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inorganic Flashcards

1
Q

trend in boiling points of METALS
in period three + reasons

A

increases
higher atomic charge
smaller atoms
more protons
more delocalised electrons per atoms
more energy to overcome

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2
Q

silicone melting point

A

high melting point
giant covalent lattice
very strong covalent bonds

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3
Q

phosphorous, sulphur, chlorine melting points

A

low melting points
weak van der waals forces
S>P>Cl, bigger molecules = more electrons = more van der waals

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4
Q

argon melting points

A

very low melting point
monatomic
very weak van der waals

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5
Q

reasons for atomic radius decreasing across period

A

more electrons and protons
same shell
equal shielding
greater attraction

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6
Q

reason electronegativity increases across a period

A

smaller radius
more protons
more attraction between nucleus and two bonding electrons in covalent bond

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7
Q

general change in ionisation energy across period 3

A

increases
more electrons and protons
no more shielding
greater attraction
more energy to overcome

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8
Q

reason for dip in ionisation energy from Mg to Al

A

Mg loses from 3s orbital
Al loses from 3p orbital
3p has higher energy than 3s

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9
Q

reason for dip in ionisation energy from P to S

A

P loses electron from orbital with one electron
S loses electron from orbital with two electrons
electron-electron repulsion in S

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10
Q

general trend in melting point of period three oxides

A

decreases

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11
Q

trend + reasons for melting point of first three period 3 oxides (ionic)

A

high (ionic structure)
MgO and Al2O3 highest due to higher charges making attraction stronger

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12
Q

reason for melting point of SiO2

A

giant covalent structure = high melting point
strong covalent bonds

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13
Q

trend and reasons for melting points of P4O10 and SO2

A

low (only van der waals forces)
simple covalent molecular structure
P4O10 higher due to being a larger molecule

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14
Q

equation for formation of sodium oxide

A

4Na + O2 -> 2Na2O

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15
Q

equation for formation of magnesium oxide

A

2Mg + O2 -> 2MgO

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16
Q

equation for formation of aluminium oxide

A

4Al + 3O2 -> 2Al2O3

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17
Q

equation for formation of silicon oxide

A

Si + O2 -> SiO2

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18
Q

equation for formation of phosphorus oxide

A

P4 + 5O2 -> P4O10

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19
Q

equation for formation of sulphur oxide (oxidation state of 5)

A

S + O2 -> SO2

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20
Q

equation for formation of sulphur oxide (equilibrium)

A

SO2 + O2 <-> SO3

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21
Q

reaction of basic oxides with water explanation

A

ions dissolve (some partially) and O2- (from compund) reacts with water
O2- + H2O -> 2OH-

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22
Q

reaction of sodium oxide with water

A

Na2O + H2O -> 2NaOH

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23
Q

reaction of magnesium oxide with water

A

MgO + H2O -> Mg(oh)2

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24
Q

why is aluminium oxide insoluble

A

high lattice enthalpy

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25
Q

why is silicone dioxide insoluble

A

giant covalent lattice

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26
Q

reaction of sodium oxide with acid

A

Na2O + 2H+ -> 2Na+ + H2O

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27
Q

reaction of magnesium oxide with acid

A

MgO + 2H+ -> Mg2+ + H2O

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28
Q

reaction of aluminium oxide with acid

A

Al2O3 + 6H+ -> 2Al3+ + 3H2O

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29
Q

reaction of aluminium oxide with base

A

Al2O3 + 2OH- -> 2Al(OH)4

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30
Q

reaction of silicone dioxide with base

A

SiO2 + 2OH- -> (SiO3)2- + H2O

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31
Q

equation of reaction of phosphorous oxide with water

A

P4O10 + 6H2O -> 4H3PO4
violent

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32
Q

reaction of acidic oxides with water

A

H2O molecules attack period 3 element
H+ ions released from water

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33
Q

reaction of sulphur dioxide with water

A

SO2 + H2O -> H2SO3

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34
Q

reaction of sulphur trioxide with water

A

SO3 + H2O -> H2SO4

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35
Q

reaction of phosphorous oxide with bases

A

P4O10 + 12OH- -> (4PO4)3- + 6H2O

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36
Q

reaction of silicone dioxide with bases

A

SO2 + 2OH- -> (SO3)2- + H2O

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37
Q

reaction of silicone trioxide with bases

A

SO3 + 2OH - -> (SO4)2- + H2O

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38
Q

name both basic group three oxides

A

sodium, magnesium

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39
Q

name intermediate group three oxide

A

aluminium

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40
Q

name all acidic oxides

A

silicone, phosphorous, silicone (tri and di)

