Inorganics Notes (Topic 4) Flashcards
so far: all inorganics notes from various sources but only pages 5-9 of chemrevise notes
1
Q
Chlorine dissolved in water/organic solvent colour
A
Pale green
2
Q
Bromine dissolved in water/ organic solvent colour
A
Orange/yellow
3
Q
In liquids that do not contain oxygen, iodine is
A
Purple
4
Q
In aqueous/alcoholic solution, iodine is
A
Brown
5
Q
Trend in oxidising agents of group 7
A
Oxidising power decrease down group because atoms get larger so weaker electrostatic attraction so harder to gain electrons (oxidising agent itself is reduced so must gain electrons)
6
Q
White phosphorus when heated with chlorine produces
A
Phosphorus trichloride
7
Q
How is HClO used for water treatment?
A
It contains the chlorate(I) ion which is strongly antibacterial so sterilises water
8
Q
Chlorine disproportionating in water ionic equation
A
Cl2 + H20 —> Cl- + 2H+ + ClO-
9
Q
What happens when bromine or iodine salts react with sulphuric acid?
A
The hydrogen halide given off is strong enough to be oxidised by the sulphuric acid so that little hydrogen halide is obtained.
With BROMIDE SALTS the bromide is oxidised to bromine while the sulfurique acid is reduced to sulfur dioxide (which is why hydrogen bromide + sulphuric acid —> bromine + sulphur dioxide + water)
don’t understand this - rating as 5 for now because I think my later flashcards cover this better
10
Q
How is the presence of HCL tested
A
Glass rod in concentrated ammonia, white smoke of ammonium chloride form
11
Q
Observation of iodide salt (any metal iodide) reacting with sulphuric acid?
A
Purple fumes, hydrogen sulfide(bad egg smell), Brown sludge
12
Q
What experiments prove the different reducing powers of halide ions?
A
Halide ions réactions with sulphuric acid and what the products are.
Cl- doesn’t reduce sulphuric acid
Br- reduces sulphuric acid to sulfur dioxide (ox no. from+6 in H2SO4 –>+4)
I- reduces sulphuric acid to hydrogen sulfide (ox no +6 in H2SO4 —-> -2)
(see reducing power INCREASES down the group because the ion reduces sulfuric acid to something with an even smaller oxidation state)
13
Q
Reactions of hydrogen halides with water general equation
A
HX(g) + H2O(L) —> H3O+(aq) + X-(aq)
14
Q
Reactions of hydrogen halides with ammonia
A
HX + NH3 —> NH4X
15
Q
Ionic equation for testing ammonium ions by adding sodium hydroxide
A
NH4+ + OH- —> NH3 + H20
Damp red litmus turns blue (because ammonia produced is alkali)
16
Q
Carbonate and hydrogencarbonate ions added to hydrochloride acid give off
A
Carbon dioxide and water
17
Q
Testing a suspected sulphate ion? And EQUATION
A
Add dilute HCL
Add barium chloride solution
SO4 2- (aq)+ Ba2+ (aq)—> BaSO4(s)
18
Q
Solubility of sulphates down group 2
A
Decreases
19
Q
Barium flame test colour
A
Apple green
20
Q
Strontium flame test colour
A
Deep Red
21
Q
Ionic equation to show reaction of bromine and silver nitrate?
A
Br- + Ag+ —> AgBr
22
Q
Which elements show no colour in a flame test and why?
A
Magnesium and beryllium because they require lots of energy to move electrons to a higher energy level, this cannot be provided by the Bunsen burner.
23
Q
Charge density
A
The quantity of charge related to the area. Eg. Something might have a high charge but be a massive ion so the charge is spread over a really large area and isn’t actually that polarising.
‘electric charge per unit area of surface’
24
Q
Elements which do not decompose at Bunsen burner temperatures and why?
A
Group 1 carbonates except from lithium.
Group 1 carbonates don’t have a large enough charge density to polarise the carbonate ion.
This is because they all have a 1+ charge which means they have even less polarising power than the group 2 elements which have a 2+ charge. This means the attraction between the cation (carbonate) and anion (oxygen) remains strong and so more energy is required to break it so it has high decomposition temperatures and is very thermally stable.
Lithium is small enough to have a polarising effect so lithium carbonate does decompose
25
Trend in first ionisation energy down group 2 and why?
