AS Chemistry Unit 3: Inorganic Chemistry Cambridge A-level Centre Section 1: Period 3 Part 1: Atomic and Physical Properties of the Elements Atomic Properties Task 1 Can you write the electronic configurations of the Period 3 elements in the table below? Na [Ne] Mg [Ne] Al [Ne] Si [Ne] P [Ne] S [Ne] Cl [Ne] Ar [Ne] In each case, [Ne] represents the complete electronic structure of a neon atom. Atomic radius The graph shows how the atomic radius changes as you go across Period 3: Task 2 An atomic radius is a measure of the distance from the nucleus to the bonding pair of electrons. i) What is the general trend in atomic radius across Period 3? ................................................................................................................................................ 0 0.05 0.1 0.15 0.2 0.25 Na Mg Al Si P S Cl Ar Atomic radius (nm) Element
Transcript
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Section 1: Period 3 Part 1: Atomic and Physical Properties of the Elements
Atomic Properties
Task 1
Can you write the electronic configurations of the Period 3 elements in the table below?
Na [Ne]
Mg [Ne]
Al [Ne]
Si [Ne]
P [Ne]
S [Ne]
Cl [Ne]
Ar [Ne]
In each case, [Ne] represents the complete electronic structure of a neon atom.
Atomic radius
The graph shows how the atomic radius changes as you go across Period 3:
Task 2
An atomic radius is a measure of the distance from the nucleus to the bonding pair of electrons.
i) What is the general trend in atomic radius across Period 3?
iii) Why do you think the value for argon does not follow the expected trend?....................................................................................................................................
The radii of the ions also decrease across a Period, but it must be remembered that elements on the left form cations and elements on the right form anions.
This has already been covered in Unit 1. You should be able to use your scientific knowledge to explain in detail the trend shown in the graph below:
Remember, first ionisation energy is governed by:
• the charge on the nucleus; • the distance of the outer electron from the nucleus; • the amount of screening by inner electrons; • whether the electron is alone in an orbital or one of a pair.
Electronegativity
Task 4
Can you write a definition for the term ‘electronegativity?
Electronegativity is .................................................................................................................
The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7.
The trend
The Electronegativity value for chlorine is given. Can you sketch the rest of the graph?
Explaining the trend
The trend is explained in exactly the same way as the trend in atomic radii.
As you go across the period, the bonding electrons are always in the same level – level 3. They are always being screened by the same inner electrons.
All that differs is the number of protons in the nucleus. As you go from sodium to chlorine, the number of protons steadily increases and so attracts the bonding pair more closely.
Task 5
Why is no value included for argon?....................................................................................
0
0.5
1
1.5
2
2.5
3
3.5
Na Mg Al Si P S Cl Ar
Paul
ing
elec
tron
egat
ivit
y va
lue
Element
AS
Che
mis
try
Uni
t 3:
Ino
rgan
ic C
hem
istr
y
Cam
brid
ge A
-lev
el C
entr
e
Physical Properties
Her
e w
e lo
ok a
t th
e el
ectr
ical
con
duct
ivit
y, a
nd t
he m
elti
ng a
nd b
oilin
g po
ints
of
the
elem
ents
. T
o un
ders
tand
the
se,
you
firs
t ha
ve t
o un
ders
tand
the
str
uctu
re o
f ea
ch o
f th
e el
emen
ts.
Tas
k 6
Use
you
r kn
owle
dge
from
Uni
t 1
to c
ompl
ete
the
tabl
e be
low
.
Na
Mg
Al
Si
P S
Cl
Ar
Structure
Type of element
Bonding
Formula
Type of force broken
on melting/boiling
Does
the
element
conduct electricity?
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Task 7
Use the information in the table above to answer the following questions.
1. (a) Explain why electrical conductivity decreases across Period 3 from sodium to phosphorus.
References A-level Chemistry pages 197-207 Chemistry in Context pages 42-47, 169-176 Learning Objectives Candidates should be able to:
• describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements (see the Data Booklet)
• explain qualitatively the variation in atomic radius and ionic radius • interpret the variation in melting point and in electrical conductivity in
terms of the presence of simple molecular, giant molecular or metallic bonding in the elements
• explain the variation in first ionisation energy.
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Section 1: Period 3 Part 2: Reactions with oxygen
Reactions of Period 3 Elements with oxygen
As we move across a period of the Periodic Table from left to right, we can see that there are small, but regular changes in atomic structure. These small changes can produce very big variations in the chemical and physical properties of the elements and their compounds. The pattern is then repeated as you go across the next period. The occurrence of periodic patterns is called Periodicity. The physical properties of the Period 3 elements have been discussed previously. Here we look at their chemical properties.
Task 1 Can you complete the table below?
