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the underlying truth of chemistry explored by me > Topic 4 - inorganic > Flashcards

Topic 4 - inorganic Flashcards

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1
Q

Why are group 1 metals stored in oil?

A

When freshly cut they are quite shiny but they rapidly tarnish by reaction with oxygen to form a oxide layer which is why they are kept in oil

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2
Q

How do the flame test colours work?

A

Electrons are promoted to higher quantum levels via collisions with the high thermal kinetic energy particles. When the promoted electron falls back to its more stable electronic level, energy is emitted. If the wavelength of the photons emitted is in the visible region of the electromagnetic spectrum a ‘flame colour results’.

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3
Q

Explain the ionisation trend down groups 1 and 2?

A

As you go down the atomic radius increases. Therefore the outer electrons are further from the nucleus and also have more shielding. Therefore less energy is required to remove them

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4
Q

Explain the trends in successive ionisation energies in groups 1 and 2 ?

A

Successive ionisation energies always increase as the same nuclear charge is attracting fewer electrons and on average closer to the nucleus. The negative electrons are being successively removed from an increasingly more positive ion so more energy is required.

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5
Q

Explain the reactivity trend in groups 1 and 2?

A

The metal gets more reactive going down the group because it is easier to lose an electrons as the atomic radius increases and the shielding increases therefore the outer electrons are further away from the nucleus.

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6
Q

Explain the melting and boiling points down groups 1 and 2?

A

Decrease.This is because the ionic radii increase down the group increasing the charge separations between metal cations of the lattice and the free delocalised electrons. This weakens the attractive force and less energy is needed to break these bonds.

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7
Q

Explain the trends in electronegativity in group 1 and 2 elements?

A

The electronegativity decreases. They get lower because the effective nuclear attractive force on the outer electron charge decreases down the group.

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8
Q

Explain the reactions of group 1 metals with oxygen?

A

4M + O2 —> 2M2O

It’s a redox reaction and simple oxide is formed.

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9
Q

Write the equation how oxide are soluble in water?

A

The oxides are soluble in water and form a strongly alkaline hydroxide (except for lithium)

M20 + H2O —> 2MOH

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10
Q

Explain the trend in reactivity of group 2 metals with water?

A

Becomes more reactive as you go down the group.
Exceptions - beryllium has a strong resistant layer of oxide on its surface which lowers its reactivity at ordinary temperatures.

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11
Q

Summarise the reason for the increase in reactivity of group 2 metals with water?

A

The reactions becomes easier as the energy needed to form positive ions falls. This mainly due to a decrease in ionisation energy as you go down the group. This leads to lower activation energies and therefore faster reactions.

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12
Q

Explain the strength for group 1 hydroxides as you go down the group?

A

Strong base getting stronger down the group

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13
Q

Write the reaction for group 2 metals with water?

A

X + 2H2O —> X(OH)2 + H2

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14
Q

Write the equation for the reaction between Group 1 and water ?

A

2X + 2H2O —-> 2XOH + H2

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15
Q

Summarise the reason for the increase in the reactivity between group 1 elements and water ?

A

The reactions become easier as the energy needed to form positive ions falls. This is in part due to the decrease in ionisation energy as you go down the group and also in part to a fall in atomisation energy reflecting weaker metallic bonds as you go down from lithium to caesium. This leads to lower activation energies and therefore faster reactions.

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16
Q

Write the equation for Group 2 metals with oxygen?

A

2X + O2 —> 2XO

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17
Q

Why do some group 2 metals form peroxides on heating in oxygen?

A

Peroxide ion - O2 2-
The covalent bond between two oxygen atoms is relatively weak.
If you bring a small 2+ ion close to the peroxide ion. Electrons in the peroxide ion will be strongly attracted towards the positive ion. This forms a simple oxide if the positive ion is small and highly charged - it has a high charge density.
However as you go down the group the positive ions get bigger, they don’t have so much effect on the peroxide ion.

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18
Q

how do group 2 metals form nitrides on heating in air?

A

Even though nitrogen is thought to be fairly unreactive , all these metals combine with it to produce nitrides, X3N2 containing X 2+ and N 3- ions

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19
Q

Why are different oxides formed as you go down the group?

A
  • lithium form simple oxides
  • sodium form peroxide
  • potassium, rubidium and caesium form superoxide
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20
Q

Write the equation for the reaction of group 2 metals (not Be) with halogens? (A redox reaction)

A

M + Cl2 —> MCl2

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21
Q

What happens when magnesium reacts with steam?

A

It rapidly forms magnesium oxide and hydrogen gas in a vigorous reaction

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22
Q

What’s the test for carbon dioxide?

A

Limewater is used to test for carbon dioxide - it goes milky as a white precipitate forms.

