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Chemistry Content > 4 - groups > Flashcards

4 - groups Flashcards

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1
Q

what do pure group 2 elements look like?

A

bright silvery solids
when exposed to air they combine with oxygen and form oxides as surface layers making them appear dull

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2
Q

what factors are considered when explaining trends in ionisation energy?

A
  • nuclear charge (number of protons)
  • orbital the electron is in
  • shielding
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3
Q

what happens to ionisation down group 2?

A

decreases
- nuclear charge increases
- as each quantum shell is added, energy of the outermost electron increases
- number of filled inner shells increases, force of repulsion of electron being removed increases

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4
Q

what happens to reactivity down group 2?

A

increases
- decrease in energy needed to remove the two electrons from each atom of the element

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5
Q

oxygen (from air) and magnesium visible reaction

A

bright flame and formation of a white solid

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6
Q

what are the reactions like between group 2 metals and air?

A

vigorous - may be hard to see if the metals are fresh samplpes
all group 2s need to be heated for the reaction to occur

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7
Q

what happen between air and group 2 metals if they aren’t heated first?

A

slow reaction forming an oxide which helps prevent further reaction

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8
Q

how is barium stored to prevent it reacting with oxygen and water vapour in the air?

A

in oil

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9
Q

general equation of group 2 and oxygen

A

2M(s) + O2(g) –> 2MO(s)

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10
Q

what are the reactions like between group 2 metals and chlorine?

A

vigorous (more down the group however harder to see the trend compared to the trend with oxygen)

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11
Q

general equation of group 2 and chlorine

A

M(s) + Cl2(g) –> MCl2(s)

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12
Q

what are the reactions like between group 2 metals and water?

A

very slow and does not proceed completely
(down the group vigour increases seen by more effervescence)

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13
Q

general equation of group 2 and water

A

M(s) + 2H2O(l) –> M(OH)2(aq) + H2(g)
*Ca(OH)2 is slightly soluble in water, so the liquid goes cloudy as a precipitate of Ca(OH)2 forms

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14
Q

what is the reaction between magnesium and steam like?

A

Mg reacts differently with steam and rapidly forms magnesium oxide (a white solid) and hydrogen gas in a vigorous reaction
Mg(s) + H2O(g) –> MgO(s) + H2(g)
INSERT IMAGE FROM TEXTBOOK??
hydrogen formed is burned as it leaves the tube - this is so the highly flammable gas doesn’t escape into the lab

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15
Q

beryllium and radium

A

not required to know any reactions however be able to predict - Be is less reactive than Mg and Ra is more reactive than Ba

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16
Q

what are group 2 oxides classed as?

A

basic oxides meaning they react with water to form alkalis - the reaction occurs when the oxide is added to water - observation: solids react to form colourless solutions

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17
Q

general equation of group 2 oxide and water

A

MO(s) + H2O(l) –> M(OH)2(aq)
it can be simplified since there is no change to the M^2+ ion
O^2- + H2O –> 2OH^-
shows hydroxide ions which is why resulting solutions are alkaline

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18
Q

what happens to solubility down group 2?

A

increases
therefore maximum alkalinity (pH value) of the solutions formed also increases down the group

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19
Q

what affects the pH value of alkaline solution formed between a group 2 oxide and water?

A

the relative amounts of oxide and water and the solubility of the hydroxide

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20
Q

what is limewater?

A

a saturated aqueous solution of calcium hydroxide

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21
Q

how do you test for carbon dioxide?

A

bubble through limewater - it goes cloudy (or milky) as a white precipitate forms
CO2 + Ca(OH)2 –> CaCO3 + H2O
as CO2 is bubbled through the amount of precipitate increases

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22
Q

what is milk of magnesia?

A

a suspension of magnesium hydroxide in water, a remedy for indigestion
a bottle contains a saturated solution of magnesium hydroxide mixed with extra solid magnesium hydroxide

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23
Q

how does milk of magnesia work?

A

neutralises some HCl in the stomach
Mg(OH)2 + 2HCl –> MgCl2 + 2H2O
OH- ions attack human tissue however the low solubility of Mg(OH)2 means that the concentration is low and does not pose a health risk

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24
Q

reaction of oxides and hydroxides of group 2 metals with acids

A

neutralisation reactions
the white solid reacts to form a colourless solution - reactions are exothermic
e.g.
MgO + H2SO4 –> MgSO4 + H2O
Ba(OH)2 + 2HCl –> BaCl2 + 2H2O

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25
Q

what is lime and how is it used in agriculture?