Group number
1 2 3 4 5 6
Element in Period 3
Nuclear charge
[Ne] electronic configuration
Trend in Atomic radius
Trend in 1st ionisation energy
Trend in electronegativity
Formula of oxide/s
The reactivity and properties of elements depend upon a combination of things: nuclear charge, size of the atom, the number of outer electrons and the amount of shielding, and these will also help to explain and predict the properties of many compounds. It is therefore very useful if you are familiar with the trends shown in the table above. You should not, however, lose sight of your general chemistry knowledge.
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Task 2 Can you complete the notes below by adding the missing words and writing the appropriate balanced equations? With oxygen The solid elements in Period 3 all burn in air or oxygen when ignited. Sodium burns with a __________ flame, forming the oxide: Magnesium, aluminium, silicon and phosphorus burn when ignited, emitting a very bright white light and white smoke of the oxides: These reactions are all very exothermic. Sulphur burns with a __________ flame but much less vigorously than the elements above, to form the pungent, colourless gas ____________________: In an excess of pure oxygen, some SO3 is also formed. This utilises the highest oxidation state of sulphur. Task 3 Can you complete the table below? Na2O MgO Al2O3 SiO2 P4O10 SO2 Tm/K 1548 3125 2345 1883 573 200
Bonding
Structure Use your scientific knowledge to explain the changes in melting point.
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
...................................................................................................................................................... ...................................................................................................................................................... ...................................................................................................................................................... ...................................................................................................................................................... ...................................................................................................................................................... Acid-base properties of the oxides of Period 3 Elements The table below shows the change in pH for solutions of the oxides across Period 3. There is an evident trend for alkaline oxides → acidic oxides as the bonding changes from ionic to covalent, but it is masked by the change in solubilities. A substance will only change the pH of water if it dissolves. Task 4 Can you write balanced equations for the reactions of each oxide with water to give the appropriate pH value. Oxide Reaction with water pH
Na2O
14
MgO
9
Al2O3
7
SiO2
7
P4O10
0
SO2
3
SO3
0
Ionic oxides The oxide ion is too highly charged to exist on its own in water. It attracts water molecules and hydrolyses to form OH-
(aq) ions:
O2− + H2O → 2 OH−
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Na2O: very soluble in water → lots of oxide ions present and therefore many OH-
(aq) ions will be formed. MgO: less soluble in water due (in part) to higher lattice energy → less oxide ions and less OH-
(aq) ions. Al2O3: insoluble → no oxide ions will be present and hence no OH-
(aq) ions will be formed. Aluminium oxide is an ionic solid with some covalent character. It is amphoteric; it will act as both an acid and a base.
BASE: Al2O3 + 3H2SO4 → Al2(SO4)3 + 3H2O
ACID: Al2O3 + 2NaOH + 3H2O → 2NaAl(OH)4 Covalent oxides SiO2: insoluble in water → no ions produced. P4O10, SO3: these acidic oxides rapidly dissolve in water and give the corresponding acids. They are very soluble, so many H+
(aq) ions are formed, giving a low pH. SO2 is less soluble so a weaker acid is formed. References A-level Chemistry pages 209-211 Chemistry in Context pages 177-184 Learning Objectives Candidates should be able to:
• describe the reactions, if any, of the elements with oxygen to give Na2O, MgO, Al2O3, P4O10, SO2 and SO3.
• state and explain the variation in oxidation number of the oxides. • describe the reactions of the oxides with water. • describe and explain the acid/base behaviour of oxides and hydroxides,
including, where relevant, amphoteric behaviour in reaction with NaOH and acids.
AS
Che
mis
try
Uni
t 3:
Ino
rgan
ic C
hem
istr
y
Cam
brid
ge A
-lev
el C
entr
e
Section 1: Period 3
Part 3: Reactions with chlorine
All
of t
he c
hlor
ides
of
the
peri
od 3
ele
men
ts c
an b
e fo
rmed
by
dire
ct c
ombi
nati
on w
ith
chlo
rine
but
you
r sy
llabu
s is
onl
y in
tere
sted
in t
he r
eact
ions
of
Na,
Mg,
Al,
Si a
nd P
. T
ask
1 U
se t
he in
form
atio
n on
pag
es 4
7-49
and
179
-180
of
‘Che
mis
try
in C
onte
xt’ a
nd p
ages
209
and
211
-212
of
‘AS
leve
l Ch
emis
try’
to
com
plet
e th
e ta
ble
belo
w:
Element
Na
Mg
Al
Si
P
Description of reaction
with chlorine
Formula of chloride/s
Oxidation state of
period 3 element
State of chloride at
r.t.p.
b.pt. of chloride (oC)
Structure of chloride
Bonding in chloride
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Task 2 Read the information sheet ‘Aluminium chloride’ and answer the question below. i. At room temperature aluminium chloride is considered to be ionic. What two
................................................................................................................................... ii. At temperatures above ~200oC aluminium chloride exists as a dimer, Al2Cl6.
This molecule contains both ordinary and dative covalent bonds. Draw a diagram to show the two types of bonding.
Reaction of the period 3 chlorides with water As Period 3 is crossed, the reactions of the chlorides with water become increasingly more violent.
