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Chemistry > Topic 4 Inorganic Chemistry & the Periodic Table > Flashcards

Topic 4 Inorganic Chemistry & the Periodic Table Flashcards

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1
Q

Appearance of group 2 elements

A

Pure: bright, silvery solids.
Exposed to air: they form oxides surface layers, so appear dull.

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2
Q

Factors to consider when explaining trends in ionisation energy

A

Nuclear charge.
The orbital in which the electron exists.
The shielding effect– the repulsion between the filled inner shells & the electron being removed.

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3
Q

Why do first and second ionisation energies decrease down group 2?

A

As the nuclear charge increases, the force of attraction for the electron being removed also increases. This is outweighed by the combined effect of:
As each quantum shell is added, the energy of the outermost electrons increases.
As the number of filled inner shells increases, their force of repulsion on the electron being removed increases.

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4
Q

Why is there a general increase in reactivity down group 2?

A

The element becomes a M2+ ion, and 2nd ionisation energy/ the energy needed to remove 2 electrons from each atom of the element decreases down the group.

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5
Q

What must be included in ionisation energy equations after each atom & ion?

A

The state symbol (g).

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6
Q

What happens when magnesium burns in air?

A

There is a very bright flame & a formation of white solid.

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7
Q

What happens when calcium, strontium & barium are burnt in air?

A

The same as magnesium, but more vigorous.

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8
Q

If the metals are fresh samples, it is hard to distinguish between their reactions in air. How can this be overcome?

A

If the burning metal is placed in a glass jar of oxygen, the same reaction occurs but more vigorously.

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9
Q

What happens to the group 2 element in oxygen without heating?

A

A slow reaction that forms a surface coating of oxide, which prevents the element from further reaction.

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10
Q

Why is barium stored under oil?

A

To prevent it reacting with oxygen and water vapour in the air, as it is the most reactive.

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11
Q

General equation for group 2 elements + oxygen:

A

2M(s) + O2(g) –> 2MO(s) (oxides).

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12
Q

What happens to reactions with chlorine down group 2?

A

They get more vigorous, but this is harder to see than in oxygen.

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13
Q

General equation for group 2 elements + chlorine (requires heating):

A

M(s) + Cl2 (g) –> MCl2 (s) (chlorides).

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14
Q

Magnesium + water

A

Very slow, does not proceed completely; not very vigorous. The piece of magnesium is covered in bubbles of hydrogen gas.

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15
Q

How can the increasing vigour of the reactions of calcium, strontium & barium with water be observed?

A

Increasing effervescence.

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16
Q

General equation for group 2 elements + water:

A

M(s) + 2H2O (l) –> M(OH)2 (aq) + H2 (g) (hydrogen gas + hydroxide).

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17
Q

Reaction of calcium + water

A

Ca (s) + 2H2O (l) –> Ca(OH)2 (s) + H2 (g) Calcium hydroxide is only slightly soluble in water, so the liquid goes cloudy as a precipitate of calcium hydroxide forms.

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18
Q

Reaction of barium + water

A

Ba (s) + 2H2O (l) –> Ba(OH)2 (aq) + H2 (g) Barium hydroxide is soluble in water, hence (aq).

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19
Q

What happens when magnesium is heated in steam?

A

It rapidly forms magnesium oxide ( a white solid) + hydrogen gas in a vigorous reaction.

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20
Q

Equation for when magnesium is heated in steam:

A

Mg (s) + H2O (g) –> MgO (s) + H2 (g)

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21
Q

In the reaction between Mg + steam, why is hydrogen burnt as it leaves the tube?

A

Safety! To prevent the release of a highly flammable gas into the lab.

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22
Q

Group 2 oxides

A

Basic oxides: they can react with water to form alkalis.

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23
Q

General equation for a group 2 oxide + water:

A

MO (s) + H2O (l) –> M(OH)2 (aq) (Hydroxide.) Colourless solution.
This can be simplified to:
O2- + H2O –> 2OH- Colourless solution.
The formation of hydroxide ions = alkaline solution.

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24
Q

How does the solubility of group 2 hydroxides change down the group?

A

Solubility of group 2 hydroxides increases down the group because the relative decrease in lattice enthalpy is more than the relative decrease in hydration enthalpies.

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25
Q

How does the maximum alkalinity/pH value of group 2 hydroxides in solution change down the group?

A

pH increases down the group, as solubility increases.

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26
Q

Limewater

A

A saturated, aqueous solution of calcium hydroxide.

