inorganic Flashcards
(136 cards)
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
2
Q
silicone melting point
A
high melting point
giant covalent lattice
very strong covalent bonds
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
4
Q
argon melting points
A
very low melting point
monatomic
very weak van der waals
5
Q
reasons for atomic radius decreasing across period
A
more electrons and protons
same shell
equal shielding
greater attraction
6
Q
reason electronegativity increases across a period
A
smaller radius
more protons
more attraction between nucleus and two bonding electrons in covalent bond
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
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
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
10
Q
general trend in melting point of period three oxides
A
decreases
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
12
Q
reason for melting point of SiO2
A
giant covalent structure = high melting point
strong covalent bonds
macromolecular
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
14
Q
equation for formation of sodium oxide
A
4Na + O2 -> 2Na2O
15
Q
equation for formation of magnesium oxide
A
2Mg + O2 -> 2MgO
16
Q
equation for formation of aluminium oxide
A
4Al + 3O2 -> 2Al2O3
17
Q
equation for formation of silicon oxide
A
Si + O2 -> SiO2
18
Q
equation for formation of phosphorus oxide
A
P4 + 5O2 -> P4O10
19
Q
equation for formation of sulphur oxide (oxidation state of 5)
A
S + O2 -> SO2
20
Q
equation for formation of sulphur oxide (equilibrium)
A
SO2 + O2 <-> SO3
21
Q
reaction of basic oxides with water explanation
A
ions dissolve (some partially) and O2- (from compund) reacts with water
O2- + H2O -> 2OH-
22
Q
reaction of sodium oxide with water
A
Na2O + H2O -> 2NaOH
23
Q
reaction of magnesium oxide with water
A
MgO + H2O -> Mg(oh)2
24
Q
why is aluminium oxide insoluble
A
high lattice enthalpy
25
why is silicone dioxide insoluble
giant covalent lattice
26
reaction of sodium oxide with acid
Na2O + 2H+ -> 2Na+ + H2O
27
reaction of magnesium oxide with acid
MgO + 2H+ -> Mg2+ + H2O
28
reaction of aluminium oxide with acid
Al2O3 + 6H+ -> 2Al3+ + 3H2O
29
reaction of aluminium oxide with base
Al2O3 + 2OH- -> 2Al(OH)4
30
reaction of silicone dioxide with base
SiO2 + 2OH- -> (SiO3)2- + H2O
31
equation of reaction of phosphorous oxide with water
P4O10 + 6H2O -> 4H3PO4
violent
32
reaction of acidic oxides with water
H2O molecules attack period 3 element
H+ ions released from water
33
reaction of sulphur dioxide with water
SO2 + H2O -> H2SO3
34
reaction of sulphur trioxide with water
SO3 + H2O -> H2SO4
35
reaction of phosphorous oxide with bases
P4O10 + 12OH- -> (4PO4)3- + 6H2O
36
reaction of silicone dioxide with bases
SO2 + 2OH- -> (SO3)2- + H2O
37
reaction of silicone trioxide with bases
SO3 + 2OH - -> (SO4)2- + H2O
38
name both basic group three oxides
sodium, magnesium
39
name intermediate group three oxide
aluminium
40
name all acidic oxides
silicone, phosphorous, silicone (tri and di)
41
magnesium and steam equation + observation
MgO + H2O -> MgO + H2
white flame
42
solubility of group two hydroxides down the group
solubility increases
conc of OH- increases (higher pH)
43
solubility of group two sulphates down the group
solubility decreases
44
group two combustion with oxygen observation
become more difficult to ignite
flame colour (Mg - Ba)
bright white -> red -> red -> green
45
group two reaction with water observation
become quicker down the group
magnesium only reacts with steam (vigorously)
46
strontium and barium reaction with sulphuric acid
observation
quickly stopped by formation of sulphate on top of metal
47
equation of reaction between group two metals and oxygen
O2 + 2M -> 2MO
48
equation of reaction between group two metals and water
M + 2H2O -> M(OH)2 + H2
49
magnesium and water reaction observation
extremely slowly
weakly alkaline
partially soluble
50
reaction of group two metals with HCl equation
2HCl + M -> H2 + MCl2
51
reaction of group two metals with H2SO4 equation
H2SO4 + M -> H2 + MSO4
52
solubility of group two sulphates
all insoluble apart from berrilyium
53
reaction of group two sulphates with acid not equation just products
will always form water and carbon dioxide
remaining elements make a compound
54
removal of acid rain
CaO/CaCO3
55
used in x-rays
barium sulphate
56
antacid
magnesium hydroxide
57
halogens - oxidising or reducing agents?