• Ionic chlorides usually dissolve in water to form neutral solutions containing the hydrated parent ions.
• Covalent chlorides are hydrolysed by water to form acidic solutions containing HCl.
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Task 3 Can you summarise these reactions in the table below? Reaction with water pH
7
6/7
3
0
0
References A-level Chemistry pages 209-210 and 211- 212 Chemistry in Context pages 47-49 and 179-180 Learning Objectives Candidates should be able to:
• describe the reactions, if any, of the elements with chlorine to give NaCl, MgCl2, Al2Cl6, SiCl4, and PCl5.
• describe and explain the reactions of the chlorides with water.
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Section 2: Group 7 Part 1: Properties of the elements
Please note that fluorine will not be included in discussions of Group VII, because its small size and high electronegativity give it some anomalous properties.
Task 1 Use your general scientific knowledge to fill in the gaps in the paragraph below. The halogens are the __________ in _______7 of the __________ Table. All
halogen atoms have 7 __________ in the outer shell. The halogens are the
most __________ group of non-metals, and none of them is found
__________ in the elemental form. They are all found in __________, often
as __________ ions (a singly __________ charged ion, e.g. Br-).
Fluorine and chlorine are the most __________ halogens, bromine occurs in
smaller __________, iodine is quite __________ and astatine is an
All the halogen elements occur as __________ molecules, e.g. F2. The atoms
are linked by a single __________ bond.
In compounds, a halogen atom can attain stability by:
• __________ an electron from a __________ atom to form a halide ion
in an __________ bonded compound,
• sharing an electron from another atom in a __________ bonded
compound.
Task 2 Draw electron dot-cross diagrams for NaCl and HCl in the space below.
AS Chemistry
Cambridge A-level Centre
Task 3
Use the information in your textbooks to complete the table below.
Element Electronic structure
Appearance at r.t.p.
Fluorine Chlorine Bromine Iodine Task 4 From your knowledge of the structure and bonding of the halogens, explain why they are more soluble in organic solvents than in aqueous solution.
You will see that both melting points and boiling points rise as you descend the Group.
Task 5
Use your scientific knowledge to explain the trend in volatility of the halogens. .......................................................................................................................................................
Watch the short video extract. For each reaction write down your observations.
Halide ion in aqueous solution Halogen Chloride, Cl- Bromide, Br- Iodide, I- Chlorine, Cl2
Bromine, Br2
Iodine, I2
The oxidising power of the halogens
Imagine the reaction between one halogen (chlorine, say) and the ions of another one (iodide ions, perhaps). The iodide ions will be in a solution of a salt like sodium or potassium iodide. The sodium or potassium ions will be spectator ions, and are completely irrelevant to the reaction. In the chlorine and iodide ion case, the reaction would be:
Cl2 + 2I- à 2Cl- + I2
AS Chemistry
Cambridge A-level Centre
The iodide ions have lost electrons to form iodine molecules. They have been __________.
The chlorine molecules have gained electrons to formbeen __________.
This is obviously a redox reaction in which chlorine is acting as _______________ agent.
Fluorine
We'll have to exclude fluorine from this descriptive bit, because it is too strong an oxidising agent. Fluorine oxido simple solution reactions with it.
Chlorine, bromine and iodine
In each case, a halogen higher in the Group can oxidise the ions of one lower down. For example, chlorine can oxidise the bromide ions potassium bromide solution) to bromine:
The bromine appears as an orange solution.
As you have seen above, chlorine can also oxidise iodide ions (in, for example, potassium iodide solution) to iodine:
The iodine appears either as a amount of chlorine you use, or as a dark grey precipitate if the chlorine is in excess.
Bromine can only oxidise iodide ions to iodine. It isn't a strong enough oxidising agent to convert chloride ions into chlorreverse of that happening.)
A red/brown solution of iodine is formed (see the note above) until the bromine is in excess. Then you get a dark grey precipitate.
Unit 3: Inorganic Chemistry
The iodide ions have lost electrons to form iodine molecules. They have been
The chlorine molecules have gained electrons to form chloride ions. They have
This is obviously a redox reaction in which chlorine is acting as _______________ agent.
We'll have to exclude fluorine from this descriptive bit, because it is too strong an oxidising agent. Fluorine oxidises water to oxygen and so it is impossible to do simple solution reactions with it.
Chlorine, bromine and iodine
In each case, a halogen higher in the Group can oxidise the ions of one lower down. For example, chlorine can oxidise the bromide ions (in, for example, potassium bromide solution) to bromine:
The bromine appears as an orange solution.
As you have seen above, chlorine can also oxidise iodide ions (in, for example, potassium iodide solution) to iodine:
The iodine appears either as a red/brown solution if you are mean with the amount of chlorine you use, or as a dark grey precipitate if the chlorine is in
Bromine can only oxidise iodide ions to iodine. It isn't a strong enough oxidising agent to convert chloride ions into chlorine. (You have just seen exactly the reverse of that happening.)
