Unit 3.3 - Chemistry of the p-block Flashcards
1
Q
Periodicity
A
How properties change across the periodic table
2
Q
Name 2 compound with more than 8 electrons in their outer shell
A
Phosphorus (V) chloride —> PCL5
Phosphorus (V) fluoride —> PF5
3
Q
What can we notice about the phosphorus compounds PCL5 and PF5?
A
There are 10 electrons around the central phosphorus atom
4
Q
What’s the name for having more than 8 electrons in an outer shell?
A
Expanding the octet
5
Q
Which elements are the only ones which can expand their octets?
A
Fluorine
Oxygen
Chlorine
6
Q
What can fluorine, oxygen and chlorine expand their octets with?
A
Highly electronegative elements
7
Q
Why are fluorine, oxygen and chlorine able to expand their octets?
A
They have large electronegativity values and have a very strong tendency to attract electrons to themselves
8
Q
What does Florine, oxygen and chlorine expanding their octets occur with?
A
Elements in the 3rd row of the periodic table (Na —> Ar)
9
Q
What does fluorine, oxygen and chlorine expanding their octets NOT occur with?
A
The elements in the 2nd row of the periodic table (Li—>Ne)
10
Q
Why are fluorine, oxygen and chlorine able to expand their octets with the elements in the 3rd row of the periodic table?
A
As the elements in the 3rd row can expand their octet of electron as vacant 3d orbitals of suitable energy are available for bonding, allowing them to hybridise
11
Q
How come Florine, oxygen and chlorine are able to expand their octet with elements from the 3rd row of the periodic table but not the second?
A
The energy difference between the 3s, 3p and 3d orbitals are small enough to overcome when bonding to strongly electronegative elements
With the 2nd row elements, the energy difference is too large to be overcome and no orbitals are available for expanding the octet
12
Q
What happens to the gap between subshells as you get further from the nucleus?
A
Decreases
13
Q
What does a smaller gap between subshells lead to?
A
Smaller energy differences
14
Q
What can happen to small energy differences?
A
Can be overcome when bonding to strongly electronegative elements
15
Q
Valency
A
Amount of bonds they can form
16
Q
Maximum valency of second row elements
A
4 pairs
17
Q
Maximum valency of 3rd row elements
A
Up to 7 pairs
18
Q
Octet expansion
A
The ability of some atoms to use d-orbitals to have more than 8 electrons in their valence shell
19
Q
Which orbitals are used to expand octets?
A
d-orbitals
20
Q
What is the maximum covalent bonds that can be formed by boron?
A
3 (this is electron deficient)
21
Q
What is the maximum covalent bonds that can be formed by carbon?
A
4
22
Q
What is the maximum covalent bonds that can be formed by nitrogen
A
3 (this is also has 1 lone pairs)
23
Q
What is the maximum covalent bonds that can be formed by oxygen?
A
2 (this is also has 2 lone pairs)
24
Q
What type of orbitals do elements in period 3 and below have and why is this important?
A
d orbitals
More bonds can be formed
25
How is PF5 formed?
An s-electron is promoted to a d-orbital
These then hybridise forming 5 sp^3d orbitals (each having the same energy)
26
Trends in bonding across the p-block
Ionic to covalent
27
Trend in bonding down the p-block
Bonding becomes more ionic (metallic)
28
Metallic character
The tendency of an element to lose electrons and form positive ions or cations
29
Trend in metallic character across the p-block
Metallic character decreases
30
Trend in metallic character down the p-block
Metallic character increases
31
What’s related to metallic character and how?
Electronegativity
As electronegativity increases, the metallic character decreases and vice versa
32
Where does amphoteric character occur mainly?
In the middle of the period table, here the elements Be, Zn, Al, Ga, In, Sn and Pb are found
33
Elements in group 2 which show amphoteric behaviour
Beryllium
34
Elements in group 3 which show amphoteric behaviour
Aluminium, gallium, indium
35
Elements in group 4 which show amphoteric behaviour
Tin, lead
36
Transition elements which show amphoteric behaviour
Zinc
37
Amphoteric behaviour
The oxide and hydroxide will react with acids and bases to form salts
38
What are oxides and hydroxides of group 1 and 2?
Basic
39
What does the fact that the oxides and hydroxides of group 1 and 2 are basic mean?
They only react as acids
40
How does an element behave if it’s demonstrating amphoteric behaviour?
As either an acid or a base
41
If something reacts with an acid, how does it behave?
As a base
42
If something reacts with a base, how does it behave?
As an acid
43
Describe the electronegativity values of elements that show amphoteric character
Intermediate
44
Describe the bonding of elements that show amphoteric character
It’s changing in character from ionic to covalent
45
Describe whether the elements that show amphoteric character are metals or non-metals
On the region of the periodic table where the elements are changing from being metals to non-metals
46
What are all of the elements which show amphoteric characters?
Metals with relatively high electronegativity
47
What type of bonding do elements that demonstrate amphoteric character have?
Ionic or covalent bonding
48
How does a hydroxide form in the amphoteric character equations?
By the addition of sodium hydroxide solution to a salt solution of the metal
49
Aluminium hydroxide showing basic behaviour equation
Al(OH)3 + 3H+ —> Al3+ + 3H2O
50
Aluminium hydroxide showing acidic behaviour equation
Al(OH)3 + OH- —> [Al(OH)4]-
51
Tetrahydroxe aluminate (III)
[Al(OH)4]-
52
Lead hydroxide showing basic behaviour equation
Pb(OH)2 + 2H+ —> Pb2+ + 2H2O
53
Lead hydroxide showing acidic behaviour equation
Pb(OH)2 + 2OH- +H2O —> [Pb(OH)4]2-
54
Tetrahydroxo plumbate (II)
[Pb(OH)4]2-
55
What happens if lead (II) hydroxide is reacted with nitric acid?
An aqueous solution of lead (II) nitrate is formed
56
What happens when lead (II) hydroxide is reacted with sodium hydroxide?
An aqueous solution of sodium plumbate is formed
57
What do we need in an question if we have a metal oxide and a base?