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27
Q

Carbon dioxide + limewater

A

Carbon dioxide reacts to form an insoluble, white precipitate of calcium carbonate.
CO2 + Ca(OH)2 –> CaCO3 + H2O

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28
Q

What happens as carbon dioxide is bubbled through limewater?

A

The amount of calcium carbonate precipitate increases.

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29
Q

What is milk of magnesia?

A

The saturated solution of magnesium hydroxide mixed with extra solid magnesium hydroxide acts as an antacid.

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30
Q

How does milk of magnesia relieve symptoms of indigestion?

A

It neutralises some of the hydrochloric acid: Mg(OH)2 + 2HCl –> MgCl2 + 2H2O.

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31
Q

Hydroxide ions attack human tissue. How does milk of magnesia not pose a risk to human health?

A

The very low solubility of magnesium hydroxide means the concentration of OH- ions is also very low.

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32
Q

What can be observed in neutralisation reactions of Group 2 oxides & hydroxides with acids?

A

Salt + water form. A white solid reacts to form a colourless solution.

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33
Q

Why are the neutralisation reactions of group 2 oxides & hydroxides with acid used in experiments to measure energy changes?

A

They are exothermic.

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34
Q

Why do farmers control soil acidity?

A

To obtain a greater yield of crops.

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35
Q

Lime is used to neutralise excess acidity in the soil. What is lime?

A

Mostly calcium hydroxide obtained from limestone (which is calcium carbonate).

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36
Q

How does lime neutralise excess acidity in the soil? Equation, where HNO3 represents the acid:

A

Ca(OH)2 + 2HNO3 –> Ca(NO3)2 + 2H2O

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37
Q

All group 2 nitrates & chlorides…

A

…are soluble!

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38
Q

How does the solubility of group 2 sulfates change down the group?

A

Sulfate solubility decreases because the relative decrease in lattice enthalpy is less than the relative decrease in hydration enthalpies.

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39
Q

Magnesium sulfate

A

Soluble.

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40
Q

Calcium sulfate

A

Slightly soluble.

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41
Q

Strontium sulfate + barium sulfate

A

Insoluble.

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42
Q

How can we test for sulfate ions in aqueous solution?

A

Add a solution containing Ba2+ ions (usually barium chloride or barium nitrate). Add dilute nitric acid or dilute hydrochloric acid. Any sulfate ions will react with the added barium ions to form a white precipitate of barium sulfate.

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43
Q

What is the ionic equation for the test for sulfate ions?

A

Ba2+ (aq) + SO4 2- (aq) –> BaSO4 (s)

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44
Q

Why is dilute hydrochloric or dilute nitric acid added as part of the test for sulfate ions?

A

Carbonate ions could also react with barium ions to form a white precipitate, so H+ ions must be present to prevent barium carbonate forming.

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45
Q

Why is barium meal not poisonous to humans?

A

The meal contains barium sulfate, which is insoluble– the ions are not free to move.

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46
Q

What is barium meal used for?

A

The dense white solids means soft tissues show up better on X-rays.

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47
Q

Why is there no trend in reactivity with water for the group 2 oxides?

A

They already contain metal ions, not metal atoms.

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48
Q

Basic oxides

A

Oxides of metals that react with water to form metal hydroxides, and with acids to form salts + water.

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49
Q

Thermal stability

A

A measure of the extent to which a compound decomposes when heated.

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50
Q

Very thermally stable compounds

A

Do not decompose at all when heated.

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51
Q

How do group 2 nitrates & carbonates compare with Group 1 chlorides, e.g., NaCl?
Why do the former decompose whilst the latter melt when heated?

A
  • The charge on a group 2 cation is double that of a group 1 cation.
  • The ionic radius of a group 2 cation is smaller than that of a group 1 cation in the same period.
  • The nitrate NO3 - & carbonate CO3 2- anions are more complex than the Cl- ion.
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52
Q

Why do group 2 nitrates and carbonates decompose when heated?

A
  • The larger, more complex nitrate ion can change into the smaller, more stable nitrite (NO2 -) ion or oxide ion (O2-) by decomposing & releasing oxygen gas and/or nitrogen dioxide gas.
  • The larger, more complex carbonate ion (CO3 2-) ion can change into the smaller, more stable oxide ion (O2-) by decomposing & releasing CO2 gas.
  • The stabilities of the nitrate & carbonate anions are influenced by the charge & size of the cations present; cations with higher charge density polarise the anions more.
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53
Q

Appearance of group 1 & 2 nitrates & carbonates

A

White solids.

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54
Q

What happens when group 1 and 2 nitrates are heated?