A red/brown solution of iodine is formed (see the note above) until the bromine is in excess. Then you get a dark grey precipitate.
Unit 3: Inorganic Chemistry
The iodide ions have lost electrons to form iodine molecules. They have been
chloride ions. They have
This is obviously a redox reaction in which chlorine is acting as
We'll have to exclude fluorine from this descriptive bit, because it is too strong dises water to oxygen and so it is impossible to
In each case, a halogen higher in the Group can oxidise the ions of one lower (in, for example,
As you have seen above, chlorine can also oxidise iodide ions (in, for example,
red/brown solution if you are mean with the amount of chlorine you use, or as a dark grey precipitate if the chlorine is in
Bromine can only oxidise iodide ions to iodine. It isn't a strong enough oxidising ine. (You have just seen exactly the
A red/brown solution of iodine is formed (see the note above) until the bromine
AS Chemistry
Cambridge A-level Centre
Iodine won't oxidise any of the other halide ions (unlesssome extremely radioactive and amazingly rare astatide ions bottom of this Group).
To summarise
• Oxidation is loss of electrons. Each of the elements (for example, chlorine) could potentially take electrons from sotheir ions (e.g. Cl-). That means that they are all potentially oxidising agents.
• Fluorine is such a powerful oxidising agent that you can't reasonably do solution reactions with it.
• Chlorine has the ability to take electrons from bothiodide ions. Bromine and iodine can't get those electrons back from the chloride ions formed. That means that chlorine is a more powerful oxidising agent than either bromine or iodine.
• Similarly bromine is a more powerful oxidising agent can remove electrons from iodide ions to give iodine get them back from the bromide ions formed.
This all means that oxidising ability falls as you go down the Group.
Reactions of the halogens with hydrogen
The halogens all react with hydrogen to give the corresponding hydrogen halide. The relative reactivity of the halog
• Fluorine explodes with hydrogen even in the dark at • The reaction with chlorine is explosive when exposed to ultraviolet light.• With bromine, the reaction occurs slowly on heating.• The reaction with iodine
Fluorine Chlorine
Bromine
Iodine
Relative oxidising power
Unit 3: Inorganic Chemistry
Iodine won't oxidise any of the other halide ions (unless you happened to have some extremely radioactive and amazingly rare astatide ions - astatine is at the
Oxidation is loss of electrons. Each of the elements (for example, chlorine) could potentially take electrons from something else to make
). That means that they are all potentially oxidising
Fluorine is such a powerful oxidising agent that you can't reasonably do solution reactions with it. Chlorine has the ability to take electrons from both bromide ions and iodide ions. Bromine and iodine can't get those electrons back from the chloride ions formed. That means that chlorine is a more powerful oxidising agent than either bromine or iodine. Similarly bromine is a more powerful oxidising agent than iodine. Bromine can remove electrons from iodide ions to give iodine - and the iodine can't get them back from the bromide ions formed.
This all means that oxidising ability falls as you go down the Group.
Reactions of the halogens with hydrogen
The halogens all react with hydrogen to give the corresponding hydrogen halide. The relative reactivity of the halogens is well illustrated by these reactions:
Fluorine explodes with hydrogen even in the dark at -200o
The reaction with chlorine is explosive when exposed to ultraviolet light.With bromine, the reaction occurs slowly on heating. The reaction with iodine is incomplete on heating.
Unit 3: Inorganic Chemistry
you happened to have astatine is at the
Oxidation is loss of electrons. Each of the elements (for example, mething else to make
). That means that they are all potentially oxidising
Fluorine is such a powerful oxidising agent that you can't reasonably do
bromide ions and iodide ions. Bromine and iodine can't get those electrons back from the chloride ions formed. That means that chlorine is a more powerful
than iodine. Bromine and the iodine can't
This all means that oxidising ability falls as you go down the Group.
The halogens all react with hydrogen to give the corresponding hydrogen halide. ens is well illustrated by these reactions:
oC. The reaction with chlorine is explosive when exposed to ultraviolet light.
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Task 7
Write equations, including state symbols, for the reactions of chlorine, bromine and iodine with hydrogen gas.
The trend is further illustrated by consideration of the ∆Hfθ values:
HCl HBr HI ∆Hf
θ (kJ mol-1) -92 -36 +26
The hydrogen halides formed are all simple molecular compounds which are gaseous at room temperature. As the size of the halogen atom increases down the Group, the H-X bond length also increases. Consequently, the H-X bond enthalpy decreases down the Group and the hydrogen halides become less stable:
• hydrogen chloride is stable at 1500oC • hydrogen bromide decomposes appreciably at 800oC, and • hydrogen iodide decomposes appreciably at 500oC.
If a red hot glass rod is placed in a gas jar containing hydrogen iodide, purple fumes of iodine are seen. This shows that iodide ions are easily oxidised and can act as strong reducing agents.