Decreases
As the you down the group there are more shells, more inner electron shielding and a larger atomic radius. There is a weaker electrostatic attraction between the outer shell electron and the nucleus so less energy is required to remove it hence ionisation energy decreases down the group.
26
Trend in reactivity down group 2 elements and why?
Increases
Because less energy is required to remove the outer shell electron (due to the increased atomic radius, more shells and more inner electron shielding), it is easier to lose two electrons and form cations. The easier it can lose electrons the more reactive it is.
27
Why are group 2 elements in the same period as group 1 elements less reactive?
because they need to lose two electrons from their outer shell rather than one
28
Finish these equations:
1) Metal + oxygen —>
2) metal + water —>
3) magnesium + [cold] water —>
4) magnesium + [hot aka steam] water —>
5) metal + chlorine —>
6) metal oxide + acid —>
7) metal oxide + water —>
1) metal oxide
2) metal hydroxide + hydrogen
3) magnesium oxide + hydrogen
4) magnesium hydroxide + hydrogen
5) metal chloride
6) metal salt + water
7) metal hydroxide
29
What is the trend in solubility of hydroxides down group 2?
It increases!
30
Explain why iodine has a stronger boiling point than chlorine.
Larger molecule
More electrons
More instantaneous dipoles
More London forces
STRONGER London forces (because more of them)
More energy required to overcome attraction
31
Explain why chlorine is a better oxidising agent than iodine when oxidising iron?
Chlorine is a smaller atom than iodine
So there’s a stronger attraction between its nucleus and electron
So it gains electrons more easily
So can take more electrons from iron than iodine (because chlorine is an oxidising agent so is reduced so gains electrons which are it takes from iron)
32
Ionic equation for chlorine plus potassium iodide
CL2 (aq) + 2I-(aq) —> 2Cl- (aq) + I2 (aq)
33
ignore this card
ignore this flashcard
34
calcium flame test colour
brick-red
35
magnesium flame test colour
no colour (bunsen doesn't provide enough energy)
36
all group 2 oxides are...
white, ionic compounds
37
When barium reacts with oxygen, what does it forms?
substantial amounts of Ba2O2 (barium peroxide) and BaO (barium oxide)
38
What happens with the reaction as you react water with the group 2 metals, down the group?
it gets more vigorous
39
How does magnesium react with cold water vs steam?
Give the equation
slowly with cold water
rapidly with steam
Mg(s) + H20(l) ---> MgO(s) + H2(g)
40
How does calcium react with cold water?
Give the equation
quickly to give a milky white suspension of calcium hydroxide, some of which dissolves.
effervescence
Ca(s) + 2 H20(l) --> Ca(OH)2 (aq) + H2(g)
41
Give the general equation for how calcium, strontium and barium react with cold water.
X(s) + 2 H20(l) ---> X(OH)2(aq) + H2(g)
42
When barium reacts with cold water what is the observation?
a colourless solution of barium hydroxide (the most soluble of the group 2 hydroxides)
43
all group 2 metals, when heated with chlorine produce...
white, ionic chlorides
44
General equation for the reaction of magnesium, calcium, strontium and barium with chlorine.
X(s) + Cl2(g) ----heat----> XCl2(s)
45
Reactions of group 2 metal oxides (magnesium oxide, calcium oxide, strontium oxide and barium oxide) with water general equation
XO(s) + H2O(l) ---> X(OH)2 (aq)
| metal oxide + water --> metal hydroxide
46
solubility of magnesium hydroxide
sparingly soluble (soluble in small/restricted quantities)
47
Reactions of group 2 metal oxides with dilute acids (word equation)
metal oxide + dilute acid ---> metal salt (so either chloride, nitrate or sulfate) + water
48
Reactions of group 2 metal hydroxides with dilute acids (word equation)
metal hydroxide + dilute acid ---> metal salt (so either chloride, nitrate or sulfate) + water
49
All group 2 metal oxides are (acidic/alkali/basic)
| All group 2 metal hydroxides are (acidic/alkali/basic)
BASIC
| BASIC
50
Finish this example of a metal oxide and dilute acid:
MgO(s) + H2SO4(aq) -------------
------> MgSO4 (aq) + H20(l)
51
Finish this example of a metal oxide and dilute acid:
BaO(s) + 2HCl ---------
-------> BaCl2(aq) + H20(l)
52
Finish this example of a metal oxide and dilute acid:
CaO(s) + 2HNO3 -------------
----------> Ca(NO3)2 (aq) + H20(l)
53
Finish this example of a metal hydroxide and dilute acid:
Mg(OH)2 (aq) + H2SO4 (aq) ----------------
-------> MgSO4 (aq) + 2H20 (l)
54
How many moles of water are generally in
| metal hydroxide + dilute acid --> metal salt + water equations?