A

They decompose to nitrites or oxides, & emit brown fumes of nitrogen dioxide and/or oxygen.

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55
Q

What is also observed when the nitrate contains water of crystallisation?

A

Steam!

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56
Q

What has occurred if NO brown fumes are observed?

A

Lesser decomposition: metal nitrate –> metal nitrite + oxygen.

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57
Q

What has occurred if brown fumes are observed?

A

Greater decomposition: metal nitrate –> metal oxide + nitrogen dioxide + oxygen.

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58
Q

Nitrate(V)

A

Nitrate

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59
Q

Nitrate(III)

A

Nitrite

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60
Q

When does greater decomposition occur?

A

When a cation has a 2+ charge, or when the cation has a 1+ charge & is the smallest group 1 cation.
I.e., All the group 2 nitrates + LiNO3.

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61
Q

What happens when group 1 & 2 carbonates are heated?

A

They either:
- Do not decompose.
- OR they decompose to an oxide + carbon dioxide gas.

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62
Q

What is observed when group 1 & 2 carbonates are heated?

A

Nothing. CO2 gas is colourless, and the carbonate & the oxide are both white solids.

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63
Q

Decomposition of group 1 carbonates

A

Li2CO3 –> Li2O + CO2
Lithium carbonate decomposes at lower temperatures than the other group 1 carbonates. Other group 1 carbonates do not decompose on heating, except at very high temperatures.

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64
Q

Decomposition of group 2 carbonates

A

All decompose in the same way, but with increasing difficulty down the group.
CaCO3 –> CaO + CO2

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65
Q

What are flame tests used for?

A

To indicate the presence of some metal cations in groups 1 & 2 of the periodic table.

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66
Q

Safety precautions for a flame test

A

Goggles, lab coat & light a Bunsen burner in a fume cupboard.

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67
Q

How do you carry out a flame test?

A
  1. Add a few drops of concentrated hydrochloric acid to the solid using a dropper.
  2. Mix together, so that the metal compound begins to dissolve.
  3. Put a platinum or nichrome wire or silica rod in the Bunsen flame, then in conc. HCl then into the mixture to obtain a sample of the compound.
  4. Hold the end of the wire or rod in the flame, & observe the colour.
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68
Q

Why is hydrochloric acid used to dissolve the metal compound in a flame test?

A

To convert any metal compound to a chloride. Chlorides are more volatile than other salts, so are more likely to give better results.

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69
Q

2 main problems with flame tests

A

Many compounds contains small amounts of sodium compounds as impurities, so the intense colour of sodium can mask other colours.
Describing colour with words is subjective; a word description of colour may mean different things to different people. People have different levels of colour vision, too.

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70
Q

Li+ flame test colour

A

red

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71
Q

Na+ flame test colour

A

yellow/orange

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72
Q

K+ flame test colour

A

lilac

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73
Q

Rb+ flame test colour

A

red (red-purple/violet)

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74
Q

Cs+ flame test colour

A

blue/violet

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75
Q

Be2+ flame test colour

A

No colour

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76
Q

Mg2+ flame test colour

A

No colour

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77
Q

Ca2+ flame test colour

A

brick red

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78
Q

Sr2+ flame test colour

A

Crimson red

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79
Q

Ba2+ flame test colour

A

apple green

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80
Q

Copper compounds flame test colour

A

blue-green

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81
Q

Ground state

A

An atom with all its electrons in their lowest possible energy levels.

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82
Q

Excited state

A

When an electron absorbed energy and moves to a higher energy level.

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83
Q

What causes the colours in flame tests?

A

Electrons absorb energy and move to a higher energy level– from the ground state to the excited state. This is immediately followed by the return of the electron to its ground state, which releases energy as light. If this energy corresponds to radiation in the visible light spectrum, the characteristic colour appears.

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84
Q

why is there no flame test colour for magnesium?

A

The electron transition corresponds to a wavelength outside the visible light spectrum (400-700nm).

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85
Q

Ammonium ions do not give a colour in the flame test. What is the test used instead?

A

Add sodium hydroxide solution, & warm the mixture.
NH4 + + OH- –> NH3 + H2O
The warming results in ammonia being released as gas. This turns damp red litmus paper blue (ammonia is the only common alkaline gas).

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86
Q

Alternative ammonium ion test

A

Hydrogen chloride gas from concentrated hydrochloric acid reacts with ammonia to form white fumes of ammonium chloride.
NH3 + HCl –> NH4Cl

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87
Q

Why are astatine & radium ignored?