References A-level Chemistry pages 237-240 Chemistry in Context pages 227-231 Learning Objectives Candidates should be able to:
• describe the trends in volatility and colour of chlorine, bromine and iodine.
• interpret the volatility of the elements in terms of van der Waals’ forces. • describe the relative reactivity of the elements as oxidising agents. • describe and explain the reactions of the elements with hydrogen.
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
• describe and explain the relative thermal stabilities of the hydrides. • interpret these relative stabilities in terms of bond energies.
AS Chemistry
Cambridge A-level Centre
Section 2: Group 7
Here we look at the redox reactions involving halide ions and concentrated sulphuric acid. We use these reactions to discuss the trend in of the ions as you go from fluoride to chloride to bromide to iodide.
The Facts
There are two different types of reaction which might go on when concentrated sulphuric acid is added to a solid ionic halide like sodium fluoride, chloride, bromide or iodide. The concentrated sulphuric acid can act both as an acid and as an oxidising agent.
Concentrated sulphuric acid acting as an acid
The concentrated sulphuric acid gives a hydrogen ion to the halide ion to produce a hydrogen halide. Because this is a gas, it immediately escapes from the system. If the hydrogen halide is exposed to moist air, you see it as sfumes.
As an example, concentrated sulphuric acid reacts with solid sodium chloride in the cold to produce hydrogen chloride and sodium hydrogensulphate.
All of the halide ions (fluoride, chloride, bromide and iodide) behave similarly.
The reduction of concentrated sulphuric acid
With fluoride or chloride ions
The fluoride and chloride ions aren't strong enough reducing agents to reduce the sulphuric acid.
Whichever way you look at it, all you get is the hydrogen halide!
That isn't true, though, w
With bromide ions
The bromide ions are strong enough reducing agents to reduce the concentrated sulphuric acid to sulphur dioxide gasoxidised to bromine.
Unit 3: Inorganic Chemistry
Part 2: Properties of the halides
Here we look at the redox reactions involving halide ions and concentrated sulphuric acid. We use these reactions to discuss the trend in of the ions as you go from fluoride to chloride to bromide to iodide.
nt types of reaction which might go on when concentrated sulphuric acid is added to a solid ionic halide like sodium fluoride, chloride, bromide or iodide. The concentrated sulphuric acid can act both as an acid and
huric acid acting as an acid
The concentrated sulphuric acid gives a hydrogen ion to the halide ion to produce a hydrogen halide. Because this is a gas, it immediately escapes from the system. If the hydrogen halide is exposed to moist air, you see it as s
As an example, concentrated sulphuric acid reacts with solid sodium chloride in the cold to produce hydrogen chloride and sodium hydrogensulphate.
All of the halide ions (fluoride, chloride, bromide and iodide) behave similarly.
oncentrated sulphuric acid
With fluoride or chloride ions
The fluoride and chloride ions aren't strong enough reducing agents to reduce
Whichever way you look at it, all you get is the hydrogen halide!
That isn't true, though, with bromides and iodides.
The bromide ions are strong enough reducing agents to reduce the concentrated to sulphur dioxide gas. In the process the bromide ions are
Unit 3: Inorganic Chemistry
roperties of the halides
Here we look at the redox reactions involving halide ions and concentrated sulphuric acid. We use these reactions to discuss the trend in reducing ability of the ions as you go from fluoride to chloride to bromide to iodide.
nt types of reaction which might go on when concentrated sulphuric acid is added to a solid ionic halide like sodium fluoride, chloride, bromide or iodide. The concentrated sulphuric acid can act both as an acid and
The concentrated sulphuric acid gives a hydrogen ion to the halide ion to produce a hydrogen halide. Because this is a gas, it immediately escapes from the system. If the hydrogen halide is exposed to moist air, you see it as steamy
As an example, concentrated sulphuric acid reacts with solid sodium chloride in the cold to produce hydrogen chloride and sodium hydrogensulphate.
All of the halide ions (fluoride, chloride, bromide and iodide) behave similarly.
The fluoride and chloride ions aren't strong enough reducing agents to reduce
The bromide ions are strong enough reducing agents to reduce the concentrated . In the process the bromide ions are
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Task 1
Can you write two electron half equations for these steps, then combine them to form the overall redox equation for the reaction.
Task 2
Complete the notes below.
The bromide ions reduce the sulphuric acid to sulphur dioxide gas. This is a decrease of oxidation state of the sulphur from _____ in the sulphuric acid to _____ in the sulphur dioxide.
What you see in this reaction are the steamy fumes of __________ __________ contaminated with the brown colour of __________ vapour. The sulphur dioxide is a colourless gas, so you couldn't observe its presence directly.
With iodide ions
Iodide ions are stronger reducing agents than bromide ions are. They are oxidised to iodine by the concentrated sulphuric acid. The reduction of the sulphuric acid is more complicated than before. The iodide ions are powerful enough reducing agents to reduce it
• first to sulphur dioxide • then to sulphur itself • and all the way to hydrogen sulphide.