2 moles
2 H20(l)
55
How are the metal oxide + dilute acid --> metal salt + water equations usually balanced?
xxxx + 2HNO3 --->
| two moles of acid reactant except if the reactant is H2S04 which is one mole
56
What happens when group 2 oxides (so barium, strontium and calcium but not magnesium) react with sulfuric acid?
they produce a coating of the insoluble metal sulfate which prevents further reaction
57
Which group 2 metal oxides produce an insoluble coating when reacting with sulfuric acid?
calcium oxide
strontium oxide
barium oxide
58
equation in the edexcel textbook for magnesium oxide reacting with sulfuric acid
Normal: MgO(s) + H2SO4(aq) --> MgSO4(aq) + H2O(l)
textbook: MgO(s) + 2H+ (aq) ----> Mg2+ (aq) + H2O(l)
59
equation in the edexcel textbook for magnesium hydroxide reacting with sulfuric acid
Normal:
Mg(OH)2 (aq) + H2SO4 (aq) ----> MgSO4 (aq) + 2H20
textbook:
Mg(OH)2 (aq) + 2H+ (aq) -----> Mg2+ (aq) + 2H2O (l)
60
Talk about decomposition of group 2 carbonates
Group 2 carbonates including lithium decompose to m
metal oxide + carbon dioxide
61
Talk about decomposition of group 1 nitrates
Group 1 nitrates except lithium decompose to
metal nitrite + oxygen
62
Talk about decomposition of group 2 nitrates
Group 2 nitrates including lithium decompose more extensively to
metal oxide + nitrogen dioxide + oxygen
63
Give the equation for lithium nitrate decomposing
4LiNO3 ---> 2Li2O + 4NO2 + O2
64
what melts into a colourless solution and then resolidifies?
white nitrate solid
65
Trend in thermal stability down group 1 and 2?
thermal stability increases down the group
charge density decreases down the group
polarising power decreases down the group
C-O/N-O bond less distorted (bond in anion) and doesn't weaken as much
so coumpound more thermally stable
66
Why does charge density decrease down the group
because although nuclear charge is increasing, the size of the atom increases with it and cancels out its effects.
67
Why do group 2 nitrates and lithium nitrate decompose more extensively?
because they have a more polarising cation
68
what is useful about barium sulfate
its high insolubility makes it useful as a test for SO4 2- ions
69
Talk about limewater
limewater is saturated calcium hydroxide, when carbon dioxide passes through it, it forms insoluble calcium carbonate which is what causes the cloudy suspension
70
Flame formation
when substances are heated, the energy provided promotes electrons to a higher energy level. When the electrons fall back down to their original energy level, they release this energy in the form of light.
71
Which ions don't produce a flame?
Mg and Be ions - not enough energy from bunsen
72
Why might there be no colour to a flame?
bunsen burner doesn't supply enough heat energy to promote electrons
73
experimental procedures to show patterns in thermal decomposition of group 1 and 2 carbonates
Heat carbonate in test tube with a delivery tube extending into limewater
Record time taken for limewater to turn cloudy
Repeat with different carbonates
Limewater will take longer to go cloudy for more thermally stable carbonates
74
experimental procedures to show patterns in thermal decomposition of group 1 and 2 nitrates
no official test
| brown gas if group 2 nitrate being decomposed
75
Trend in melting point and boiling points of halogens as you go down the group?
increases
76
Explain the trend in melting point and boiling points of halogens as you go down the group?
As you go down the group
- the molecules become larger
- so there are more electrons
- and more instantaneous dipoles so there are more london forces (larger + stronger)
- more energy is required to break these intermolecular forces
77
Appearence at room temperature + extra stuff
- fluorine
- chlorine
- bromine
- iodine
- fluorine = pale yellow gas - highly reactive
- chlorine = green-yellow gas - reactive and poisonous in high concentrations
- bromine = brown volatile liquid - gives off dense brown/orange fumes
- iodine = shiny grey solid; purple gas
78
Chlorine appearance in aqueous solution
pale green (almost colourless)
79
Bromine appearance in aqueous solution
orange or yellow
80
Iodine appearance in aqueous/alcoholic solution
brown
81
Chlorine appearance in hydrocarbon solution
pale green (almost colourless)
82
Bromine appearance in hydrocarbon solution
red (edexcel tbook says orange or yellow so check)
83
Iodine appearance in hydrocarbon solution
violet
84
potassium iodide + aqueous iodine equation
final colour and name of product
KI(aq) + I2(aq) ---> KI3(aq)
brown solution of potassium triiodide
85
What is the trend in electronegativity down group 7?