A

They only exist as radioactive isotopes.

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88
Q

How do boiling and melting point change down group 7?

A

Boiling & melting temperatures increase down group 7.

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89
Q

How does electronegativity change down group 7?

A

Electronegativity decreases.

90
Q

Why are halogen molecules non-polar, on average?

A

The two atoms in the diatomic molecule are identical; the pair of electrons forming the covalent bond between them is shared equally between the two atoms.

91
Q

How do temporary/instantaneous dipoles arise in halogen molecules?

A

The positive charges of the protons in the two nuclei are in fixed positions, but the electron density continuously fluctuates. Sometimes the centres of positive & negative charge do not coincide.

92
Q

How do instantaneous dipole-induced dipole forces of attraction form between halogen molecules?

A

If a molecule on the left becomes an instantaneous dipole, it will cause an induced dipole in the molecule on the right. There is a force of attraction between the two molecules.

93
Q

Why do melting 7 boiling point of halogens increase down group 7?

A

Weak instantaneous dipole-induced dipole forces increase as the number of electrons & the size of the electron cloud increase, so the forces increase in strength down group 7 as the number of electrons in the molecules increases.

94
Q

Why does bromine give off brown vapour at room temperature?

A

Its boiling point (59 degrees C) is not much higher than room temperature.

95
Q

When iodine is warmed, what is the equation for its change in state?

A

I2 (s) –> I2 (g). Iodine sublimes as it changes directly into a vapour without melting.

96
Q

Equations for changes in state.

A

Write the halogen as a diatomic molecule. Use state symbols.

97
Q

What factors does electronegativity depend upon?

A
  • Its nuclear charge– the bigger the nuclear charge, the higher the electronegativity.
  • The distance between the nucleus & the bonding pair of electrons– the shorter the distance, the higher the electronegativity.
  • The shielding effect of electrons in inner energy levels– the fewer the energy levels, the higher the electronegativity.
    Less shielding/fewer energy levels, shorter distance; higher nuclear charge = more electronegative.
98
Q

The electronegativity of group 7 elements

A

The highest of any group in the periodic table.

99
Q

Why does reactivity decrease down group 7?

A

Due to their high electronegativity, most reactions of halogens involve them acting as oxidizing agents to form negative ions, or becoming partially negative as part of a polar molecule. Thus, reactivity is affected by the same factors as electronegativity.

100
Q

Between which elements are reactions most vigorous?

A

Elements at the bottom of group 1 & 2 and elements at the top of group 7. The products of these reactions are salts– ionic, usually white solids.

101
Q

Reactions between group 1 & 2 metals + halogens

A

All involve electron transfer to the halogen, so are redox reactions in which the halogen acts as an oxidizing agent.
The oxidation number of the halogen decreases from 0 to -1. The oxidation number of the metal increases from 0 to +1 or +2, depending on the group.

102
Q

Halogen displacement reactions

A

Take place in aqueous solution, so are indicated by a colour change. A more reactive halogen displaces a less reactive halogen from one of its compounds.

103
Q

What are the problems with interpreting colour changes in halogen displacement reactions?

A

Similarity of some colours and the variation in colour with concentration. E.g. liquid bromine is red-brown, and iodine dissolved in water may also appear brown at some concentrations.

104
Q

How can we make the colour change more visible in halogen displacement reactions?

A

Add an organic solvent, such as cyclohexane, after the reaction. Shake the tube. Halogens are more soluble in cyclohexane than in water, so the halogen dissolves in the organic upper layer where its colour can more easily be seen.

105
Q

How do oxidation numbers change in halogen displacement reactions?

A

Reacting halogen: 0 to -1.
reacting halide: -1 to 0.

106
Q

Of chlorine, bromine & iodine, why is chlorine most reactive?

A

It’s the smallest atom, so the incoming electron gets closer and is more attracted to the protons in the nucleus.
It has the smallest number of complete inner energy levels of electrons, so the incoming electron experiences the least repulsion.

107
Q

Disproportionation

A

The simultaneous oxidation & reduction of an element in a single reaction.

108
Q

When chlorine dissolves in water, some of the dissolved chlorine reacts to form a mixture of hydrochloric acid + chloric (I) acid. How can this disproportionation reaction be written as an equation?

A

Cl2 + H2O –> HCl + HClO

109
Q

Sodium chlorate (I) NaClO

A

Disinfectant used in industry & in bleach.

110
Q

Disproportionation reaction when chlorine is added to cold, dilute; aqueous NaOH.
(Hint: this forms the salts of the acids made when chlorine dissolves in water.)