Task 3
Can you give the oxidation state of sulphur in each of the species above?
The most important of this mixture of reduction products is probably the hydrogen sulphide.
AS Chemistry
Cambridge A-level Centre
Task 4
Can you write the half equations for these reactions and combine them to give the overall redox equation?
This time what you see is a trace of steamy fumes of hydrogen iodide, but mainly lots of iodine. The reaction is exothermic and so purple iodine vapour is formed, and probably dark grey solid iodine condensing around the top of the tube.
You won't see the colourless hydrogen sulphide gas, but might pick up its "bad egg" smell if you were foolish enough to smell the intensely poisonous gases evolved!
Summary of the trend in reducing ability
• Fluoride and chloride ions won't reduce concentrated sulphuric acid• Bromide ions reduce the sulphuric acid to sulphur dioxide. In the process,
the bromide ions are oxidised to bromine.• Iodide ions reduce the sulphuric acid to a mixture of products including
hydrogen sulphide. The iodide ions are oxidised to iodine.• Reducing ability of the halide ions increases as you go down the Group.
(Obviously, this is the opposite direction to their oxidising ability!!)
Relative reducingpower
Unit 3: Inorganic Chemistry
Can you write the half equations for these reactions and combine them to give the overall redox equation?
This time what you see is a trace of steamy fumes of hydrogen iodide, but mainly lots of iodine. The reaction is exothermic and so purple iodine vapour is formed, and probably dark grey solid iodine condensing around the top of the
colourless hydrogen sulphide gas, but might pick up its "bad egg" smell if you were foolish enough to smell the intensely poisonous gases
Summary of the trend in reducing ability
Fluoride and chloride ions won't reduce concentrated sulphuric acidBromide ions reduce the sulphuric acid to sulphur dioxide. In the process, the bromide ions are oxidised to bromine. Iodide ions reduce the sulphuric acid to a mixture of products including hydrogen sulphide. The iodide ions are oxidised to iodine. Reducing ability of the halide ions increases as you go down the Group.(Obviously, this is the opposite direction to their oxidising ability!!)
Fluoride Chloride Bromide Iodide
Relative reducing power
Unit 3: Inorganic Chemistry
Can you write the half equations for these reactions and combine them to give
This time what you see is a trace of steamy fumes of hydrogen iodide, but mainly lots of iodine. The reaction is exothermic and so purple iodine vapour is formed, and probably dark grey solid iodine condensing around the top of the
colourless hydrogen sulphide gas, but might pick up its "bad egg" smell if you were foolish enough to smell the intensely poisonous gases
Fluoride and chloride ions won't reduce concentrated sulphuric acid. Bromide ions reduce the sulphuric acid to sulphur dioxide. In the process,
Iodide ions reduce the sulphuric acid to a mixture of products including
Reducing ability of the halide ions increases as you go down the Group. (Obviously, this is the opposite direction to their oxidising ability!!)
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Testing for halide ions
Using silver nitrate solution
Carrying out the test
This test has to be done in solution. If you start from a solid, it must first be dissolved in pure water.
The solution is acidified by adding dilute nitric acid. (Remember: silver nitrate + dilute nitric acid.) The nitric acid reacts with, and removes, other ions that might also give a confusing precipitate with silver nitrate.
Silver nitrate solution is then added.
Task 5
Can you complete the table below to show what you would observe in each case?
ion present observation
F-
Cl-
Br-
I-
All of the precipitates change colour if they are exposed to light - taking on grey or purplish tints.
The absence of a precipitate with fluoride ions doesn't prove anything unless you already know that you must have a halogen present and are simply trying to find out which one. All the absence of a precipitate shows is that you haven't got chloride, bromide or iodide ions present.
The chemistry of the test
The precipitates are the insoluble silver halides - silver chloride, silver bromide or silver iodide.
AS Chemistry
Cambridge A-level Centre
Silver fluoride is soluble, and so you don't get a precipitate.
Confirming the precipitate using amm
Carrying out the confirmation
Ammonia solution is added to the precipitates.
Task 6
Can you complete the table below to show what you would observe in each case?
original precipitate
AgCl
AgBr
AgI
The chemistry of the test
When the precipitate of silver chloride or silver bromide dissolves, the silver ion forms a complex ion:
E.g. AgCl(s) + 2NH
The reactions of chlorine with sodium hydroxide
Chlorine is used in the manufacture of bleach, which is widely used as a disinfectant. Bleach is produced by passing chlorine gas up a tower, down which cold, dilute aqueous sodium hydroxide is flowing. The equation for the reaction which takes place in the tower is as follows:
Cl2(g) + 2NaOH(aq) →
Unit 3: Inorganic Chemistry
Silver fluoride is soluble, and so you don't get a precipitate.
Confirming the precipitate using ammonia solution
Carrying out the confirmation
Ammonia solution is added to the precipitates.