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23
Q

Write the equation for the reaction of group 2 elements with dilute hydrochloric acid?

A

X + 2HCl ——> XCl2 + H2

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24
Q

What happens in a between group 2 elements and dilute HCl?

A

All the metals react with dilute hydrochloric acid to give bubbles of hydrogen and a colourless solution of metal chloride.
Reaction get more vigorous down the group.

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25
Q

Explain the reaction between dilute Sulphuric Acid and beryllium and magnesium?

A

Bubbles of hydrogen gas are formed, together with colourless solutions of beryllium or magnesium sulphate.

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26
Q

Explain the reaction between sulphuric acid and calcium, strontium and barium?

A

Calcium sulphate is sparingly soluble, and strontium and barium sulphate as being insoluble. That means that you’ll get a layer of insoluble sulphate on all these which slow down the reaction or stop it entirely.
In calcium’s case you’ll get a white precipitate and some hydrogen

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27
Q

State the trend in the solubility of hydroxides of group 2 elements?

A

The hydroxides become more soluble as you go down the group.

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28
Q

Explanation of the solubility trends in group 2?

A

Consider lattice enthalpy - ‘the energy released when 1 mole of an ionic lattice is formed from its constituent gaseous ions at standard pressure’.
The lattice enthalpy decreases down the group as the cation radius increases. Therefore energetically the solvation in terms of lattice energy is increasingly favoured down the group.

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29
Q

State the trends in solubility of sulphates in group 2?

A

The sulphates become less soluble as you go down the group

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30
Q

What reaction is used as a test for sulfate?

A

Add acidified barium chloride/dil, HCl or barium nitrate/dil nitric acid to a solution of the suspected sulphate. A dense white precipitate of barium sulphate forms in a positive result.

barium carbonate is insoluble as well as barium sulfate, which is why if we are testing for the presence of a sulfate ion in solution, we have to acidify it first before adding the barium chloride or barium nitrate solution to form a barium sulfate precipitate

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31
Q

State the solubility of the carbonates of group 2 elements?

A

The carbonates tend to become less soluble as you go down the group.

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32
Q

Explain the effect of heat on the group 2 carbonates?

A

All the carbonates in this group undergo thermal decomposition to give the metal oxide and carbon dioxide gas.

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33
Q

Write an equation for the thermal decomposition of group 2 carbonates?

A

XCO3 —-> XO + CO2

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34
Q

State the trend in the thermal stability of group 2 carbonates?

A

The carbonates become more stable to heat as you go down the group

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35
Q

What does the the thermal stability of group 1 and 2 carbonates depend on ?

A

The polarising effect or ability of the cation.

36
Q

What is polarising effect/power?

A

It is the ability to distort the electron cloud around the anion.
The greater the polarising effect on the anion the weaker the bonds in the anion become.

37
Q

State and explain what the polarising effect of a cation depends on and what?

A

Size and charge.
As the size (ionic radius) increases the polarising effect of a cation decreases.
As the charge increases the polarising power also increases.
Both of these increase the intensity of the electric field - a high charge density.

38
Q

Explain the trend in the thermal stability of group 2 nitrates?

A

Same argument as carbonates.
The small positive ions at the top of the group polarise the nitrate ions more than the larger positive ions at the bottom
More stable as you go down

39
Q

List out the flame test colours for group 1 compounds?

A 
Li-red
Na-yellow-orange
K- Lilac
Rb - red/purple
Cs - blue/violet
40
Q

State and explain the trends in electronegativity of group 7 elements?

A

Electronegativity falls as you go down the group. The atoms become less good at attracting bonding pairs of electron. Due to the nucleus is further away from the bonding pair therefore it isn’t as strongly attracted.

41
Q

State and explain the trends in boiling points in group 7?

A

The melting and boiling points increase as you go down the group.
As all the halogens exist as diatomic molecules. The intermolecular forces are London forces. As the molecules get bigger there are more electrons which can move around and set up temporary dipoles which create theses attraction.
The stronger the intermolecular forces the bigger energy you have to supply to break the bonds.

42
Q

Explain the solubility of group 7 elements in water?

A

Fluorine reacts violently with water to give hydrogen fluoride gas and a mixture of oxygen and ozone.
Chlorine,bromine and iodine all dissolve in water to some extend but no patterns.

43
Q

Why do you add cyclohexane in the displacement reaction of halogens

A

The halogens are much more soluble in organic solvents. Both hexane and the halogens are non-polar molecules attracted to each other by van der waals dispersion forces.

44
Q

State and explain the trend in oxidising ability of group 7 elements?

A

The oxidising ability falls as you go down the group. Whenever the halogens is involved in oxidising something in solution it ends up as halide ions with water molecules attached to them. As you go down the group the ease with which these hydrated ions are formed falls and so the halogens become less good as oxidising agents.