A

Cl2 + 2NaOH –> NaCl + NaClO + H2O. Forms sodium chloride + sodium chlorate.

111
Q

Disproportionation reaction when chlorine is added to hot, concentrated NaOH.

A

3Cl2 + 6NaOH –> 5NaCl + NaClO3 + 3H2O
Sodium chloride + sodium(V) chlorate form.

112
Q

How does the reducing power of the halides change down the group?

A

tIt increases down the group. Astatide has he highest reducing power.

113
Q

Half-equation for the reducing action of halides:

A

2X- –> X2 + 2e-

114
Q

Equation for the partial ionisation of sulfuric acid:

A

H2SO4 –><– H+ + HSO4 -

115
Q

Properties of concentrated sulfuric acid

A

It can act as an oxidising agent as well as an acid. The extent of its reduction & the the products formed depends on the species being oxidised.

116
Q

3 possible reduction products of sulfuric acid

A

Sulfur dioxide
Sulfur
Hydrogen sulfide

117
Q

In half-equations involving sulfuric acid, what patterns are there regardless of which product (+ water is formed)?

A

The decrease in oxidation number of the sulfur is equal to the number of + ions and electrons on the LHS.

118
Q

Concentrated sulfuric acid + NaCl

A

Misty fumes. Produces HCl.

119
Q

Concentrated sulfuric acid + NaBr

A

Misty fumes = HBr
Brown fumes = Br2
Colourless gas with a choking smell = SO2.

120
Q

Concentrated sulfuric acid + NaI

A

Misty fumes = HI
Purple fumes/black solid = I2
Colourless gas with a choking smell= SO2
Yellow solid = sulfur, S
Colourless gas with a rotten egg smell= hydrogen sulfide, H2S.

121
Q

With NaCl, why does the concentrated sulfuric acid behave only as an acid?

A

The chloride ions have low reducing power, so the H2SO4 does not act as an oxidising agent.
NaCl + H2SO4 –> NaHSO4 + HCl

122
Q

With the NaBr, what happens to the concentrated sulfuric acid?

A

The greater reducing power of the bromide ions causes the H2SO4 to be reduced.
H2SO4 + 2H+ + 2e- –> 4H2O + SO2

123
Q

With NaI, what happens to the concentrated sulfuric acid?

A

The much greater reducing power of the I- causes the sulfuric acid to be reduced.
H2SO4 + 2H+ + 2e- –> 4H2O + SO2
H2SO4 + 8H+ + 8e- –> 4H2O + H2S
H2SO4 + 6H+ + 6e- –> 4H2O + S

124
Q

When testing for halide ions in solution, why is nitric acid added beforehand?

A

To ensure that any other anions, especially carbonate ions, are removed, as they would also form precipitates.

125
Q

How can you test for halide ions in solution?

A

Add dilute nitric acid.
Add silver nitrate solution.
If a precipitate is obtained, add ammonia solution.

126
Q

Why can the test for halide ions not be used for fluoride?

A

Silver fluoride is soluble.

127
Q

Chloride ions– result in the halide ions test.

A

White precipitate soluble in aqueous ammonia.

128
Q

Bromide ions– result in the halide ions test.

A

Cream precipitate. Insoluble in dilute, but soluble in concentrated aqueous ammonia.

129
Q

Iodide ions– result in the halide ions test.

A

Yellow precipitate. Insoluble in aqueous ammonia.

130
Q

General ionic equation for halides + silver nitrate:

A

Ag+ (aq) + X- (aq) –> AgX (s)

131
Q

Equation for the dissolving of silver chloride in aqueous ammonia:

A

AgCl (s) + 2NH3 (aq) –> [Ag(NH3)2]+ (aq) + Cl- (aq)

132
Q

Common properties of hydrogen halides

A

Exist as polar, diatomic molecules. Colourless gases that react with water to form acidic, colourless solutions.

133
Q

Acid formed from hydrogen iodide

A

Hydriodic acid

134
Q

General equation for acid formation from hydrogen halides:

A

HX + H2O –> H3O+ + X-

135
Q

Hydrofluoric acid is a weak acid. Write its formation equation:

A

HF + H2O reversible arrow <– –> H3O+ + F-

136
Q

Hydrogen halides + ammonia gas

A

Form salts, all of which are white ionic solids.
NH3 (g) + HCl (g) –> NH4Cl (s)

137
Q

Why does barium metal stop reacting with sulfuric acid even when there is an excess of acid?

A

A precipitate of barium sulfate forms around the barium metal, which stops it reacting with the acid.