Can you complete the table below to show what you would observe in each case?
observation
The chemistry of the test
When the precipitate of silver chloride or silver bromide dissolves, the silver
2NH3(aq) → [Ag(NH3)2]+(aq) + Cl
The reactions of chlorine with sodium hydroxide
Chlorine is used in the manufacture of bleach, which is widely used as a disinfectant. Bleach is produced by passing chlorine gas up a tower, down which
aqueous sodium hydroxide is flowing. The equation for the reaction e tower is as follows:
→ NaClO (aq) + NaCl(aq) +
Sodium chlorate (I)
Unit 3: Inorganic Chemistry
Can you complete the table below to show what you would observe in each case?
When the precipitate of silver chloride or silver bromide dissolves, the silver
+ Cl-(aq)
Chlorine is used in the manufacture of bleach, which is widely used as a disinfectant. Bleach is produced by passing chlorine gas up a tower, down which
aqueous sodium hydroxide is flowing. The equation for the reaction
H2O(l)
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Task 7
i. Using oxidation states can you identify which element is oxidised and which is reduced in the equation above?
2. Draw a dot-cross diagram to show the bonding in a molecule of nitrogen:
3. Nitrogen is used as an inert atmosphere in food storage and during some metalworking processes where oxidation must be prevented. Why do you think nitrogen gas is so unreactive? ................................................................................................................................................ ................................................................................................................................................
4. Can you list two situations in which nitrogen gas will react with oxygen gas?
Why do these reactions occur? ................................................................................................................................................
................................................................................................................................................ 5. Atmospheric nitrogen is ‘fixed’ as nitrate ions by bacteria in the roots of
certain plants such as peas, beans and clover. How can bacteria enable this to happen? ................................................................................................................................................
6. Nitrogen is converted into ammonia in the Haber Process. Can you write a
balanced equation for this reaction. 7. Ammonia readily forms ammonium ions through co-ordinate bonds. Can you
draw a diagram to represent this below:
Ammonia is acting as what type of species in this reaction?
................................................................................................................................................ 8. How would you test for the presence of an ammonium ion?
Can you name three of these salts? .......................................................................................................................................... ..........................................................................................................................................
10. Ammonia is also converted into nitric acid. Can you write a balanced equation for the overall reaction and give the conditions needed?
11. There are concerns over the use of nitrate fertilisers as they have been
linked to the process of eutrophication. Can you draw a flow diagram for this process below:
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Sulphur
Use your scientific knowledge and the information on pages 262-264 of your textbook to help you answer the following questions.
1. The Earth’s atmosphere contains oxides of nitrogen and sulphur dioxide from natural sources. Can you name some of these natural sources?
a. nitrogen oxides........................................................................................................
b. sulphur dioxide........................................................................................................
2. Sulphur dioxide in the atmosphere has been linked to human respiratory problems such as bronchitis and asthma. It is however, used as a food preservative. How does sulphur dioxide preserve food?
In each case, the two outer electrons feel a core charge of 2+ from the nucleus. The positive charge on the nucleus is shielded by the inner electrons.
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
This is equally true for all the other atoms in Group 2. Work it out for calcium if you aren't convinced.
The only factor which is going to affect the size of the atom is therefore the number of layers of inner electrons which have to be fitted in around the atom. Obviously, the more layers of electrons you have, the more space they will take up - electrons repel each other. That means that the atoms are bound to get bigger as you go down the Group.
Trends in First Ionisation Energy
First ionisation energy is the energy needed to remove the most loosely held electron from each of one mole of gaseous atoms to make one mole of singly charged gaseous ions - in other words, for 1 mole of this process:
X(g) → X+(g) + e-
Notice that first ionisation energy falls as you go down the group.
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. All of these elements have a low electronegativity (due to their small core charge).
As the metal atoms get bigger, any bonding pair gets further and further away from the metal nucleus, and so is less strongly attracted towards it. In other words, as you go down the Group, the elements become less electronegative.
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
As you go down the Group, the bonds formed between these elements and other things such as chlorine become more and more ionic. The bonding pair is increasingly attracted away from the Group 2 element towards the chlorine (or whatever).
References A-level Chemistry pages 215-218 Chemistry in Context pages 211-213 Learning Objectives Candidates should be able to interpret, and make predictions from, the trends in physical and chemical properties of the elements and their compounds.
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
Section 4: Group 2 Section 2: Chemical properties
When Group 2 metals react, they are oxidised from a state of 0 to +2, forming M2+ ions. The elements are powerful reducing agents. Group 2 metals become increasingly reactive as you go down the Group. This is because the distance between the nucleus and the outer electrons increases and so they become easier to remove. Beryllium is markedly different from the other members of the group and will not be considered in detail.
Reactions of the metals with water
Beryllium
Beryllium has no reaction with water or steam even at red heat.
Magnesium
Magnesium burns in steam to produce magnesium oxide and hydrogen.
Very clean magnesium has a very slight reaction with cold water. The reaction soon stops because the magnesium hydroxide formed is almost insoluble in water and forms a barrier on the magnesium preventing further reaction.