Water to balance it out
58
Equation for lead oxide showing basic behaviour
PbO + 2H+ —> Pb2+ + 2H2O
59
Equation for lead oxide showing acidic behaviour
PbO + 2OH- + H2O —> [Pb(OH)4]2-
60
Which elements react with sodium hydroxide solution violently to form a salt and release hydrogen gas?
Aluminium, zinc and beryllium
61
What do alumnimum, zinc and beryllium do when reacting with sodium hydroxide?
React violently to form a salt and release hydrogen gas
62
What is it that usually reacts with NaOH to release hydrogen gas?
Acids
63
What is an additional characteristic property of amphoteric behaviour?
Aluminium, zinc and beryllium reacting with sodium hydroxide solution violently to form a salt and release hydrogen gas
64
Equation for the reaction between aluminium and hydroxide
2Al + 2OH- + 6H2O —> 2[Al(OH)4]- + 3H2
65
Which elements does the violent reaction with sodium hydroxide NOT occur with?
Tin and lead
66
What is the only amphoteric group II element?
Beryllium
67
What is the only thing really that beryllium has in common with the other group II elements?
It has 2 electrons in the s-orbital
68
How is beryllium different to the other group 2 elements?
It’s oxide and hydroxide are amphoteric
69
How are many of beryllium’s compounds?
Covalent
70
Example of a covalent compound contains beryllium
BeCl2
71
Why are many of beryllium’s’ compounds covalent?
A consequence of its small size and comparatively large electronegativity which favours the formation of covalent bonds
72
Why does beryllium favour the formation of covalent bonds?
Small size and comparatively large electronegativity
73
What are the properties of beryllium closer to than the properties of other group II elements?
Closer to those of aluminium
74
Alkali
A soluble base
75
Silicon oxide (IV)
SiO2
76
Phosphorus pentoxide
P2O5
77
Describe the elements on the right hand side of the periodic table
Acidic
78
How can we investigate amphoteric behaviour in a practical?
By adding aqueous sodium hydroxide solutions to a number of aqueous salt solutions
79
Method for investigating amphoteric behaviour by adding aqueous sodium hydroxide solutions to a number of aqueous salt solutions
1.) place 4cm^3 of the zinc sulfate solution in a test tube
2.) to this solution, add sodium hydroxide solution drop wise until a precipitate forms
3.) divide the precipitate into 2 equal amounts
4.) to the first portion, add nitric acid slowly until it is in excess
5.) to the second portion, add sodium hydroxide solution slowly until it is in excess
6.) note all observations
7.) repeat the experiment with lead (II) nitrate and aluminium chloride
80
What do zinc, lead and aluminium all have in common?
All show amphoteric behaviours
81
What was the observation for adding the initial sodium hydroxide to zinc sulfate, lead (II) nitrate and aluminium chloride?
White precipitate
82
Why do we add the initial sodium hydroxide in the investigating amphoteric behaviour experiment?
To get the elements which show amphoteric behaviour (zinc, lead (II) and aluminium) to be as hydroxides
83
What’s the observation when adding both excess nitric acid and excess sodium hydroxide to the zinc, lead (II) and aluminium hydroxides?
Precipitate dissolves
Colourless solution
84
What word do we always use to describe a “clear” solution?
Colourless
85
What’s happened to a precipitate when it’s no longer visible?
Dissolved
86
Why do we add both excess nitric acid and excess sodium hydroxide in the investigating amphoteric behaviour experiment?
To see the hydroxides of zinc, lead (II) and aluminium react in the same way with both acids and bases = amphoteric behaviour
87
equation for the reaction between lead (II) ions and hydroxide ions
Pb ^2+ (aq) + 2OH- (aq) --> Pb(OH)2 (s)
88
equation for the reaction between lead (II) hydroxide and sodium hydroxide solution
Pb(OH)2 + 2OH- --> [Pb(OH)4]^2- (aq)
(sodium is a spectator ion)
89
equation for the reaction between aluminium ions and hydroxide ions
Al^3+ (aq) + 3OH- (aq) --> Al(OH)3 (s)
90
equation for the reaction between aluminium hydroxide and sodium hydroxide solution
Al(OH)3 + OH- (aq) --> [Al(OH)4]- (aq)
91
how do we identify complex ions?
placed in square brackets
92
what type of ion is one in square brackets?
a complex one
93
[Pb(OH)4]2-
tetrahydroxo plumbate (II)
94
[Al(OH)4]-
tetrahydroxo aluminate (III)
95
what do the roman numerals represent in the complex ions equations?
the oxidation state of the lead/aluminium
96
tetrahydroxo plumbate (II)
[Pb(OH)4]2-
97
tetrahydroxo aluminate (III)
[Al(OH)4]-
98
latinised name of lead
plumbum
99
latinised name of aluminium
alum
100
how do we explain why metals form amphoteric oxides and hydroxides?
their electronegativity values are intermediate (lie between a metal and a non-metal
101
what happens to an aluminium metal when exposed to oxygen?
forms an oxide spontaneously
102
why would there be no initial reaction when trying to react an aluminium metal with something?
aluminium forms an oxide spontaneously when exposed to oxygen --> need to get through the oxide layer first
103
what are the group II elements we will study the most?
boron and aluminium
104
boron electronic configuration
1s^22s^22p^1
105
aluminium electronic configuration
[Ne]3s^23p^1
106
what do both boron and aluminium have in common?
both have 3 electrons in their outer shell
107
oxidation state of boron and aluminium in their compounds
+3
108
general formula of group III halides
MX3
(x is the halide)
109
structure of boron (III) chloride or boron fluoride
trigonal planar
110
number of electrons around boron and aluminium
6
111
what can we see about boron and aluminium due to the fact that they have 6 electrons around them?
they're short of an octet = electron deficient
112
what does it mean if something is short of an octet?
is electron deficient
113
can boron and aluminium gain a pair of electrons readily?
yes
114
how come boron and aluminium can gan a pair of electrons readily?
they have strong electron acceptor properties
115
what are the electron acceptor properties of aluminium chloride reflected in?
the ready formation of the dimer Al2Cl6
116
what is Al2Cl6?
a dimer
117
what is the dimer Al2Cl6 made up of?