It decreases
86
Explain the trend in electronegativity down group 7
As you go down the group, the molecules have a larger atomic radius, more shells, more inner electron shielding so the attraction between the nucleus and bonding pair of electrons that it wants to gain is much weaker and so they are less electronegative
87
What is the trend in reactivity down group 7?
it decreases
88
Explain the trend in reactivity down group 7
Reactivity decreases down the group because as you go down the group
-the atoms become larger
- there are more shells
-more inner electron shielding
-so they less easily attract and accept electrons (form 1- ions less easily) ///
so a weaker attraction between the nucleus and the (8th) electron they want to gain so they are less reactive. (Because easier to gain electrons means more reactive)
89
trend in oxidising agents down group 7
As you go down group 7, the strength of the oxidising agent decreases
90
explain the trend in oxidising agents down group 7
Oxidising agent strength decreases down the group because
oxidising agents themsleves are reduced (they gain electrons)
but as you go down group 7, due to the increasing atomic radius, inner electron shielding and shells, the atoms less easily attract and accept electrons
91
explain the trend in reducing agents down group 7
reducing agent strength increases down the group
reducing agents themselves are oxidised and lose electrons
and as you go down group 7 (due to larger atomic radius, more shells, more inner electron shielding) it becomes easier to lose electrons because of the weaker attraction
92
sodium and chlorine equation + nature of reaction
2Na + Cl2 --> 2NaCl
burns violently
93
reaction vigor down group 7 (reacting with metals)
and exothermicness...
reaction of group 1 and 2 metals with halogens become less vigorous down group 7 (because the halogens are less reactive)
less exothermic
94
metal + halogen -->
| eg. 2Na + Cl2 -------
metal halide
| ----> 2NaCl
95
group 2 metal + halogen example equation
Ca(s) + Cl2(g) ---> CaCl2(s)
96
Ion formed in reaction of iron with a more powerful oxidising agent (further up group)
forms Fe 3+
97
Ion formed in reaction of iron with a less powerful oxidising agent (further down group)
forms Fe 2+
98
iron (Fe3+) + chlorine reaction equation
2Fe(s) +3Cl2(g) --> 2FeCl3(s)
99
iron plus iodine equation
Fe(s) + I2(g) ---> FeI2(s)
100
halogen + metal (what is happening to the metal)
it is being oxidised
101
Br2(l) + 2Na(s) ---> 2NaBr(s)
reduction and oxidation equations for this (work them out)
2Na ---> 2Na+ + 2e-
| Br2 + 2e- ---> 2Br-
102
reactions of non-metals + halogens give...
covalent compounds
103
reactions of non-metals and halogens example
hydrogen and chlorine reacting
H2(g) + Cl2(g) --->2HCl(g)
104
Reactions of non-metals and halogens example (two equations)
white phosphorus
white phosphorus heated with chlorine:
P4 (s)+ 6Cl2 (g)—> 4PCl3(L)
the PCl product reacting with excess chlorine:
PCl3(l) + Cl2(g) ---> (equilibrium arrow) PCl5(g)
105
when halogens react with group 1 and 2 metals, the halogens are acting as
when this reaction occurs, what are the group 1 and 2 oxidised from and to?