A

mostly calcium hydroxide
neutralises excess acidity in the soil
example reaction with nitric acid as the acid in soil:
Ca(OH)2 + 2HNO3 –> Ca(NO3)2 + 2H2O

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26
Q

what happens to solubility of group 2 sulphates down the group?

A

decreases
- Magnesium sulfate is classed as soluble
- Calcium sulfate is slightly soluble
- Strontium sulfate and barium sulfate are insoluble
the very low solubility of barium sulfate is used in a test for sulfate ions in solution

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27
Q

how to test for sulfate ion?

A

sulfate ions in an aqueous solution can be shown by adding a solution containing barium ions (usually barium chloride/nitrate)
any sulfate ions will react with the barium ions to form a white precipitate of barium sulfate
Ba^2+(aq) + SO4^2-(aq) –> BaSO4(s)

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28
Q

why must H+ ions be present when testing for sulfate ions?

A

other anions can form a white precipitate with barium ions - the H= prevent barium carbonate from forming as a white precipitate
dilute nitric acid or dilute hydrochloric acid is therefore added

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29
Q

method for testing for sulfate ions

A
  • add dilute nitric acid and barium nitrate solution
  • a white precipitate forms
    Ba(NO3)2(aq) + Na2SO4(aq) –> BaSO4(s) + 2NaNO3(aq)
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30
Q

what are barium meals?

A

barium ions are poisonous to humans
in hospitals patients are sometimes given a barium ‘meal’ containing barium sulfate (not poisonous as insoluble)
soft tissues will show up more clearly in an x-ray because of the dense white solid

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31
Q

what is meant by thermal stability?

A

how stable a compound is when it is heated
- doesn’t decompose = thermally stable
- decomposes as much as poss = not thermally stable

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32
Q

what happens when group 2 nitrates and carbonates are heated?

A

they do not melt, they decompose
- larger, complex nitrate ion can change into smaller, more stable nitrate ion or oxide ion by decomposing and releasing gas
- larger, complex carbonate ion can change to smaller more stable oxide ion by decomposing and releasing gas
- stabilities of nitrate and carbonate anions are influenced by charge and size of cations present (smaller and higher charge cations affect the anions more)

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33
Q

group 2 v group 1 nitrates and carbonates??

A
  • group 2 cation has double the charge
  • group 2 cation is smaller (ionic radii)
  • nitrate and carbonate anions are more complex than Cl- ion
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34
Q

which group 2 cation has the greatest influence on an anion?

A

Be2+
biggest charge, smallest size

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35
Q

what do the nitrates and carbonates of group 1 and 2 elements look like?

A

white solid

36
Q

what happens to group 1 and 2 nitrates or oxides when heated?

A

decompose to nitrates or oxides and give off nitrogen dioxide (brown fumes) and/or oxygen
(steam observed if nitrate contains water of crystallisation)
No brown fumes indicates a lesser decomposition
metal nitrate –> metal nitrite + oxygen
brown fumes indicates a greater decomposition
metal nitrate –> metal oxide + nitrogen dioxide + oxygen

37
Q

which group 1 and 2 nitrates when heated can you observe brown fumes?

A

Lithium nitrate, all group 2 nitrates (all greater decompositions)

38
Q

example reactions of group 1 and 2 nitrates decomposing

A

4LiNO3 —> 2Li2O + 4NO2 + O2 (only g1 nitrate to decompose completely)
2NaNO3 —> 2NaNO2 + O2 (all other g1 nitrates decompose this way)
2Be(NO3)2 –> 2BeO + 4NO2 + O2 (all g2 nitrates decompose tis way)

39
Q

what happens to group 1 and 2 carbonates when heated?

A

do not decompose or decompose to oxides and give off carbon dioxide
gas is colourless and carbonate and oxide are both white solids therefore no observations

40
Q

which group 1 and 2 carbonates when heated decompose?

A

Lithium carbonate and all group 2 carbonates
Li2CO3 –> Li2O + CO2
CaCO3 –> CaO + CO2 (typical g2 equation)

41
Q

what is a flame test and what does it indicate?

