium if the value of Kc is larger ?

A

The equilibrium is further to the right, there are more products then reactants.

195
Q

Where is the equilibrium if the value of Kc is smaller ?

A

The equilibrium is further to the left, there are more reactants then products.

196
Q

Hydrogen and iodine are placed in a closed flask and allowed to react. What is the equilibrium constant, Kc, for the reaction at 640K ?

The equilibrium constants are:
[HI] = 0.8mol/dm3
[H2] = 0.1 mol/dm3
[I2] = 0.1 mol/dm3

A

Hydrogen + Iodine –> Hydrogen iodide
H2 + I2 –> 2HI

[0.8]^2 / [0.1] x [0.1] = 64

197
Q

What happens to the equilibrium if you change the concentration of the products and reactants ?

A

The equilibrium will shift to keep Kc constant

198
Q

What happens if you increase the pressure ?

A
  • Shifts the equilibrium to the side with fewer molecules of gas.
199
Q

What happens if you decrease the pressure ?

A
  • Shifts the equilibrium to the side with more gas molecules
200
Q

What happens if you increase the the temperature

A
  • The equilibrium will shift in the direction of the endothermic reaction (positive).
201
Q

What happens if you decrease the temperature ?

A
  • The equilibrium shifts in the direction of the exothermic (negative)
202
Q

Explain what is happening when a reaction reaches a ‘dynamic equilibrium’ ?

A

The forward reaction is going at the same rate as the reverse reaction.

203
Q

Describe how changing the temperature affects the equilibrium position of a reversible reaction where the forward reaction is exothermic ?

A

Decreasing the pressure, favours the exothermic reaction. The equilibrium shifts to the right, there are more products than reactants.

204
Q

Dichromate(VI) ions and water were mixed together and allowed to reach equilibrium at a fixed temperature.

Cr2O72-(aq) + H2O(l) –>
2CrO42-(aq) + 2H+(aq)

a) Write an expressions fro Kc

b) Calculate the value of the equilibrium constant RTP, given that the equilibrium concentrations of the reactants and products are:
[Cr2O72-] = 0.20mol/dm3
[H2O] = 0.60mol/dm3
[H+] = 0.30mol/dm3
[CrO42-] = 0.10mol/dm3

c) What would happen to the value of Kc if the concentration of the water increased ? Explain your answer.

A

a) [CrO42-]^2 x [H+]^2 / [Cr2O72-] x [H2O]

b) Kc = [0.20]^2 x [0.30]^2 / [0.20] x [0.60]
Kc = 0.0075

c) There would be no change as the concentrations of the other reagents would change to keep Kc constant.

Chemistry (8 decks)

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