2 aluminium chloride molecules have combined
118
what type of bonds form to form an Al2Cl6 dimer?
co-ordinate bonds
119
how are co-ordinate bonds formed in the Al2Cl6 dimer?
by each aluminium atom accepting a lone pair of electrons from one of the chlorine atoms in the other molecule
120
why is the Al2Cl6 dimer able to form?
aluminium has a vacant 3d orbital
chlorine has a lone pair of electrons
121
is aL2cL6 electron deficient?
no
122
bonding in AlCl3
covalent
123
bonding in AlF3
ionic
124
bonding in Al2O3
ionic
125
properties of AlCl3 with its covalent bonding
sublimes at 180 degrees celcius (fairly low temperature)
decomposes in water
exists as dimers, Al2Cl6, when solid and as monomers
AlCl3 in the gaseous state
126
state of AlCl3
gaseous
127
when does AlCl3 exist as Al2Cl6??
when solid and as monomers
128
properties of AlF3 with its ionic bonding
melts at around 1300 degrees celcius
dissolves readily in water
molten alumnium fluoride is a good conductor of electricity
129
properties of Al2O3 with its ionic bonding
melts at 2100 degrees celcius
insoluble in water
conducts electricity when molten
130
which bonding is exhibited between aluminium and oxygen and fluorine and why?
ionic
large electronegativity
131
with which elements does aluminium exhibit ionic bonding and why?
with oxygen and fluorine
large electronegativity
132
what type of bonding does a large electronegativity usually lead to?
ionic
133
which type of bonding happens with less electronegative elements?
covalent
134
apart from with oxygen and fluorine, when else does aluminium exhibit ionic bonding?
when hydrated with water
135
why does ionic bonding occur in aluminium when hydrated with water?
large hydration energy owing to the very small size of the aluminium ion
136
list some aluminium elements that exhibit ionic bonding
AlF3
Al2O3
[Al(H2O)6]^3+
137
list some aluminium elements that exhibit covalent bonding
Al2Cl6
AlBr3
Al2H6
138
give an example of a donor-acceptor compound
ammonia-boron trifluoride
139
ammonia-boron trifluoride equation
NH3BF3
140
when is ammonia-boron trifluoride formed?
when ammonia reacts with boron trifloride, forming a co-ordinate bond
141
how are donor-acceptor compounds formed?
when one atom, having a lone pair of electrons to donate, reacts with another atom, which is electron deficient (i.e - has an empty orbital to accept the electron)
142
what type of bonding is in a donor-acceptor compound?
coordinate
143
what does each central atom have around it in a donor-acceptor compound?
an octet
144
why is ammonia-boron trifloride able to form?
nitrogen in ammonia: has a lone pair of electrons which is available for bonding
boron in boron trifluroride: has only 6 electrons in its outer shell
therefore, boron attains a full shell of electrons in forming a co-ordinate bond with the lone pair of electrons from the nitrogen
145
which group element is nitrogen?
group V
146
describe boron
electron deficient group III element
147
how many electrons does aluminium have in its outer shell in aluminium chloride and what does this mean?
6
is electron deficient
148
what happens when aluminium chloride reacts with a chloride ion?
aluminium will readily accept a lone pair of electrons from a chloride ion to form the complex ion tetrachloroaluminate (III)
149
[AlCl4]-
tetrachloroaluminate (III)
150
tetrachloroaluminate (III) ion
[AlCl4]-
151
equation for the reaction between aluminium chloride and chloride ions
AlCl3 + Cl- ⇌ AlCl4-
152
why will aluminium readily accept a lone pair of electrons from a chloride ion to form the complex ion tetrachloroaluminate (III)?
since aluminium in aluminium chloride only has 6 electrons in its outer shell (is electron deficient)
153
What leads to aluminium chloride being an industrially important catalyst?
the affinity of AlCl3 for chlorine species
154
what does the affinity of aluminium chloride for chlorine species lead to?
it's an industrially important catalyst
155
examples of industrially important catalysts that aluminium chloride forms
AlCl3 - catalysed Friedel-crafts reactions. there is no need to heat when carrying out reactions = saves energy
Low melting temperature ionic liquids, which contain the chloroaluminate (III) ion, [AlCl4]- are being developed as "clean technology" catalysts for the polymerisation of alkenes. the ionic liquid contains the chloroaluminate (III)) ion and a large organic cation.
156
why is boron nitride of interest to us?
as the B-N bond is similar to the C-C bond
157
which 2 words can be used to describe boron nitride and its comparisons to carbon?
isoelectronic (with the elemental forms of carbon)
isomorphism (occurs between the 2 species)
158
isoelectronic
the same number of electrons or the same electronic structure
159
isomorphism
similar chemical composition and exist in the same crystalline form
160
the same number of electrons or the same electronic structure
isoelectronic
161
similar chemical composition and exist in the same crystalline form
isomorphism
162
how many electrons do both C-C and B-N bonds have between the 2 atoms?
12
163
describe the structure of boron nitride
giant covalent
164
what are the giant covalent structures of boron nitride?
hexagonal
cubic
165
what is hexagonal boron nitride isomorphic with?
graphite
166
what is hexagonal boron nitride with graphite?
isomorphic
167
similarities between hexagonal boron nitride and graphite
high melting and boiling points
softness
insoluble
168
why do both hexagonal boron nitride and graphite have high melting and boiling points?
each atom forms 3 or 4 covalent bonds which require a lot of heat energy to overcome
169
why are both hexagonal boron nitride and graphite soft?
weak van der waal forces allow them to slide over each other
170
what does the feature that both graphite and hexagonal boron nitride are soft allow them to be used as?
a lubricant
171
how come hexagonal boron nitride and graphite can be used as a lubricant? explain this
they're both soft
weak van der waal forces allow them to slide over each other
172
why are both hexagonal boron nitride and graphite insoluble?
no charges particles to react with the permanent dipole of water molecules
173
differences between hexagonal boron nitride and graphite
hexagonal boron nitride is an electrical insulator, whereas graphite is an electrical conductor
atoms in adjacent layers in hexagonal boron nitride are in register. in graphite, they're out of register
electronegativity difference between boron and nitrogen = polar B-N bonds. no electronegativity difference between the carbon atoms in graphite (all atoms are the same).