45
Q

How do you carry out a flame test?

A

-wear safety glasses and a lab coat. within a fume cupboard, light a bunsen burner.
-Use a nichrome wire (nichrome is an unreactive metal and will not give out any flame colour)
-Clean the wire by dipping in concentrated hydrochloric
acid and then heating in Bunsen flame. (hydrochloric acid is also used to convert any metal compound to a chloride)
-If the sample is not powdered then grind it up.
-Dip wire in solid and put it over a roaring blue Bunsen flame and observe flame

46
Q

flame test colours of Group 1 elements?

A 
Li - red
Na - yellow/orange
K - lilac
Rb - red/purple
Cs - blue/violet
47
Q

flame test colours of group 2 elements?

A 
Be - no colour
Mg - no colour
Ca - brick red
Sr - crimson red
Ba - apple green
48
Q

what’s the test for ammonium ions?

A
  • add sodium hydroxide and warm the mixture.
  • ammonia gas is releases.
  • use damp red litmus paper , which turns blue.
49
Q

What do group 1 nitrates decompose do?

A

lihtium nitrate:

  • oxide
  • nitrogen dioxide
  • oxygen

other group 1 nitrates:

  • nitrite (III)
  • oxygen
50
Q

state and explain trends in reactivity in group 7?

A

decrease down the group.

same reasons as the decreasing electronegativity down the group

51
Q

explain the trend in oxidising strength of group 7 elements?

A

the oxidising strength decreases down the group.

52
Q

halogens and halide displacement reactions. what displaces what?

displacement reactions are redox reactions.

A

a halogen that is a strong oxidising agent will displace a halogen that has a lower oxidisng power from one of its compounds.

  • chlorine displaces bromine and iodine
  • bromine displaces iodine only
  • iodine displaces neither

these reactions occur in aqeous solution

53
Q

observations of halogens in aqeous solution?

A

chlorine - colourless
bromine - orange/brown
iodine - red

54
Q

what happens and what are the observations when you had an organic solvent to halogen halide? why add an organic solvent?

A

halogens more soluble in organic solvents so it dissolves in the organic upper layer, where its colour can more easily be seen.

chlorine - 2 layers form with hexane,darker green layer
bromine - 2 layers orange colour bromine
iodine- 2 layers form
changes to purple or violet

55
Q

explain the reactivity of chlorine, bromine and iodine?

A

chlorine is the most reactive as it is the smallest atom so the incoming electron gets closer to and is more attracted by 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.

56
Q

colours of halogen in organic solvent?

A

chlorine - pale green
bromine - orange brown
iodine - violet

57
Q

explain the reaction of halogens with cold dilute alkali?

A
  • the colour of the halogen solution will fade to colourless.
58
Q

reaction of chlorine with water?

A

Cl2(g) + H2O(l) –> HClO(aq) + HCl (aq)
If some universal indicator is added to the solution it will
first turn red due to the acidity of both reaction products. It
will then turn colourless as the HClO bleaches the colour

59
Q

how is chlorine used in water treatment?

A

to kill bacteria. It has been used to treat drinking water and the water in
swimming pools. The benefits to health of water treatment by chlorine outweigh its toxic effects.

60
Q

reaction of chlorine with cold NaOH?

A

Cl2(aq) + 2NaOH(aq) —> NaCl (aq) + NaClO (aq) + H2O(l)

The mixture of NaCl and NaClO is used as Bleach and to disinfect/ kill bacteria

61
Q

reaction of chlorine with hot alkali?

A

With hot alkali disproportionation also occurs but the halogen that is oxidised goes to a higher
oxidation state.
3Cl2 (aq) + 6 NaOH(aq) –> 5 NaCl (aq) + NaClO3 (aq) + 3H2O (l)

62
Q

What is dispropotionation?

A

is the simultaneous oxidation and reduction of an element in a single reaction?

63
Q

What do group 2 nitrates decompose to?

A
  • metal oxide
  • nitrogen dioxide
  • oxygen
64
Q

state and explain the trends in the reducing power of halides?

A

increases down group 7.
A reducing agent donates electrons.
The reducing power of the halides increases down group 7.They have a greater tendency to donate electrons.This is because as the ions get bigger it is easier for the
outer electrons to be given away as the pull from the nucleuson them becomes smaller.

65
Q

explain the reactions of sodium fluoride and soidum chloride with concentrated sulfuric acid?
plus observations?

A

NaF(s) + H2SO4
(l) –>NaHSO4(s) + HF(g)
Observations: White steamy fumes of HF are evolved.
NaCl(s) + H2SO4(l) –> NaHSO4(s) + HCl(g)
Observations: White steamy fumes of HCl are evolved.

66
Q

sodium bromide + concentratd sulfuric acid? and observation?