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41
Q

magnesium and steam equation + observation

A

MgO + H2O -> MgO + H2
white flame

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42
Q

solubility of group two hydroxides down the group

A

solubility increases
conc of OH- increases (higher pH)

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43
Q

solubility of group two sulphates down the group

A

solubility decreases

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44
Q

group two reaction with oxygen observation

A

become more difficult to ignite
flame colour (Mg - Ba)
bright white -> red -> red -> green

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45
Q

group two reaction with water observation

A

become quicker down the group
magnesium only reacts with steam (vigorously)

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46
Q

strontium and barium reaction with sulphuric acid
observation

A

quickly stopped by formation of sulphate on top of metal

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47
Q

equation of reaction between group two metals and oxygen

A

O2 + 2M -> 2MO

48
Q

equation of reaction between group two metals and water

A

M + 2H2O -> M(OH)2 + H2

49
Q

magnesium and water reaction observation

A

extremely slowly
weakly alkaline
partially soluble

50
Q

reaction of group two metals with HCl equation

A

2HCl + M -> H2 + MCl2

51
Q

reaction of group two metals with H2SO4 equation

A

H2SO4 + M -> H2 + MSO4

52
Q

solubility of group two sulphates

A

all insoluble apart from berrilyium

53
Q

reaction of group two sulphates with acid not equation just trend

A

will always form water and carbon dioxide
remaining elements make a compound

54
Q

removal of acid rain

55
Q

used in x-rays

A

barium sulphate

56
Q

antacid

A

magnesium hydroxide

57
Q

halogens - oxidising or reducing agents?

A

oxidising agents

58
Q

power of halogens as oxidising agents moving down the group

A

decreases down the group
increasing atomic size

59
Q

halides - oxidising or reducing agents?

A

reducing agents

60
Q

power of halides as reducing agents moving down the group

A

increases
more shielding - easier to accept electrons

61
Q

test for halides

A

nitric acid - remove false positives
silver nitrate - precipitate - positive

62
Q

colours of halide precipitates

A

chloride - white
bromide - cream
iodide - yellow

63
Q

distinguishing between halide precipitates

A

dilute ammonia - only chloride redissolves
conc ammonia - bromde + chlodie redissolves
iodide never redissolves

64
Q

reaction with H2SO4 of sodium chloride

A

HCl gas - white fumes
H2SO4 + NaCl -> HCl + NaHSO4

65
Q

reaction with H2SO4 of sodium bromide
2x equations + observations

A

brown/red fumes
H2SO4 + NaBr -> HBr + NaH2SO4
2HBr + H2SO4 -> Br2 + SO2 + 2H2O

66
Q

reaction with H2SO4 of iodide

A

violet purple vapor
yellow solid
eggy smell
H2SO4 + Nl -> I + NaHSO4
2HI + H2SO4 ->I2 + SO2 + 2H2O
6HI + H2SO4 → 3I2 + S + 4H2O
8HI + H2SO4 → 4I2 + H2S + 4H2O

67
Q

chlorine and water no sunlight

A

Cl2 + H2O -> HCl + HClO

68
Q

chlorine and water in sunlight

A

2Cl2 + 2H2O -> 4HCl + O2

69
Q

chlorine and sodium hydroxide

A

Cl2 + 2NaOH -> NaCl + NaClO + H2O

70
Q

extraction of titanium with magnesium chloride

A

2Mg + TiCl4 -> 2MgCl2 + Ti

71
Q

definition of a transition metal

A

forms at least one ion with a part-full d-sub shell

72
Q

reason for ability to act as catalysts (transition metals)

A

easily transitions between oxidation states

73
Q

define ligand

A

particle with a lone pair of electrons that bonds to metals by a co-ordinate bond

74
Q

define complex

A

metal ion with co-ordinately bonded ligands

75
Q

co-ordination number

A

number of co-ordination bonds from ligands to metal ions

76
Q

lewis base

A

lone pair donor

77
Q

lewis acid

A

lone pair acceptor

78
Q

name of lewis acid/base in ion complex

A

ligands = lewis base
ion = lewis acid

79
Q

formation of NO in engines, how?

A

burning of N2 and O2 from air at high temperatures

80
Q

chelating effect

A

delta H negligible (same no. of smilier types of bonds)
delta S high positive (go from small no. of particles to high no. of particles)
delta G highly negative = very feasible

81
Q

when will small ligands substitute large molecules

A

when highly in excess

82
Q

why do large molecules generally sub small molecules

A

less ligands form
reaction more feasible

83
Q

plancks constant

A

6.63 x 10 ^-34

84
Q

why do transition metals have colour?