174
how come the bonds between B and N in boron nitride are polar?
electronegativity difference
175
how come hexagonal boron nitride is an electrical insulator whilst graphite is an electrical conductor even though they're isomorphic?
hexagonal boron nitride only has localised electrons due to the large electronegativity difference between the nitrogen and boron atoms = no free electrons
in graphite, electrons are evenly delocalised (non-bonding electrons)
176
uses of hexagonal boron nitride
used in electronics as a substrate for semi-conductors
ceramics
microwave windows
catalyst carrier in fuel cells and batteries
single layers can be wrapped to create nanotubes
177
how are nanotubes created?
single layers of hexagonal boron nitride are wrapped
178
which property of hexagonal boron nitride allow them to be used in electronics as a substrate for semi-conductors?
electrical insulator
179
what can hexagonal boron nitride be used for due to it being an electrical insulator?
can be used in electronic as a substrate for semi-conductors
180
what is cubic boron nitride isomorphic with?
diamond
181
what is cubic boron nitride with diamond?
isomorphic
182
similarities between cubic boron nitride and diamond
hard
strong
extremely high melting and boiling points
insoluble
electrical insulator
Tetrahedral arrangement of atoms
183
why are both cubic boron nitride and diamond hard?
each atom is bonded to 4 others with a strong covalent bond = structure held together in a rigid 3D structure
184
why do both cubic boron nitride and diamond have extremely high melting and boiling points?
each atom has 3 or 4 covalent bonds which require a lot of heat energy to overcome
185
why are both cubic boron nitride and diamond electrical insulators?
no delocalised electrons in the structure to carry voltage
186
differences between cubic boron nitride and diamond
cubic boron nitride isn't as hard as diamond = preferred for grinding certian materials
cubic boron nitride is more stable due to the lack of electronegativity difference between the atoms
diamond can react with transition metals like iron and above 800 degrees celcius can react with air to form CO2
Cubic boron nitride is a good thermal conductor
187
why is cubic boron nitride more stable than diamond?
due to the lack of electronegativity difference between the atoms
188
what can diamond do above 800 degrees celcius?
can react with air to form CO2
189
uses of cubic boron nitride
wear-resistant coating
industrial abrasive
cutting tools
190
How many oxidation states do p-block elements usually show?
2
191
What does the higher oxidation state of p-block elements correspond to?
The group number
192
What is the lower oxidation state of p-block elements usually?
2 and lower
193
Oxidation state of group 3 elements
+3 and +1
194
Oxidation state of group 4 elements
+4 and +2
195
Oxidation states of group 5 elements
+5 and +3
196
Inert pair effect
The increasing reluctance, as you move down the group, of the s2 pair of electrons in the binding level to become involved in bonding
197
What does the inert pair effect cause in groups 3, 4 and 5?
The lower oxidation states of the element are more stable the lower down the group due to the inert pair effect
198
What happens to the metallic character of group 4 elements as the atomic number increases?
Increases
199
What happens to the bonding as the atomic number increases down group 4?
Covalent bonding
Metallic bonding
200
What happens as the atomic number increases down group 4 to the lower oxidation states?
Increase in stability
201
Which oxidation states increase in stability down group 4?
Lower oxidation states
202
Non-mets of group 4
Carbon and silicon
203
Semi-metal of group 4
Germanium
204
Metals of group 4
Tin and lead
205
Describe the electronic structure of group 4 elements
4 electrons in the outer shell of their atoms
2 in the s-sub-shell, 2 in the p sub-shell
206
2 allotropic forms of carbon
Diamond and graphite
207
What structure do both of the allotropes of carbon have (diamond and graphite)?
Giant covalent structure
208
Describe the structure, appearance, melting point and conductivity of silicon
Structure similar to that of diamond
Black crystalline solid
High melting point
Semi-conductor
209
Describe the appearance and melting point of lead and tin
Soft shiny metals
Fairly low melting points
210
Metallic character
The tendency of an element to lose electrons and form positive ions or cations
211
What can explain the change form non-metallic to metallic character as the atomic number increases down group 4?
Electronegativity
212
Why do the elements in group 4 change or non-metallic to metallic as the atomic number increases?
On going down the groups there is a decrease in electronegativity
213
What does a decrease in electronegativity down group 4 lead to?
A change from non-metallic to metallic character
214
Main oxidation states of carbon
+4
+2
215
Main oxidation state of silicon
+4
216
Main oxidation states of tin
+4
+2
217
Main oxidation states of lead
+4
+2
218
What happens to the stability of the +4 oxidation state down group 4?
Decreases
219
What happens to the stability of the +2 oxidation state down group 4?
Increases
220
Why does the stability of the +4 oxidation state decrease down group 4, and the +2 oxidation state increases in stability?
Due to the inert pair effect
221
What happens in terms of oxidation states down groups 3, 4 and 5?
Increasing tendency for the lower oxidation state to become more stable than the group valency
222
What’s the outer electronic structure of all of the elements in group 4?
ns^2npx^1npy^1
Where n varies from 2 (for carbon) and 6 (for lead)
223
What does the oxidation state of +4 imply in group 4?
Shows that all of the outer electrons are directly involved in bonding
224
What happens to the s^2 pair in group 4 as you go down the group?
There’s an increased tendency for it not to be involved in bonding
225
What does it lead to if there’s a decreased tendency for a pair of electrons not to be used in bonding?
Has a lower oxidation state
226
Explain the inert pair effect
As you get closer to the bottom of the group, there’s an increasing tendency for the s^2 pair not to be used in bonding. This leads to the element having a lower oxidation state —> the inert pair effect
227
What happens to the oxidation state when going from gallium to indium to thallium and why?
+1 state becomes more stable
Inert pair effect
228
What happens to the oxidation state when going from tin to lead and why?
+2 state becomes more stable
Inert pair effect
229
What happens to the oxidation state when going from arsenic, antimony to bismuth and why?