A

Acid- base step: NaBr(s) + H2SO4(l) –> NaHSO4(s) + HBr(g)
Redox step: 2HBr + H2SO4 –> Br2(g) + SO2(g)+2H2O(l).
mistry fumes of HBr
brown fumes of Bromine
sulfur dioxide - colourless with choking smell

67
Q

sodium iodide + concetrated sulfuric acids?

A

I- ions are the strongest halide reducing agents. They can reduce the sulfur from +6 in H2SO4 to + 4 in SO2, to 0 in S and -2 in H2S.

Observations:
White steamy fumes of HI are evolved.
Black solid and purple fumes of Iodine are
also evolved.
A colourless, acidic gas SO2
A yellow solid of Sulphur
H2S (Hydrogen Sulphide), a gas with a bad egg
smell,
68
Q

what role does sulfuric acid play in the reaction with group 1 halides?

A

it plays the role of acid in the first step and then acts as an oxidising agent in the redox steps.

69
Q

explain the testing for halides in solution? what you would do not the observations.

A

reagent - silver nitrate solution. Nitric acid to make sure any iother anions are removed.
if precipitate is obtained then add ammonia solution either dilute or conc

70
Q

observations from the test for halide ions in solutions?

A

chloride ion
AgNO3 (solution) - white precipitate .
soluble in dilute aqueous ammonia.

bromide ion
AgNO3 (solution) - cream precipitate
soluble in CONC aqueous ammonia

iodide ions
AgNO3 (solution) - yellow precipitate
insoluble in both dilute and conc aqueous solution

71
Q

reactions of hydrogen halides with water?

A

the readily react with water to form acids all of which are colourless.

72
Q

reaction of hydrogen halides with ammonia?

A

all hydrogen halides react readily with ammonia gas to form salt(white solids)

73
Q

why can’t you use sulfuric acid to test for sulfate?

A

Sulfuric acid cannot be used to acidify the mixture because it contains sulfate ions which would form a precipitate.

74
Q

what are the test for negative ions?

A

add aqueous acid and observe effervescence.

bubble gas through limewater to test for CO2 will turn limewater cloudy.

75
Q

whats the test for positive ions?

A

Test for ammonium ion NH4+ , by reaction with cold
NaOH(aq) forming NH3
Ammonia gas can be identified by its
pungent smell or by turning damp red litmus paper blue.

NH4+ +OH- –> NH3 + H2O

76
Q

testing for the presence of a sulfate?

A

Acidified BaCl2 solution is used as a reagent to test for sulfate ions
If Barium Chloride is added to a solution that contains sulfate ions a white precipitate forms.
The acid is needed to react with carbonate impurities that are often found in salts which would form a white Barium carbonate precipitate and so give a false result

77
Q

describe what happens when ammonia and hydrogen bromide gases are mixed together?

A

they react to form a white solid

-ammonium bromide.

78
Q

explain what kind of reaction occurs when concentrated sulfuric acid is added to sodium fluoride?

A
  • sulfuric acid would act as an acid, not as an oxidising agent because fluoride ions have a very low reducing power and cannot reduce the sulfuric acid.
79
Q

Why is concentrated hydrochloric acid used in flame tests?

A

It is used to convert the metal compound being tested into a chloride, because chlorides are usually volatile and
so more likely to produce a better flame colour.

80
Q

suggest why it is not good idea to use water to put out a fire involving burning magnesium ?

A

-magnesium and water react at high temperatures to form hydrogen, which would make the fire worse.

81
Q

Explain fully why beryllium is less reactive than barium?

A

Beryllium is less reactive than barium because more energy is needed to remove the electrons in beryllium (than
in barium) during a reaction. This is because, even though the nuclear charge of beryllium is less than that of
barium (which decreases the ionisation energy), the electron being removed from beryllium is closer to the
nucleus than in barium. Also, there is only one filled shell of electrons in beryllium (compared with five filled
shells in barium) to repel the electron being lost. The last two factors outweigh the first, so the ionisation energy
of beryllium is higher than that of barium.

82
Q

Why cannot you use water to put out a fire in which magnesium metal is burning?

A

Mg + 2H2O —-> Mg(OH)2 + H2

-produces hydrogen gas. This hydrogen the burns by combining with oxygen gas in the air makes the fire hotter.

83
Q

Describe a standard test for chlorine gas?

A

-in a fume cupboard test the gas with moist litmus paper. Chlorine turns moist litmus paper red then bleaches it white.

84
Q

What is observed when iodine crystals are heated?

A

-the dark shiny great dagger like crystals sublime, producing violet fumes.

85
Q

Explain why hydrogen halides, as covalent compounds, dissolve in water to give solutions of low pH values?

A

-the hydrogen halides react in water and dissociate, so acting as proton donors (i.e acids)

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