A

identifies the presence of a cation in a compound
it can indicate the presence of some metal cations in groups 1 and 2 but not all

42
Q

flame test method

A
  • wear goggles, within a fume cupboard, light a bunsen
  • using a dropper, add c HCl to the solid and mix so the metal compound begins to dissolve (HCl used to convert metals to metal chloride as chlorine is more volatile than other salts so likely to give better results)
  • dip a clean metal wire (nichrome or platinum) into mixture to obtain sample
  • hold end of wire in the flame and observe colour
43
Q

what are the problems with a flame test?

A
  • many compounds contain small amounts of sodium compounds as impurities, so the intense colour of sodium can mask other colours
  • describing colours with words is subjective - people have different levels of colour vision, and a word description of a colour may mean different colours to different people
44
Q

flame test: Li+

A

red

45
Q

flame test: Na+

A

yellow/orange

46
Q

flame test: K+

A

lilac

47
Q

flame test: Rb+

A

red/purple

48
Q

flame test: Cs+

A

blue/violet

49
Q

flame test: Be2+, Mg2+

A

no colour

50
Q

flame test: Ca2+

A

brick red

51
Q

flame test: Sr2+

A

crimson red

52
Q

flame test: Ba2+

A

apple green

53
Q

what causes the colours in flame tests?

A

electron transitions
electrons absorb energy and move to higher energy levels, an ‘excited state’
the movement is immediately followed by the return to the ‘ground state’ which releases energy
if the energy corresponds to radiation in the visible light spectrum then a characteristic colour appears
wavelength range in visible spectrum is 400-700 nm

54
Q

how to test for ammonium ions?

A

add sodium hydroxide solution and warm the mixture
NH4^+ + OH^- –> NH3 + H2O
the warming releases ammonia gas
ammonia can be recognised by its smell, but a simple chemical test is to use damp litmus paper which turns blue (ammonia is only common alkaline gas)
OR
hydrogen chloride gas reacts with ammonia to form white fumes of ammonium chloride
NH3 + HCl –> NH4Cl

55
Q

which elements from the halogens are often ignored?

A

fluorine and astatine
fluorine becuase it sometimes behaves differently and astatine because it only exists as radioactive isotopes

56
Q

what state are the halogens in at room temp?

A

F, Cl = gas
Br = liquid
I, At = solid

57
Q

what affects the melting and boiling points of the halogens?

A

diatomic molecules - mbpt depend on the strength of intermolecular forces of attraction (London forces) between these molecules

58
Q

what happens to the melting and boiling points down group 7?

A

increases

59
Q

explain the bonding of g7 diatomic molecules and how it affects melting and boiling points

A

halogen molecules are nonpolar - two atoms identical meaning pair of electrons shared equally
protons positive charge is permanent however the electron density in a halogen molecule continuously fluctuates - temporary dipole (instantaneous dipole)
you would expect no interaction between two halogen molecules as both are non-polar, if one becomes an instantaneous dipole it will cause an induced dipole = force of attraction between the two molecules = instantaneous dipole-induced dipole attraction
these weak forces increase as the number of electrons in the molecules increase - mbpt increases down group

60
Q

equations for halogen change of state

A

Br2(l) –> Br2(g)
I2(s) –> I2(g)

61
Q

what affects electronegativity?

A
  • nuclear charge - bigger charge = higher electronegativity
  • distance between nucleus and bonding pair of electrons - shorter = higher electronegativity
  • shielding in inner energy levels - fewer energy levels = higher electronegativity
62
Q

what happens to electronegativity down group 7?

A

decreases

63
Q

what happens to reactivity down group 7?

A

decreases
- high electronegativity
- nuclear charge - bigger charge = higher electronegativity
- distance between nucleus and bonding pair of electrons - shorter = higher electronegativity
- shielding in inner energy levels - fewer energy levels = higher electronegativity
- halogens act as oxidising agent and gain electrons to form negative ions

64
Q

key facts about g1/2 and g7 reactions

A
  • most vigorous between bottom of 1/2 and top of 7
  • products are salts - ionic solids that are usually white
  • all reactions involve electron transfer to the halogen (redox reactions)
  • oxidation of halogen goes from 0 to -1, oxidation of metal goes from 0 to +1 OR +2
65
Q

what can chlorine displace?

A

bromine and iodine

66
Q

what can iodine displace?