oxidising agents
58
power of halogens as oxidising agents moving down the group
decreases down the group
increasing atomic size
59
halides - oxidising or reducing agents?
reducing agents
60
power of halides as reducing agents moving down the group
increases
more shielding - easier to accept electrons
61
test for halides
nitric acid - remove false positives
silver nitrate - precipitate - positive
62
colours of halide precipitates
chloride - white
bromide - cream
iodide - yellow
63
distinguishing between halide precipitates
dilute ammonia - only chloride redissolves
conc ammonia - bromde + chlodie redissolves
iodide never redissolves
64
reaction with H2SO4 of sodium chloride
HCl gas - white fumes
H2SO4 + NaCl -> HCl + NaHSO4
65
reaction with H2SO4 of sodium bromide
2x equations + observations
brown/red fumes
H2SO4 + NaBr -> HBr + NaH2SO4
2HBr + H2SO4 -> Br2 + SO2 + 2H2O
66
reaction with H2SO4 of iodide
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
chlorine and water no sunlight
Cl2 + H2O -> HCl + HClO
68
chlorine and water in sunlight
2Cl2 + 2H2O -> 4HCl + O2
69
chlorine and sodium hydroxide
Cl2 + 2NaOH -> NaCl + NaClO + H2O
70
extraction of titanium with magnesium chloride
2Mg + TiCl4 -> 2MgCl2 + Ti
71
definition of a transition metal
forms at least one ion with a part-full d-sub shell
72
reason for ability to act as catalysts (transition metals)
easily transitions between oxidation states
73
define ligand
particle with a lone pair of electrons that bonds to metals by a co-ordinate bond
74
define complex
metal ion with co-ordinately bonded ligands
75
co-ordination number
number of co-ordination bonds from ligands to metal ions
76
lewis base
lone pair donor
77
lewis acid
lone pair acceptor
78
name of lewis acid/base in ion complex
ligands = lewis base
ion = lewis acid
79
formation of NO in engines, how?
burning of N2 and O2 from air at high temperatures
80
chelating effect
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
when will small ligands substitute large molecules
when highly in excess
82
why do large molecules generally sub small molecules
less ligands form
reaction more feasible
83
plancks constant
6.63 x 10 ^-34
84
why do transition metals have colour?
gap in energy between d orbital same as energy of visible light
85
why is carbon monoxide toxic
haemoglobin subs water ligands for carbon monoxide ligands
86
heterogenous catalyst
in a different state to reactants its catalysing
87
where does reaction occur on heterogenous catalyst
surface area of catalyst
88
homogenous catalyst
in same phase as reactants
89
why would an ion be blue?