138
Q

Why does thermal stability of group 2 carbonates and nitrates increase down the group?

A

The polarizing power of the M2+ cation decreases down the group due to increasing ionic radius. The electron cloud of the anion is less polarized, so the C-O/N-O bond is weakened less.

139
Q

How are group 1 salts prepared?

A

Acid-base titration. The solution is boiled to evaporate most of the water, cooled to form crystals, filtered to remove any more water and dried with a paper towel.

140
Q

How are group 2 & transition metal salts prepared?

A

Acid + insoluble metal base. The base is added in excess. Then filtered to remove unreacted base. The solution is boiled to evaporate most of the water, cooled to form crystals, filtered to remove any more water and dried with a paper towel.

141
Q

Why does the thermal stability of group 1 carbonate & nitrates increase down the group?

A

Only Li+ is sufficiently polarising. Ionic radius increases down the group. The polarizing ability of the cation decreases, so the anion’s electron cloud is less polarised down the group. The C-O/N-O bond is not sufficiently weakened.

142
Q

How can you identify a sulfate?

A

Add HCl + BaCl2 in equal volumes. A white precipitate of barium sulfate forms. The HCl is to remove the BaCO3 formed due to carbonate impurities.

143
Q

Cl2 colours

A

Standard state: green. Pale green in organic solvents or water.

144
Q

Br2 colours

A

Standard state: reddish/brown. Orange in organic solvent or water.

145
Q

I2 colours

A

Standard state: grey. In water: brown. In organic solvent: violet.

146
Q

How does electron affinity change down group 7?

A

It becomes less exothermic.

147
Q

Alternative test for carbonate/hydrogencarbonate ions

A

Add any dilute acid –> effervescence as CO2 is emitted.

148
Q

Test for SO2

A

Turns orange potassium dichromate paper green.

149
Q

Test for O2

A

Relights a glowing splint.

150
Q

Test for H2

A

It will explode with a squeaky pop when ignited.

151
Q

Test for hydrogen halide gases

A

White smoke is produced when they react with vapours from concentrated NH3.

152
Q

Precipitate formed when Mg2+ is reacted with NaOH (aq)

A

white

153
Q

Precipitate formed when Ca2+ is reacted with NaOH (aq)

A

white

154
Q

Precipitate formed when Cu2+ is reacted with NaOH (aq)

A

blue

155
Q

Precipitate formed when Fe2+ is reacted with NaOH (aq)

A

green

156
Q

Precipitate formed when Fe3+ is reacted with NaOH (aq)

A

brown

157
Q

Precipitate formed when Al3+ is reacted with NaOH (aq)

A

White, but is the only one to dissolve in excess NaOH.

158
Q

Why might adding sulfuric acid to what is thought to be barium carbonate NOT be a valid test for the a carbonate ion?

A

The expected result may not occur as BaCO3 is insoluble, so may not react with the acid.

159
Q

Summary of how to carry out a flame test

A

Dip a nichrome wire loop into clean/fresh/concentrated HCl.
Dip the wire into the solid, and place into a roaring flame.

160
Q

Equation for the thermal decomposition of lithium nitrate:

A

2 LiNO3 –> 2 NO2 + Li2O + 1/2 O2
Forms nitrogen dioxide lithium oxide + oxygen.

161
Q

Why might experimental gas volume collected from thermal decomposition differ from the theoretical volume?

A

Incomplete reaction/ decomposition.

162
Q

What does acidified barium chloride test for?

A

Sulfate ions.

163
Q

Describe how to compare the thermal stability of group 2 nitrates.

A

Use a fume cupboard.
Method 1: same volume of nitrate in different test tubes. Same heat applied. First to relight a glowing splint.
Method 2: Use a gas syringe, and time how long it takes to reach a certain gas volume.

164
Q

Why is a lack of flame colour not a positive test for magnesium ions?

A

Lots of other ions do not produce a flame colour, such as beryllium, Be2+.

165
Q

How can the ease of thermal decomposition of carbonates be compared?
State a control variable.

A

Heat a mass of carbonate in a test tube with a bung. There is a delivery tube to another test tube in which there is limewater. Time how long it takes for a cloudy precipitate to form.
Same distance between the flame, and the test tube.

166
Q

How can atoms of different elements produce characteristic flame colours when heated?

A

Different energy gaps in different elements, so different colours are emitted.

167
Q

Why does carrying out a flame test on a mixture of potassium & strontium chloride not clearly show two different metal ions are present?

A

The red colour masks the lighter lilac colour.