A

gap in energy between d orbital corresponds to energy of visible light

85
Q

why would an ion be blue?

A

absorbs all orange wavelengths, reflects all blue wavelengths

86
Q

reaction of chlorine with sodium hydroxide

A

2NaOH + Cl2 -> CaClO + NaCl + H2O

87
Q

reaction of chlorine in water (no sunlight)

A

H2O + Cl2 -> 2H+ + Cl- + ClO-

88
Q

reaction of chlorine in water (sunlight)

A

2H2O + Cl2 -> 4H+ + 2Cl- + O2

89
Q

pros of cleaning water with bleach

A

kills disease causing micro-organisms
reduces growth of algae
lasts a long time

90
Q

cons of cleaning water with bleach

A

could react with organic compound to form carcinogens
toxic + burns skin
respiratory problems

91
Q

why can’t you clean water with chlorine in sunlight

A

ClO- not formed
no diseases killed

92
Q

trend in oxidising agents of halogens

A

get less powerful moving down group

93
Q

how to test which halide ion you have

A

add samples of halogen gases
more reactive (more powerful halogens) will displace ions

94
Q

colour of Br2 gas

95
Q

colour of I2

96
Q

observations when reaction sodium chloride with sulphuric acid

A

white misty fumes

97
Q

observations when reaction sodium bromide with sulphuric acid

A

orange vapor (Br2)

98
Q

observations when reaction sodium iodide with sulphuric acid

A

yellow solid
rotten egg smell

99
Q

how to make test for halide ions certain

A

add nitric acid, remove false positives by reacting with other positive ions

100
Q

test for ammonium ions

A

add NaOH
turns litmus paper blue
NH4+ + OH- -> NH3 + H2O

101
Q

test for carbonates

A

add HCl
forms CO2 (cloudy limewater)
CO3 (2-) + 2H+ -> CO2 + H2O

102
Q

test for hydroxides

A

all alkali turns litmus paper blue
hydroxides will aswell
further tests to confirm

103
Q

test for sulphates

A

add HCl
ad barium chloride
white precipitation = positive

Ba (2+) + SO4 (2-) _> BaSO4

104
Q

define periodicity

A

repeating trends in period table

105
Q

amphoteric definition

A

can react as both an acid and a base

106
Q

role of H2SO4 in first reaction with halide ions

A

proton donor

107
Q

why is aluminium white

A

full electron shell in d-sub shell
cannot absorb visible light

109
Q

Define ligand

A

Particle with lone pair of electrons that can form a co-ordinate bond with a transition metal

110
Q

Co-ordination number

A

Number of co-ordinate bonds to one metal ion

111
Q

Lewis

112
Q

Which ions can form square planar complexes

A

Pt2+ and Ni2+

113
Q

Chelating effect

A

Delta H is negligible
Delta S increases as there is more particles formed
Delta G is very negative
Highly likely to be feasible

114
Q

Colours of Fe2+ ion
Water, NaOH, NH3, Na2CO3, conc HCl

A

Water - green sol
Add NaOH - green precip ([Fe(HO)4(OH)3]) darkens on oxidation
Add NH3 - green precip same as above
Add Na2CO3 - green precip
Add conc HCl - yello sol ([FeCl4]2-)

115
Q

Colours of Cu2+ ion
Water, NaOH, NH3, XS NH3, Na2CO3, HCl

A

Water - blue sol
Add NaOH - blue precip ([Cu(H2O)4(OH)2])
Add NH3 - blue precip same as above
XS NH3 - deep blue solution ([Cu(H2O)2(NH3)4]+)
Add Na2CO3 - green blue precip CuCO3
Add conc HCl - green solution [CuCl4]2-

116
Q

Colours of Fe3+ ion
Water, NaOH, NH3, Na2CO3, conc HCl

A

Water - purple but appears orange due to hydrolysis
NaOH - brown precip ([Fe(H2O)3(OH)3]
Add NH3 - brown precip same as above
Add Na2CO3 - brown precip and effervescence ([Fe(H2O)3(OH)3] + CO2
Add HCl - yellow sol [FeCl4-]

117
Q

Colours of Al3+ ion
Water, NaOH, XS NaOH, NH3, Na2CO3, conc HCl

A

Water - colorless solution
Add NaOH - white precip ([Al(H2O)3(OH)3])
XS NaOH - redissolves to colourless solution ([Al(H2O)2(OH)4]-)
Add NH3 - white precip same as above
Na2CO3 - white precip same as above and CO2 gas
Add conc HCl - [AlCl4]-

chemistry (8 decks)

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