absorbs all orange wavelengths, reflects all blue wavelengths
90
reaction of chlorine with sodium hydroxide
2NaOH + Cl2 -> NaClO + NaCl + H2O
91
reaction of chlorine in water (no sunlight)
H2O + Cl2 -> 2H+ + Cl- + ClO-
92
reaction of chlorine in water (sunlight)
2H2O + 2Cl2 -> 4H+ + 4Cl- + O2
93
pros of cleaning water with bleach
kills disease causing micro-organisms
reduces growth of algae
lasts a long time
94
cons of cleaning water with bleach
could react with organic compound to form carcinogens
toxic + burns skin
respiratory problems
95
why can't you clean water with chlorine in sunlight
ClO- not formed
no diseases killed
96
trend in oxidising agents of halogens
get less powerful moving down group
97
how to test which halide ion you have
add samples of halogen gases
more reactive (more powerful halogens) will displace ions
98
colour of Br2 gas
orange
99
colour of I2
brown
100
observations when reaction sodium chloride with sulphuric acid
white misty fumes
101
observations when reaction sodium bromide with sulphuric acid
orange vapor (Br2)
102
observations when reaction sodium iodide with sulphuric acid
yellow solid
rotten egg smell
103
how to make test for halide ions certain
add nitric acid, remove false positives by reacting with other positive ions
104
test for ammonium ions
add NaOH
turns litmus paper blue
NH4+ + OH- -> NH3 + H2O
105
test for carbonates
add HCl
forms CO2 (cloudy limewater)
CO3 (2-) + 2H+ -> CO2 + H2O
106
test for hydroxides
all alkali turns litmus paper blue
hydroxides will aswell
further tests to confirm
107
test for sulphates
add HCl
ad barium chloride
white precipitation = positive
Ba (2+) + SO4 (2-) _> BaSO4
108
define periodicity
repeating trends in period table
109
amphoteric definition
can react as both an acid and a base
110
role of H2SO4 in first reaction with halide ions
proton donor
111
why is aluminium white
full electron shell in d-sub shell
cannot absorb visible light
112
Define ligand
Particle with lone pair of electrons that can form a co-ordinate bond with a transition metal
113
Co-ordination number
Number of co-ordinate bonds to one metal ion
114
Lewis
Lone pair
115
Which ions can form square planar complexes
Pt2+ and Ni2+
116
Chelating effect
Delta H is negligible
Delta S increases as there is more particles formed
Delta G is very negative
Highly likely to be feasible
117
Colours of Fe2+ ion
Water, NaOH, NH3, Na2CO3, conc HCl
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-)
118
Trend in reactivity of group two
Increases
Easier to lose electrons down the group
119
Colours of Cu2+ ion
Water, NaOH, NH3, XS NH3, Na2CO3, HCl
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-
120
Colours of Fe3+ ion
Water, NaOH, NH3, Na2CO3, conc HCl
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-]
121
Colours of Al3+ ion
Water, NaOH, XS NaOH, NH3, Na2CO3, conc HCl
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]-
122
homogenous catalyst
catalyst in the same state as the reactants
123
heterogenous catalyst
catalyst in different state to the reactants
124
autocatalysis
when a catalyst is produced during a reaction
125
ways to maximise use of heterogenous catalysts
increasing SA
126
limits of heterogenous catalysts
stop working over time
tiny bits of impurities limit effectivity
127
how do heterogenous catalysts work
adsorption of reactants onto surface of catalyst
this takes place at an active site
lowers activation energy
reaction takes place and products leave
128
how could a heterogenous catalyst lower the activation energy
concentrate reactants (put them all in the same place)
weaken some of the bonds in reactants
position the molecule in a favourable orientation
129
why might a heterogenous catalyst not work
adsorption is too weak
adsorption too strong - molecules can't move
130
haber process reactions
3H2 + N2 -> 2NH3
catalysed by Fe
131
contact process reactions
2SO2 + O2 -> 2SO3
V2O5 catalyst
to make H2SO4
V2O5 + SO2 -> V2O4 + SO3
V2O4 + 1/2 O2 -> V2O5
132
how does a homogenous catalyst work?
usually forms an intermediate which then reacts further
very often aqueous
133
how does a transition metal homogenous catalyst work?
by variable oxidation states
134
reaction between iodine ions and peroxodisulphate
2I- + (S2O8)2- -> I2 + 2(SO4)2-
2Fe2+ + (S2O8)2- -> 2Fe3+ + 2(SO4)2-
2Fe3+ + 2I- -> 2Fe2+ + I2
very slow due to repel of two negative ions
catalysed by Fe 2+ reaction which can vary oxidation state
135
reaction between manganate ions and ethanediote ions
2MnO4- + 16H+ + 5 (C2O4)2- -> 2Mn 2+ + 8H2O + 10CO2
4Mn2+ + MnO4- + 8H+ -> 5Mn3+ + 4H2O
2Mn3+ + (C2O4)2- -> 2CO2 + 2MN2+
reaction is very slow because two negative ions repel each other
catalysed by Mn2+ ions which acts as a catalyst because it can easily change between oxidation states
the catalyst is a product of the reaction (autocatalysis)
reaction is slow until Mn2+ is made then it speeds up
136
what happens when a heterogenous catalyst becomes poisoned
impurities bind to the surface of the catalyst
blocks active site for reactants