oxidising agents
group 1 elements oxidised from 0 to +1
group 2 elements oxidised from 0 to +2
106
chlorine and water disproportionation reaction
-------------------------------------------------------------
what happens when universal indicator is added
-----------------------------------------------------------------
oxidation numbers of chlorine
------------------------------------------------------------------
HClO name
----------------------------------------------------------
colour of solution
--------------------------------------------------------------------
antibacterial properties
Cl2(g) + H2O(l) ---> (equilibrium arrow) HClO(aq) + HCl(aq)
--------------------------------------------------------------------
solution turns red because products are acidic but then colourless because HClO bleaches the colour
-----------------------------------------------------------------
O to +1 and -1
---------------------------------------------------------------
Hypochlorous acid/ chloric (I) acid
----------------------------------------------------------------
pale green solution
---------------------------------------------------------------
HClO is used in water treatment because it contains the chlorate (I) ion which is strongly bacterial so it sterilises water
107
Which equation produces a product used in water treatment
Cl2(aq) + H2O(l) -----> (equilibrium) HClO(aq) + HCl(aq)
108
chlorine and sodium hydroxide disproportionation equation (cold)
------------------------------------------------------------------
reactant
----------------------------------------------------------------
NaClO name
-------------------------------------------------------------
forms?
------------------------------------------------------------------
oxidation numbers of chlorine
Cl2(g) + 2NaOH(aq) ----> NaCl(aq) + NaClO(aq) + H2O(l)
---------------------------------------------------------------------------------
cold dilute aqueous sodium hydroxide (cold alkali)
----------------------------------------------------------------------------
sodium chlorate (I)
------------------------------------------------------------------------------
bleach
---------------------------------------------------------------------------
0 to -1 and +1
109
chlorine and sodium hydroxide disproportionation equation (hot)
------------------------------------------------------------------
reactant
----------------------------------------------------------------
NaClO3 name
-------------------------------------------------------------
forms?
---------------------------------------------------------------------
oxidation numbers of chlorine
Cl2(g) + 6NaOH(aq) ----> 5NaCl(aq) + NaClO3(aq) + 3H2O(l)
---------------------------------------------------------------------------------
hot concentrated aqueous sodium hydroxide (hot alkali)
----------------------------------------------------------------------------
sodium chlorate (v)
------------------------------------------------------------------------------
--------------------------------------------------------------------------
0 to -1 and +5
110
which reactions demonstrate the halide's increasing reducing strength down group 7?
the reactions of solid halides with sulfuric acid
111
[metal halides reacting with sulfuric acid]
Which ions cannot be oxidised by H2SO4, what does and what doesn't occur?
The H2SO4 is not strong enough an oxidising reagent to oxidise the chloride and fluoride ions.
No redox reactions occur.
Only acid-base reactions occur.
112
example reaction equation:
of metal(sodium) halide (chloride) reacting with sulfuric acid where the acid isn't strong enough to oxidise the halide ion
observation
NaCl(s) + H2SO4(l) -----> NaHSO4(s) + HCl(g)
Observations: White steamy fumes of HCl are evolved.
(redox NOT occurring, this is an acid-base equation)
113
example reaction equation:
of metal(sodium) halide (fluoride) reacting with sulfuric acid where the acid isn't strong enough to oxidise the halide ion
observation
NaF(s) + H2SO4(l) ----->NaHSO4(s) + HF(g)
Observations: White steamy fumes of HF are evolved.
(redox NOT occurring, this is an acid-base equation)
114
[metal halide + sulfuric acid]
Two equations to show the reducing ability of bromide ions. What are they called?
1) acid-base step
| 2) redox step
115
Acid- base step equation
(for bromide ions reacting with sulfuric acid as part of the metal halide + sulfuric acid sub topic)
-----------------------------------------------------------------------
observation
NaBr(s) + H2SO4(l) ----> NaHSO4(s) + HBr(g)
--------------------------------------------------------------------
white steamy fumes of HBr
116
Redox step equation
(for bromide ions reacting with sulfuric acid as part of the metal halide + sulfuric acid sub topic)
--------------------------------------------------------------------
observation
2HBr + H2SO4 ------> Br2(g) + SO2(g) + 2H2O(l)
----------------------------------------------------------------------
red fumes of bromine, colourless acidic SO2 gas
117
What do bromide ions reduce sulphuric acid to in the second step (redox step)
after the initial acid-base reaction, Bromide ions
reduce the
sulfur in H2SO4 from +6 to + 4 in SO2
118
bromide ions (in metal halide) reacting with sulfuric acid
Reduction half equation (it's the same doesn't matter what group 1 metal it is)
H2SO4 + 2 H+ + 2 e- --------> SO2 + 2 H2O
119
bromide ions (in metal halide) reacting with sulfuric acid
Oxidation half equation (its the same doesn't matter what group 1 metal it is)
2Br - ---------> Br2 + 2e-
120
strongest halide reducing agent
I- ions
121
What can I- ions, the strongest halide reducing agents, reduce when metal halides (iodide) react with sulfuric acid?
the sulfur from +6 in H2SO4 to
+ 4 in SO2
to 0 in S
to -2 in H2S.