168
Q

Why is a nichrome wire used in flame tests?

A

It is inert and produces no colour when heated in a flame test.

169
Q

Why is the nichrome wire dipped into acid and heated in the first stage of a flame test?

A

To keep the wire clean & uncontaminated. To remove the residue of any previous sample being tested.

170
Q

What is observed when a hydrated group 2 nitrate is heated?

A
  • H2O is emitted & seen as condensation.
  • Nitrogen dioxide is released as brown fumes.
  • A white solid forms.
171
Q

Why is fresh concentrated HCl used in the 2nd stage of a flame test?

A

The acid can become contaminated by residue from previous tests giving incorrect results.

172
Q

What is observed when a metal nitrate is heated?

A

Brown fumes of nitrogen dioxide gas.

173
Q

Why is a high temperature required for the decomposition of carbonates?

A

Strong bonds within the carbonate ion.

174
Q

What type of reaction occurs when ammonia reacts with HCl gas?

A

Acid-base.

175
Q

Which is the least soluble group 2 hydroxide?

A

Mg(OH)2

176
Q

In the test for sulfate ions, which acids can be used to acidify Ba(NO3)2/BaCl2?

A

HCl or HNO3

177
Q

Why does bromine disproportionation still occur to a small extent, even though it is not thermodynamically feasible?

A

Disproportionation is an equilibrium system.

178
Q

What must always go before oxidation number?

A

+ or -

179
Q

HCl gas does not conduct electricity. Why is HCl a good electrical conductor as an acid?

A

The covalent bond in hydrogen chloride changes to an ionic bond in aqueous solution.

180
Q

Titration tips

A

White tile. Add slowly. Swirl the flask.

181
Q

Hazards associated with bromine

A

It is corrosive. Wear gloves.

182
Q

How are London forces produced?

A

Uneven distribution of electrons due to their random movement results in a temporary dipole in the 1st molecule.
This induces a dipole on another molecule.

183
Q

How can a less soluble compound be extracted from a reaction mixture?

A

Cool the mixture to give time for crystallisation.
Filter.
Pat dry with a paper towel.

184
Q

How can you show that I- ions are present in an aqueous solution of KI?

A

Add AgNO3. A yellow precipitate forms.

185
Q

When KBr (aq) is added to I2 (aq) there is no reaction. Why?

A

Iodine cannot oxidise Br- ions.

186
Q

Trend in the reducing ability of hydrogen halides

A

Hydrogen halides are better reducing agents down the group.
E.g, H2SO4 is reduced further by HI than HBr or HCl.

187
Q

F- ions + H2SO4

A

NaF(s) + H2SO4(l) —> NaHSO4(s) + HF(g)
White, steamy fumes of HF are evolved.
Only acid-base. Not strong enough to reduce S in H2SO4.

188
Q

Cl- ions + H2SO4

A

NaCl (s) + H2SO4(l) —> NaHSO4(s) + HCl(g)
White, misty fumes of HCl are evolved.
Only acid-base. Not strong enough to reduce S in H2SO4.

189
Q

Br- ions + H2SO4

A

Acid-base step: NaBr(s) + H2SO4(l) —> NaHSO4(s) + HBr(g)
Redox step: 2HBr + H2SO4 —> Br2(g) + SO2(g) + 2H2O(l)
White fumes of HBr, red fumes of Br2 & a colourless, acidic SO2 gas.
Ionic: 2Br- + 2H+ + H2SO4 —> Br2 + SO2 + 2H2O

190
Q

Half-equations for Br- ions + H2SO4

A

Oxidation half-equation: 2Br- –> Br2 + 2e-
Redox half-equation:
H2SO4 + 2H+ + 2e- –> SO2 + 2H2O

191
Q

Half-equations for I- ions + H2SO4

A

Oxidation half-equation: 2I- –> I2 + 2e-
Redox half-equations:
H2SO4 + 2H+ + 2e- –> SO2 + 2H2O
H2SO4 + 6H+ + 6e- –> S + 4H2O
H2SO4 + 8H+ + 8e- –> H2S + 4H2O

192
Q

Equations for I- ions + H2SO4

A

Acid-base step: NaI (s) + H2SO4(l) —> NaHSO4(s) + HI(g)
3 redox steps:
2HI + H2SO4 —> I2(s) + SO2(g) + 2H2O(l)
6HI + H2SO4 —> 3I2(s) + S(s) + 4H2O(l)
8HI + H2SO4 —> 4I2(s) + H2S(g) + 4H2O(l)

193
Q

Observations for I- ions + H2SO4

A

White, steamy fumes of HI, purple fumes & a black solid of I2, a colourless, acidic gas of SO2, yellow sulfur solid & the rotten egg smell of hydrogen sulphide.