+3 state becomes more stable
Inert pair effect
230
Stability of CO2
Stable
231
Stability of PbO2
Strong oxidising agent
232
Equation for the reaction between lead oxide and hydrochloric acid
PbO2 + 4HCl —> PbCl2 + Cl2 + 2H2O
233
What type of bonding predominates in the +4 state of group 4 oxides?
Covalent bonding
234
When does covalent bonding predominate in group 4 oxides?
In the +4 state
235
Bonding of group 4 oxides in the +2 state
Tendency to be ionic
236
Proof that there’s a tendency to be ionic in the +2 state of group 4 oxides
+2 compounds of tin and lead contain the ions Sn^2+ and Pb^2+
e.g - SnCl2 and PbCl2
237
What happens to +4 compounds in group 4’s stability down the group and why?
Decreases down the group
Are covalent
The oxidising power of the state increases down the group
238
What happens to the stability of +2 compounds in group 4 down the group and why?
In the +2 stat, the compounds are more ionic than in the +4 state and their stability increases down the group (i.e - the reducing power of the state decreases)
239
Which group 4 oxides do we need to be aware of?
Carbon monoxide
Carbon dioxide
Lead (II) oxide
Lead (IV) oxide
240
Carbon monoxide:
Formula
Structure
Appearance/25°C
Redox behaviour
Acid-base behaviour
CO
Simple molecular
Colourless gas
Reducing agent
N/A
241
Carbon dioxide:
Formula
Structure
Appearance/25°C
Redox behaviour
Acid-base behaviour
CO2
Simple molecular
Colourless gas
Stable
Weak acid
242
Lead (II) oxide:
Formula
Structure
Appearance/25°C
Redox behaviour
Acid-base behaviour
PbO
Ionic lattice
Yellow solid
Stabl
Amphoteric
243
Lead (IV) oxide:
Formula
Structure
Appearance/25°C
Redox behaviour
Acid-base behaviour
PbO2
N/A
Dark brown solid
Strong oxidising agent
N/A
244
Why are CO and CO2 gas at room temperature?
Weak forces between separate molecules
245
Why are lead (II) oxide and lead (IV) oxide solid at room temperature?
Low ionisation energy = can form positive ions
Oxides are ionic with strong attractions between ions
Solids
246
Oxides of Carbon to learn
Carbon monoxide
Carbon dioxide
247
What is carbon monoxide?
A very poisonous gas
248
Why is carbon monoxide so dangerous?
Reacts in preference to oxygen with haemoglobin in the blood
249
What is carbon monoxide send as?
A reducing agent industrially - extraction of iron in the blast furnace (reduction of iron (III) oxide to iron)
250
What happens to carbon monoxide during the extraction of iron in the blast furnace?
Is oxidised to CO2
251
Reaction in the blast furnace
Fe2O3 + 3CO —> 2Fe + 3CO2
252
What is carbon monoxide easily oxidised from and to and why is this good?
+2 to +4
+4 is more stable
253
Describe CO2 in water
Fairly soluble
254
Is CO2 acidic? Explain
Weakly acidic
Is partially hydrolysed by water -> about 1% of the molecules
255
HCO3- (aq)
Carbonic acid
256
Equation for the reaction between CO2 and water
CO2 (g) + H2O (l) ⇌ H+ (aq) + HCO3 - (aq)
257
What does CO2 react with alkalis for?
To form salts
258
How does CO2 form salts?
Reacts with alkalis
259
What type of salts does CO2 react with alkalis to form?
Carbonates
Hydrogen-carbonates
260
CO2 reaction with sodium hydroxide
CO2 + NaOH —> NaHCO3
261
CO2 reaction with 2 sodium hydroxides (concentrated)
CO2 + 2NaOH —> Na2CO3 + H2O
262
What do carbonates and hydrogencarboantes react with acids to produce?
CO2
263
When do carbonates and hydrogencarbonates release CO2?
When reacting with acids
264
What does CO2 do to limewater?
Turns milky
265
Lime water
Calcium hydroxide
266
What turns limewater milky?
CO2
267
Which reaction is used to test for acids?
When carbonates and hydrogencarbonates react with acids, it releases CO2
The CO2 turns limewater milky
268
Equation for the reaction between Ethanoic acid and sodium hydrogencarbonate
CH3COOH + NaHCO3 —> CH3COONa + CO2 + H2O
269
Oxidation state of carbon dioxide
+4
270
What is the most stable, carbon monoxide or carbon dioxide?
Carbon dioxide
271
Lead (II) oxide formula
PbO
272
Lead (II) oxide at room temperature
Stable
273
Lead (II) oxide description
Ionic yellow solid
274
Lead (I) oxide in water
Insoluble
275
Oxidation state of lead (II) oxide
+2
276
Reducing or oxidising properties of lead (II) oxide
None - stable
277
Acid-base behaviour of lead (II) oxide
Amphoteric
(Forms salts with acids and alkalis)
278
Equation for the reaction between lead (II) oxide and an acid
PbO + 2H+ —> Pb2+ + H2O
279
Equation for the reaction between lead (II) oxide and an alkali?
PbO + 2OH- + H2O —> [Pb(OH)4]^2-
280
Description of lead (IV) oxide
Dark brown solid
281
Lead (IV) oxide formula
PbO2
282
Redox behaviour of PbO2
Powerful oxidising agent
283
Example of PbO2 being a powerful oxidising agent
Will oxidise chloride ions to chlorine. The lead (IV) is reduced to lead (II) chloride which is found as white precipitate.
284
How is lead (II) chloride found?
White precipitate
285
What has to be done to the reagents in the reaction between PbO2 and chloride ions?
Have to be heated
286
Equation for the reaction between lad (IV) oxide and hydrochloric acid
PbO2 + 4HCl —> PbCl2 + Cl2 + 2H2O
287
Oxidation state of PbO2
+4
288
What is lead easily reduced from and to in PbO2 and why?