122
I- ions reducing sulfur from +6 in H2SO4to + 4 in SO2 equation
----------------------------------------------------------
reduction half equation
2HI + H2SO4 ----> I2(s) + SO2(g) + 2H2O(l)
-------------------------------------------------------------------
H2SO4 + 2 H+ + 2 e- -------> SO2 + 2 H2O
123
I- ions reducing sulfur from +6 in H2SO4 to to 0 in S equation
----------------------------------------------------------------
reduction half equation
6HI + H2SO4 ----> 3 I2(s) + S (s) + 4 H2O (l)
--------------------------------------------------------------------
H2SO4 + 6 H+ + 6 e- --------> S + 4 H2O
124
I- ions reducing sulfur from +6 in H2SO4 to -2 in H2S
---------------------------------------------------------------------------------
reduction half equation
8HI + H2SO4 -----> 4I2(s) + H2S(g) + 4H2O(l)
------------------------------------------------------------------------
H2SO4 + 8 H+ + 8 e- --------> H2S + 4 H2O
125
iodide ions react with sulphuric acid and there are
_4 EQUATIONS_
What is the first? (the product then goes on to be the reactant for the other three)
NaI(s) + H2SO4(l) -----> NaHSO4(s) + HI(g)
126
In the metal halide (iodide) reacting with sulfuric acid reactions what are the observations for these products:
-Hydrogen iodide (HI)
-Iodine (I2)
-sulphur dioxide (SO2)
-Sulphur (S)
Hydrogen Sulphide (H2S)
-White steamy fumes of HI
-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
127
In the reaction of iodide ions (from metal iodide (halide)) with sulphuric acid, in the first equation sulfuric acid acts as the...........
and in the last three redox equations it acts as the..........
first equation H2SO4 acts as the acid
Last three - the oxidising agent
128
Across the 4 equations for when iodide ions in a metal halide reacts with sulphuric acid, the three important reduction products are
sulfur dioxide, sulfur and hydrogen sulfide
129
test used to identify which halide ions are present
test solution made acidic with nitric acid
silver nitrate solution added dropwise
precipitates then treated with ammonia solution
130
fluorides, chlorides, bromides, iodides
example ionic equation with chlorides
what does fluoride do?
fluoride forms no precipitate
Ag+ (aq) + Cl- (aq) --------> AgCl(s)
131
When testing for halide ions, why do we add nitric acid first?
-------------------------------------------
equation
to remove carbonate ions which would otherwise form a white precipitate of silver carbonate
----------------------------------------------------------------------
2 HNO3 + Na2CO3 ---> 2 NaNO3 + H2O + CO2
132
why are halide precipitates treated iwth ammonia solution?
give an example
to help differentiate between them if the colours are similar
eg. you can differentiate between silver chloride and silver bromide depending on if they dissolve in diluted or concentration solutions of ammonia
133
[part of the halide ion tests]
equation when silver chloride dissolves in XXXX ammonia
--------------------------------------
observation
AgCl(s) + 2NH3(aq) ----> [Ag(NH3)2]+ (aq) + Cl- (aq)
------------------------------------------------------------------------
colourless solution // DILUTE ammonia
134
[part of the halide ion tests]
equation when silver bromide dissolves in XXXX ammonia
------------------------------------------------
observation
AgBr(s) + 2NH3(aq) ----> [Ag(NH3)2]+ (aq) + Br - (aq)
------------------------------------------------------------------------
colourless solution // CONCENTRATED (not dilute) ammonia
135
[part of the halide ion tests]
equation when silver iodide dissolves in ammonia
SILVER IODIDE DOES NOT REACT WITH AMMONIA - IT IS TOO INSOLUBLE
136
After the halide ion test (nitric acid and silver nitrate), which two extra tests can be done to further confirm the result?
leave precipitate in sunlight
OR
add ammonia solution (either dilute or concentrated)
137
Chloride positive sunlight test
white precipitate turns purple - grey
138
Bromide positive sunlight test
cream precipitate turns green-yellow
139
Iodide positive sunlight test
no effect *sad emoji*
140
solubility of hydrogen halides in water? What does it dissolve to?