A

neither chlorine or bromine

67
Q

halogen displacement reaction info

A

occur in aqueous solution
indicated by a colour change
often done in an organic solvent (cyclohexane) as halogens are more soluble in cyclohexane than water, the halogen dissolves in the organic upper layer where its colour can be seen clearly

68
Q

why is chlorine more reactive than bromine and iodine?

A
  • smallest atom so incoming electron gets closer to and is more attracted by protons in the nucleus
  • smallest number of complete inner energy levels of electrons so incoming electron experiences the least repulsion
69
Q

chlorine and water reaction

A
  • disproportionation
  • no visible change
    Cl2 + H2O –> HCl + HClO
    HClO = chloric acid
    chlorine added to water to help disinfect for drinking purposes, kills pathogens such as cholera
70
Q

chlorine and a cold alkali

A

(cold dilute aqueous sodium hydroxide)
- disproportionation
Cl2 + 2NaOH –> NaCl + NaClO + H2O
sodium chlorate used as disinfectant, active ingredient in bleach

71
Q

chlorine with a hot alkali

A

(hot c. sodium hydroxide)
- disproportionation
3Cl2 + 6NaOH –> NaClO3 + 3H2O
sodium chlorate used in bleaching and as a weed killer

72
Q

what do halide ions act as? (reducing agent or oxidising agent)

A

reducing agent (reducing power increases down the group)
2X- –> X2 + 2e-
greater distance and shielding therefore more readily loses an e- to facilitate the reduction of another species.

73
Q

what do halogen molecules act as? (reducing or oxidising agent)

A

oxidising (decreases down the group)
greater distance and shielding therefore less readily gains an e- to facilitate the oxidation of another species.

74
Q

observation when c.H2SO4 is added to NaCl

A

misty fumes
products = HCl

75
Q

observation when c.H2SO4 is added to NaBr

A

misty fumes, brown fumes, colourless gas with choking smell
products = HBr, Br2, SO2

76
Q

observation when c.H2SO4 is added to NaI

A

misty fumes, purple fumes/black solid, colourless gas with choking smell, yellow solid, colourless gas with rotten egg smell
products = HI, I2, SO2, S, H2S

77
Q

sodium chloride and H2SO4 equation

A

NaCl + H2SO4 –> NaHSO4 + HCl

78
Q

sodium bromide and H2SO4 equations

A

NaBr + H2SO4 –> NaHSO4 + HBr
half equations:
2Br- –> Br2 + 2e-
H2SO4 + 2H+ + 2e- –> 2H2O + SO2
final equation:
2HBr + H2SO4 –> 2H2O + SO2 + 2Br

79
Q

sodium iodide and H2SO4 equations

A

NaI + H2SO4 –> NaHSO4 + HI
half equations:
2I- –> I2 + 2e-
H2SO4 + 6H- + 6e- –> 4H2O + S
equations:
2H+ + 2I- + H2SO4 –> So2 + 2H2 + I2
6H+ + 6I- + H2SO4 –> S + 4H2O + 3I2
8H+ + 8I- + H2SO4 –> H2S + 4I2 + 4H2O

80
Q

how to test for halide ions

A

silver nitrate solution (dilute nitric acid is added beforehand to make sure other anions are removed (e.g. carbonate))
chloride = white precipitate
bromide = cream precipitate
iodide = yellow precipitate

81
Q

what happens when dilute aqueous ammonia is added to a halide test?

A

chloride ions - soluble therefore solution turns colourless
bromide and iodide - insoluble

82
Q

what happens when concentrated aqueous ammonia is added to a halide test?

A

iodide ions - insoluble
chloride and bromide - soluble

83
Q

ionic equation of halide precipitates

A

Ag+(aq) + X-(aq) –> AgX(s)
specific example:
AgNO3(aq) + NaCL(aq) –> AgCl(s) + NaNO3(aq)

84
Q

equation of dissolution of silver chloride

A

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

85
Q

hydrogen halides acting as acids..

A

all are colourless gases and exist as polar diatomic molecules

86
Q

hydrogen halides and water

A

all react readily
e.g. HF + H2O <-> H3O+ + F-
hydrofluoric acid, hydrochloric acid, etc.
hydriodic acid NOT hydroiodic

87
Q

hydrogen halides and ammonia

A

all react with ammonia gas to form salts (all are white ionic solids)
e.g. NH3(g) + HCl(g) –> NH4Cl(s)

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