From +4 to +2
+2 is more stable
289
Method used to identify Pb2+ ions and results
Pb2+ solution and I- ions (usually KI is used)
= bright yellow precipitate of lead (II) iodide
290
Equation for the reaction between lead ions and iodide ions
Pb2+ (aq) + 2I- (aq) —> PbI2 (s)
291
Group 4 halides to remember
Tetrachloromethane
Silicon (IV) chloride
Lead (I) chloride
292
Tetrachloromethane:
Formula
Bonding
Structure
Appearance
Reaction with water
CCl4
Covalent
Simple molecular
Colourless liquid
No reaction - forms a separate layer
293
Silicon (IV) chloride:
Formula
Bonding
Structure
Appearance
Reaction with water
SiCl4
Covalent
Simple molecular
Colourless liquid
Violent reaction
294
Lead (II) chloride:
Formula
Bonding
Structure
Appearance
Reaction with water
PbCl2
Ionic
Ionic lattice
White solid
Insoluble - cold water
Soluble- hot water
295
How com Tetrachloromethane has a simple molecular structure?
Covalent
Low boiling point
296
Why can’t Tetrachloromethane react with water?
Carbon cannot expand is outet to form a coordinate bond with a lone pair of electrons from the water molecule
Carbon has no empty d levels of suitable energy available for bonding
297
How come silicon (IV) chloride is able to react violently with water?
Has empty d levels of suitable energy available and the tetra slide reacts with water violently, forming co-ordinate bonds
298
Reaction between silicon (IV) chloride and water equation
SiCl4 + 2H2O —> SiO2 + 4HCl
299
What does SiO2 (product of the reaction between solid ion (IV) chloride and water) appear as?
White precipitate
300
How does HCl show up after the reaction between silicon (IV) chloride and water?
White fumes
301
Structure of silicon (IV) chloride
Tetrahedral (like all chlorides)
302
Appearance of lead (II) chloride
White ionic solid
303
Describe the solubility of lead (II) chloride in water
Cold water = insoluble
Hot water = partially soluble
304
Equation for the reaction between lead (II) chloride and water
H2O
PbCl2 (s) ——> Pb^2+ (aq) + 2Cl- (aq)
305
Natural source of silicon
Sand (silicon dioxide)
306
Describe lead as a metal
Dense
307
What is lead used in?
Roofing
308
What does a +2 oxidation state imply?
2 electrons used in bonding, 2 not
309
What does something being stable imply?
Unreactive
310
What does something being unstable imply?
Reactive
311
CO2 reacting with dilute NaOH
NaHCO3
312
CO2 reacting with concentrated NaOH
Na2CO3
313
How do we know if NaOH is concentrated?
“2” in front
314
Why do carbonates and hydrogencarbonates do in the body?
Control blood pH
315
Under which conditions alone can we extract elements by reduction?
With elements less reactive than carbon
316
What must we do to an equation if we have anything reacting as an acid?
Must include water to balance it out
317
When must we include water in these equations in this unit?
If we have anything reacting as an acid
318
What must we back any answers up with?
Equations
319
How come silicon (IV) chloride has a violent reaction with water but Tetrachloromethane doesn’t?
Silicon is a bigger atom than carbon
Can surround more atoms around it
320
What are group 7 elements otherwise known as?
The halogens
321
How are the halogens always written?
Diatomically
322
Are the halogens metals or non-metals?
Non-metals
323
Are the halogens similar to one another?
Yes - they show gradual changes only with increasing atomic mass
324
What type of bonds do the halogens form with metals?
Ionic
325
What type of bonds do the halogens form with non-metals?
Covalent bonds
326
Structure of halogens
Simple molecular
327
Change in atomic radius down the group of halogens + explanation
Atomic radius increases down the group
More shells present
328
Describe the melting + boiling points of the halogens + explain this
Low melting and boiling points
Diatomic molecules and a simple molecular structure
329
What happens to the melting and boiling points of the halogens down the group? Explain this
Increases
On going down the group, the size increases with a subsequent increase in the Van der Waal’s forces (induced dipole-induced dipole)
It’s these forces that are broken when the halogens are changing state
330
Type of Van der Waal forces in halogens + why
Induced dipole-induced dipole
Non-polar covalent molecules
331
Explain the formation of the induced-dipole induced-dipole Van der Waal forces
Slight uneven distribution of electrons = temporary dipole
Causes a chain of induced dipole-induced dipole bonds
332
In which halogen are the induced-dipole induce-dipole bonds the most significant and why?
Iodine
Largest molecule = most electrons
333
States of the different halogens + explanation
F2, Cl2 = gases
Br2 = liquid
I2 = solid
On going down the group, the Van der Waal forces increase, increasing the melting points down the group
334
What happens to electronegativity down the group of halogens?
Decreases
335
What happens down group 7 when the electronegativity decreases down the group?
Decrease in reactivity from F to I
336
What happens to the electrode potentials values down group 7?
The positive value of the electrode potential decreases
337
What does the decrease in the positive value of the electrode potentials down group 7 reflect?
The decrease in reactivity and oxidising power of the elements down the group
338
What are all of the halogens?
Oxidising agents
339
What’s different about the oxidising properties of the halogens?
Have different strengths as oxidising agents
340
What happens to the oxidising power of the halogens down group 7?
Decreases
341
How can we measure the halogens ability to remove electrons from other species?
Measure using the standard electrode potentials for the halogens
342
Most powerful oxidising agent of group 7
Cl2
343
Most powerful reducing agent of group 7
I-
344
In the reaction below, what is the oxidising agent and what is the reducing agent?
Cl2 (g) + 2e- —> 2Cl- (aq)
<—
Cl2 - oxidising agent
Cl- - reducing agent
345
Chlorine physical appearance at room temperature
Green gas
346
Bromine physical appearance at room temperature
Orange liquid
347
Iodine physical appearance at room temperature
Black/purple solid
348
Which halogen is a green gas at room temperature?
Chlorine
349
Which halogen is an orange liquid at room temperature?
Bromine
350
Which halogen is a black/purple solid at room temperature?
Iodine
351
What can all of the halogens do at room temperature?
Produce gases
352
What can bromine easily do at room temperature and what word describes this?
Easily vaporises at room temperature
Volatile
353
What can iodine do at room temperature and what does this mean?
Sublime
Straight from solid to gas
354
Which halogen can sublime at room temperature?
Iodine
355
What halogen is volatile (easily vaporises at room temperature)?
Bromine
356
What type of reaction is the reaction which occurs between a halogen and a halide ion?