HCl (g)+ H2O(l) ----> H3O+ (aq)+ Cl- (aq)
a strongly acidic solution !!
141
hydrogen halides and water - liberation of energy idea?
the formation of the hydrated H3O- ion and the X- ion liberates enough energy for the compensating of the H-X bond breaking in the reactants.
HI is the stongest acid of the hydrogen halides because it has the weakest bond and is the easiest to break.
142
hydrogen halides reacting with ammonia:
general equation
----------------------------------------------------
observation
HX(g) + NH3(g) ---> NH4X(s)
---------------------------------------
white smoke of ammonium halide
143
Talk about silver halides and photography
Photochromic lenses contain silver chloride, which dissociates to form silver in the glass. Light source removed and silver recombines with the chlorine.
The precipitates (accept AgI) darken in sunlight. The reaction is used in photography.
144
other than the sunlight, ammonia and nitric acid & silver nitrate tests, what can be used to be identify silver halides?
hydrogen halide + ammonia ---> ammonium halide (white smoke forms)
145
The properties of fluorine and astatine can be deduced from patterns seen with chlorine, bromine and iodine...
fluorine - strongest oxidising agent but weakest reducing agents
astatine - weakest oxidising agent but strongest reducing agent
146
What is the strength of the F-F bond like?
weaker than Cl-Cl because it is so short that the lone pairs of electrons in the two atoms repel each other and weaken the bond
147
Fact about fluorine and oxidation states
fluorine usually gives rise to the highest oxidation number in the element it combines with (SF6 vs SCl4 , S has ox state of +6 vs +4)
148
astatine is...x...isotope?....
radioactive and has a short half life so only a few of its atoms exist at any one time. Longest lived isotope (astatine-211) has a half life of 8.3 hours!
149
how to test for the presence of carbonate and hydrogencarbonate (HCO3-) ions
- add dilute hydrochloric acid
- both give off carbon dioxide
- see if limewater turns cloudy (also observe effervescence when acid added)
150
testing presence of the carbonate ions
| IONIC equation
2H+ + CO3 2- ------> H2O + CO2
151
testing presence of the hydrogencarbonate ions IONIC equation
H+ + HCO3- -------> H2O + CO2
152
testing presence of the carbonate ions
| FULL equation
2HCl + Na2CO3 -----> 2NaCl + H2O + CO2
153
testing presence of the hydrogencarbonate ions FULL equation
HCl + NaHCO3 -------> NaCl + H2O + CO2
154
testing for the presence of sulfate ions + result
add dilute HCl
add barium chloride solution
presence of sulfate ions = white precipitate
155
When testing for the presence of sulfate ions, why do we add acid first?
it reacts with carbonate impurities which would otherwise react to form barium carbonate and a white precipitate
156
ionic equation for barium sulfate formation in the sulfate ion test
Ba2+ (aq) + SO42-(aq) ----> BaSO4(s).
157
ammonium ion test + results
warm X with sodium hydroxide solution
hold damp red litmus to mouth of tube
red litmus goes blue
158
ammonium ion test IONIC equation
NH4+ +OH- -----> NH3 + H2O
159
chlorine, bromine, iodine displacement of potassium halide reaction colours
What the colour of the test tube solution shows?
-three colours
1) The colour of the solution in the test tube shows which free halogen is present in solution.
- Chlorine =very pale green solution (often colourless),
- Bromine = yellow solution
- Iodine = brown solution (sometimes black solid present)
160
chlorine, bromine, iodine displacement of potassium halide reaction colours - when an ORGANIC SOLVENT is added
What the colour of the test tube solution shows?
-three colours
The colour of the organic solvent layer in the test tube
shows which free halogen is present in solution.
- Chlorine = colourless
- Bromine = yellow
- Iodine = purple
161
flame test procedure
Use a nichrome wire
Clean the wire by dipping in concentrated hydrochloric
acid and then heating in Bunsen flame
Dip wire in powdered solid and put in Bunsen flame and observe flame
162
why do we use a nichrome wire in flame tests?
nichrome is an unreactive metal
| and will not give out any flame colour
163
thermal decomposition
the use of heat to break down a reactant into more than one product
164
Which chlorine disproportionation reaction has an equilibrium arrow?
chlorine + water --> HClO + HCl
165
when F, Cl, Br and I react with sulphuric acid (in the reducing agent equations), what is the cation in the first equation for all of these reactions (before they go on to all the other equations)
Na