194
Q

When might the test for carbonate ions fail?

A

When the salt is insoluble, e.g., BaCO3, so it may not react with the acid.

195
Q

Flame tests

A

Dip the nichrome wire loop into clean/fresh/concentrated HCl. Dip the wet wire into the solid, and place into the roaring flame.

196
Q

Equation for the decomposition of lithium nitrate

A

2LiNO3 –> 2NO2 + Li2O + 1/2O2

197
Q

Why might the experimental gas volume produced by decomposition differ from the actual value?

A

Incomplete decomposition.H

198
Q

How can you experimentally determine the thermal stability of group 2 nitrates?

A

Same volume of nitrate in each test tube. Same heat applied. The first to relight a glowing splint, or use a gas syringe and time how long it takes for a certain volume to be produced. Ensure there is the same distance between the flame & each tube.

199
Q

Why do different elements produce different flame colours when heated?

A

Different energy gaps in different elements, so different colours are emitted.

200
Q

Why is a nichrome wire used?

A

It is inert and doesn’t produce a colour when heated in a flame test.

201
Q

Why must the HCl be fresh in a flame test?

A

The acid can become contaminated with residue from previous tests, giving incorrect results.

202
Q

Why is the nichrome wire dipped into acid and then heated?

A

To remove residue of any previous sample tested.

203
Q

A compound heated in a Bunsen burner flame gave off a brown gas & caused a glowing splint to relight. Name that ion.

A

NO3^-

204
Q

Equation for the decomposition of NaNO3:

A

2NaNO3 –> 2NaNO2 + O2

205
Q

How can we compare the rate of decomposition of carbonates?

A

Use a delivery tube to bubble gas into limewater. Time how long it takes for limewater to turn cloudy.

206
Q

What must be added before BaCl2 (or BaNO3) in the test for sulfate ions?

A

HCl (or HNO3).

207
Q

Bromine disproportionates to a small extent, despite the reaction being thermodynamically not feasible, at RTP. How?

A

Disproportionation is an equilibrium system.

208
Q

Ionic equation for the disproportionation of chlorine with HOT NaOH(aq).

A

3Cl2 + 6OH- –> 5Cl- + ClO3^- + H2O

209
Q

Hydrogen chloride gas doesn’t conduct electricity as a gas. Hydrochloric acid conducts electricity. How?

A

The covalent bonding in HCl(g) changes to an ionic bond in aqueous solution.

210
Q

What is observed when HCl is added to solid Na2CO3?

A

Effervescence, and the solid dissolves to form a colourless solution.

211
Q

Watch out for the question phrasing.

A

Is it asking for a a name or a formula?

212
Q

Ionic equation for the reaction of chlorine with cold, aqueous hydroxide ions:

A

Cl2 + 2OH- –> Cl- + ClO- + H2O

213
Q

What type of reaction occurs when ammonia reacts with HCl?

A

Acid-base.

214
Q

All group 2 metals follow the same general equation for the reaction with oxygen:

A

2M(s) + O2(g) –> 2MO(s)
Except Sr & Ba also form peroxides: Ba(s) + O2 –> BaO2(s)

215
Q

Be + water

A

Beryllium doesn’t react with water.

216
Q

General equation for the reactions of group 2 elements with water:

A

M(s) + H2O(l) –> M(OH)2 (aq) + H2(g)

217
Q

Magnesium + water

A

With cold water, Mg reacts really slowly to form a weakly alkaline solution as Mg(OH)2 is only slightly soluble: Mg(s) + 2H2O(g) –> Mg(OH)2(aq) + H2(g)
With steam, it reacts vigorously according to the general equation for the reactions of group 2 elements with water.

218
Q

Group 2 metals react with chlorine gas to give…

A

… a metal chloride.

219
Q

All group 2 oxides are basic…

A

… except BeO, which is amphoteric. Oxide + water –> hydroxide. The solutions formed become more alkaline down the group.

220
Q

When group 2 oxides react with sulfuric acid, a sulfate + water is formed.

A

The insoluble sulfate at the surface, means the solid oxide beneath cannot react. This can be partially avoided by using a powdered oxide & stirring.

221
Q

Sulfates

A

Decrease in solubility down group 2.

222
Q

What do group 2 hydroxides form when they react with dilute acids?

A

Colourless solutions of metal salts.

Chemistry (23 decks)

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