A redox reaction
357
How are redox reactions explained?
Using electrode potentials
358
Electrode potential
The tendency of the electrode to attract electrons to itself (to be reduced)
359
What do the electrode potentials tell us about redox reactions?
The trend in the oxidising power of the elements
360
What will a halogen displace in a displacement reaction?
A halogen lower in the group form one of its ions
361
Chlorine displacement reaction with bromide ions equation
Cl2 + 2Br- —> 2Cl- + Br2
362
Explain why chlorine is able to oxidise bromide ions using electrode potential values
The electrode potential value for the 1/2Cl2/Cl- half cell is more positive than the value for the 1/2Br2/Br- half cell,so chlorine will be able to oxidise the bromide ions
363
Which ions is bromine able to oxidise and which isn’t it able to?
Oxidise iodide ions
Cannot oxidise chloride ions
364
What happens to the oxidising power of the halogens down the group and why?
Oxidising power decreases own the group
The positive value for the electrode potential decreases
365
Chlorine and iodide ions displacement equation
Cl2 + 2I- —> 2Cl- + I2
366
Chlorine reaction with…
Chloride ions
Bromide ions
Iodide ions
(No reaction)
Yellow orange Br2 formed
Black I2 formed
367
Bromine reaction with…
Chloride ions
Bromide ions
Iodide ions
No reaction
No reaction
Black I2 formed
368
Iodine reaction with…
Chloride ions
Bromide ions
Iodide ions
No reaction for all
369
Describe iodine in water
Fairly insoluble
370
What does iodine dissolve in since it’s fairly insoluble in water?
In aqueous potassium iodide
371
What happens when iodine is dissolved in aqueous potassium iodide?
Forms a red-brown solution, containing the ion I3-
372
How do we form a red-brown solution containing the ion I3-?
Iodide dissolved in aqueous potassium iodide
373
1st stage of the reaction between sodium chloride (or any chloride) with concentrated sulphuric acid
Hydrogen halide is formed in a displacement reaction
374
Sodium chloride and concentrated sulphuric acid equation
NaCl + H2SO4 —> NaHSO4 + HCl
375
HCl formed in the displacement reaction between sodium chloride and concentrated sulphuric acid
Hydrogen chloride gas, HCl
(Not hydrochloric acid!)
376
What does hydrogen chloride gas form as?
White misty fumes
377
Why does hydrogen chloride gas form white misty fumes?
Reacts with water vapour in the air
378
What can also be observed apart from misty fumes when sodium chloride reacts with sulphuric acid? Why?
Effervescence
Gas is released
379
What type of reaction is that between sodium chloride and sulphuric acid?
Displacement
Exothermic
380
Explain the second stage of the reaction between sulphuric acid and sodium chloride
The extent of further reaction depends on the reducing power of the halide ion
381
What happens to the reducing power of the halides down the group + why?
Increases down the group
The electrode potential becomes less positive
382
When does the second reaction occur between sulphuric acid and sodium chloride?
When the reducing power of the halide ion is high enough, increasing the tendency for the ion to be oxidised
383
First stage of the reaction between concentrated sulphuric acid and sodium bromide
Produces HBr gas
384
Equation for the reaction between sodium bromide and sulphuric acid
NaBr + H2SO4 —> NaHSO4 + HBr
385
Hydrogen bromide gas
HBr
386
How does hydrogen bromide gas appear?
White misty fumes
387
What’s the second stage of the reaction between sodium bromide and sulphuric acid? Explain
Further reaction with HBr
Redox reaction
HBr is oxidised
H2SO4 is reduced
388
Equation for the second reaction between sodium bromide and sulphuric acid (the HBr reacting further)
2HBr + H2SO4 —> SO2 ++ Br2 + 2H2O
389
Observations during the further reaction of HBr gas after reacting sodium bromide and sulphuric acid
Brown/orange fumes of the Br2 gas
SO2 —> pungent smelling, acidic gas
390
Describe SO2
Pungent smelling, acidic gas
391
1st reaction for the addition of sodium iodide to sulphuric acid
Produced HI gas
392
Equation for the reaction between sodium iodide and sulphuric acid
NaI + H2SO4 —> NaHSO4 + HI
393
Hydrogen chloride gas
HCl
394
Hydrogen iodide gas
HI
395
What does HI form as?
White misty fumes
396
Describe the second stage of the reaction between sodium iodide and sulphuric acid
Further reaction with HI
Redox reaction
HI —> oxidised
H2SO4 —> reduced
397
Reduction of H2SO4 equation
H2SO4 —> SO2 + S + H2S
398
Equation for the reaction between HI and sulphuric acid
2HI + H2SO4 —> SO3 + I2 + 2H2O
399
Observations for the reaction between HI and sulphuric acid
Dark purple fumes/brown solid as I2 forms
SO2 = pungent smell
When sulphur is reacted further i.e - H2SO4 —> SO2 + S + H2S, a H2S yellow solid may be seen, with a smell of rotten eggs
400
Describe H2S
Solid yellow
Smells of rotten eggs
401
Which is the only halide to react with sulphuric acid to produce a rotten egg smell?
Iodine
402
What smell does iodide form when reacting with sulphuric acid?
Rotten eggs
403
If there’s no rotten egg smell when reacting halogens with sulphuric acid, what can we deduce?
It’s not iodide, so it must be one of the other halides
404
What does how chlorine reacts with sodium hydroxide depend on?
The conditions
405
Which type of reaction occurs between chlorine and sodium hydroxide no matter the conditions?
Disproportionation reaction
406
Disproportionation reaction
Same element is oxidised and reduced during the reaction
407
What happens to chlorine during a disproportionation reaction?
Undergoes self oxidation-reduction
408
Which ions are formed when chlorine reacts with cold, dilute aqueous sodium hydroxide?
Chlorate (I) ion, ClO-
Chloride ion
409
Under which conditions does chlorine reactions with sodium hydroxide form the chlorate (I) ion, ClO- and the chloride ion?
When reacting with cold, dilute aqueous sodium hydroxide
410
Ionic equation for the reaction between chlorine and cold, dilute sodium hydroxide
Cl2 + 2OH- —> Cl- + ClO- + H2O
411
Alternative equation for the reaction between chlorine and sodium hydroxide, including the sodium ions (spectator ions)
Cl2 + 2NaOH —> NaCl + NaClO + H2O
412
chlorate (I) ion
ClO-
413
ClO-
Chlorate (I) ion
414
What is sodium chlorate (I) used in?
Bleach, germicide
415
What is sodium chloride used for?
Treating icy roads in the winter
Food preservative
416
How does sodium chloride help treat icy roads in the winter?
The impurity lowers the freezing temperature
417
Why is chlorate (I) ion important in bleaches and bactericides?
Powerful oxidising agent
418
How is the chlorate (I) ion in domestic bleaches?
As an aqueous solution of sodium chlorate (I)
419
What happens to the chlorate (I) ion at higher temperatures?
Disproportionation
420
Reaction for the disproportionation of the ClO - ion
3ClO- —> 2Cl- + ClO3-
421
What is chlorine used for and wh?
Used to kill bacteria in swimming pols and purify domestic water supplies
Strong oxidising power
422
Describe reaction with chlorine and hot, concentrated aqueous sodium hydroxide
The chlorate (I) ion disproportionates as it is formed and the chlorate (v) ion is formed with the chloride ion
423
Ionic equation for the reaction between chlorine and hot, concentrated aqueous sodium hydroxide
3Cl2 + 6OH- —> 5Cl- + ClO3- + 3H2O
424
Alternative equation for the reaction of chlorine with hot, concentrated sodium hydroxide
3Cl2 + 6NaOH —> 5NaCl + NaClO3 + 3H2O
425
Sodium chlorate (V)
NaClO3
426
NaClO3
Sodium chlorate (V)
427
Wha’s sodium chlorate (V) (NaClO3) used in?
Weed killers
428
Why is sodium chlorate (V) used as a weed killer?
Removes weeds and prevents new ones from growing
Powerful oxidising agent
429
Problem with sodium chlorate (V)
Bad for the environment (as a weed killer)
430
Commercial use of chlorine
Used to purify drinking water and swimming pools
431
Equation for the reaction between chlorine and water, showing how it’s used to purify drinking water and swimming pools
Cl2 + H2O —> HOCl + HCl
HOCl = oxidising agent
432
What is sodium chloride used for?
Enhancing the flavour of food
As a food preservative
For de-icing roads in wintry weather
433
List 3 chlorine compounds and their uses
Chlorothene = used to make plastic PVC
Organochlorine compounds have found widespread use as CFC’s, solvents, pesticides, anaesthetics
Sodium chlorate (I) is used in bleach
434
What’s chloroethene used for?
To make plastic PVC
435
Uses of organochlorines
CFC’s, solvents, pesticides, anaesthetics
436
Commercial uses of bromine
Used to make…
Bromide salts = used in photography
Organobromide compounds such as 1,2 dibromoethene (petrol additive)
437
Commercial uses of iodine
In alcoholic solution is used to disinfect wounds
Potassium iodide is used in the photographic industry
438
How do we test for halide ions?
1.) nitric acid
2.) add silver nitrate solution
= precipitate of filter halide forms
439
Why do we add nitric acid when testing for halide ions?
To dilute any OH- ions present
440
Colour of silver halide precipitate with Cl-
White
441
Colour of silver halide precipitate with Br-
Cream precipitate
442
Colour of silver halide precipitate with I-
Yellow precipitate
443
Equation for the reaction between silver nitrate and chloride ions
Ag+ (aq) + Cl- (aq) —> AgCl (s)
444
Testing silver halide ions further after nitric acid + observations
Add ammonia (NH3)
Cl- = soluble
Br- = mostly insoluble
I- = insoluble
445
What type of analysis are flame tests?
Qualitative
446
What are flame tests used for?
To identify group I and II cations
447
flame test method
1. Dip a platinum wire/wet wooden splint into concentrated HCl
2. Dip the top of the wire in the solid to be tested
3. Tip of the wire is placed in the blue Bunsen flame
4. Note the colour of the flame
448
Lithium cation flame colour
Red
449
Sodium cation flame colour
Intense yellow
450
Potassium cation flame colour
Lilac
451
Magnesium cation flame colour
No colour
452
Calcium cation flame colour
Brick red
453
Strontium cation flame colour
Crimson
454
Barium cation flame colour
Apple green
455
Beryllium showing acidic behaviour equation
Be(OH)2 + 2OH- —> [Be(OH)4] 2-
Tetrahydroxo beryllate (II)
456
Tin showing acidic behaviour equation
Sn(OH)2 + 2OH- —> [Sn(OH)4]2-
Tetrahydroxo stannate (II)
457
Equation aluminium showing acidic behaviour
Al(OH)3 + OH- —> [Al(OH)4]-
Tetrahydroxo laminate (III)
458
When showing basic or acidic behaviour is there always water formed?
Basic
459
Electron deficiency
*Outer shell* of electrons is not full (short of an octet)
460
List and explain 4 reaction of Pb2+ ions
NaOH
Pb2+ + OH- (aq) —> Pb(OH)2 (s) white ppt
XsNaOH
Precipitate dissolves to give colourless solution
Pb(OH)2 (s) + OH- (aq) —> [Pb(OH)4]^2- (aq)
HCl
Pb2+ (aq) + Cl (aq) —> PbCl2 (s) white ppt
KI
Pb2+ (aq) + 2I- —> PbI2 (s) yellow ppt
461
How does H2SO4 undergo further reaction in the reaction with sodium iodide?
HI reduces the sulphur in H2SO4
462
Is it iodine or iodide ions that make a good reducing agent?
Iodide ions
463
Why are the oxides of lead both solids?
PbO = yellow solid
PbO2 = dark brown solid
They’re ionic with strong attractions between the ions
464
What is the only soluble lead salt?
Lead nitrate
465
What occurs during all of the halide ions and concentrated sulphuric acid reactions that we tend to not really mention?
White solid forms
466
Why can’t nitrogen not expand its octet?
No available d orbitals in the outer shell
467
Describe the observations of the reaction between SiCl4 and water
Violent reaction —> white precipitate and steamy fumes