Task 1
Can you write equations for the reaction of Mg with both cold and hot water?
As a general rule, if a metal reacts with cold water, you get the metal hydroxide. If it reacts with steam, the metal oxide is formed. This is because the metal hydroxides thermally decompose (split up on heating) to give the oxide and water.
Calcium, strontium and barium
These all react with cold water with increasing vigour to give the metal hydroxide and hydrogen. Strontium and barium have reactivities similar to lithium in Group 1 of the Periodic Table.
AS Chemistry
Cambridge A-level Centre
The equation for the reacti
The hydroxides aren't very soluble, but they get more soluble as you go down the Group. The calcium hydroxide formed shows up mainly as a white precipitate (although some does dissolve). You get less precipitate as yobecause more of the hydroxide dissolves in the water.
Reactions of the metals with oxygen
The group 2 elements all burn in oxygen to form white solid oxides.
2M
Beryllium is reluctant to burn unless it is in the form of dust or powder. flame colour is white. Beryllium has a very strong (but very thin) layer of beryllium oxide on its surface, and this prevents any new oxygen getting at the underlying beryllium to react with it.
Magnesium burns with a brilliant white flame. The others burn with characteristic flame colours:
Calcium: brick red Strontium: crimson red Barium: apple green Reactions of the oxides with water
The metal oxides are all basic and each reacts with water, forming a solution of the hydroxide which is alkaline. The reaction of magnesium oxide with water is slow; the other oxides react readily with water.
The solubilities of the hydroxides formed incrsolutions become more alkaline (approx. pH9
Thermal stability of Group 2 carbonates and nitrates
Thermal decomposition is when a substance breaks down (decomposes) when heated. The more thermally stable a substance is, thbreak it down.
You are only expected to describe the trend in thermal stability, not explain it.
Unit 3: Inorganic Chemistry
The equation for the reactions of any of these metals would be:
The hydroxides aren't very soluble, but they get more soluble as you go down the Group. The calcium hydroxide formed shows up mainly as a white precipitate (although some does dissolve). You get less precipitate as you go down the Group because more of the hydroxide dissolves in the water.
Reactions of the metals with oxygen
The group 2 elements all burn in oxygen to form white solid oxides.
2M(s) + O2(g) → 2MO(s)
Beryllium is reluctant to burn unless it is in the form of dust or powder. Beryllium has a very strong (but very thin) layer of
beryllium oxide on its surface, and this prevents any new oxygen getting at the to react with it.
Magnesium burns with a brilliant white flame. The others burn with characteristic flame colours:
brick red crimson red apple green
Reactions of the oxides with water
The metal oxides are all basic and each reacts with water, forming a solution of the hydroxide which is alkaline. The reaction of magnesium oxide with water is slow; the other oxides react readily with water.
The solubilities of the hydroxides formed increase down the Group and the solutions become more alkaline (approx. pH9 – 12).
Thermal stability of Group 2 carbonates and nitrates
Thermal decomposition is when a substance breaks down (decomposes) when heated. The more thermally stable a substance is, the more heat it will take to
You are only expected to describe the trend in thermal stability, not explain it.
Unit 3: Inorganic Chemistry
The hydroxides aren't very soluble, but they get more soluble as you go down the Group. The calcium hydroxide formed shows up mainly as a white precipitate
u go down the Group
The group 2 elements all burn in oxygen to form white solid oxides.
Beryllium is reluctant to burn unless it is in the form of dust or powder. Its Beryllium has a very strong (but very thin) layer of
beryllium oxide on its surface, and this prevents any new oxygen getting at the
Magnesium burns with a brilliant white flame. The others burn with
The metal oxides are all basic and each reacts with water, forming a solution of the hydroxide which is alkaline. The reaction of magnesium oxide with water is
ease down the Group and the
Thermal decomposition is when a substance breaks down (decomposes) when e more heat it will take to
You are only expected to describe the trend in thermal stability, not explain it.
AS Chemistry Unit 3: Inorganic Chemistry
Cambridge A-level Centre
The thermal stability increases down the group.
Group 2 carbonates decompose to from the oxide and carbon dioxide. Group 2 nitrates decompose to form the oxide, nitrogen dioxide and oxygen: Task 2 Can you write equations for the thermal decomposition of calcium carbonate and calcium nitrate? Uses Task 3 Compounds of the Group 2 elements have many uses. Use the information spread throughout your textbooks to answer the following questions.
1. Magnesium is the most commonly used metal in Group 2. Can you list 2
References A-level Chemistry pages 217-222 Chemistry in Context pages 213-219 Learning Objectives Candidates should be able to
• describe the reactions of the elements with oxygen and water. • describe the behaviour of the oxides with water. • describe the thermal decomposition of the nitrates and carbonates. • interpret, and make predictions from, the trends in chemical properties
of the elements and their compounds. • explain the use of magnesium oxide as a refractory lining material and
calcium carbonate as a building material. • describe the use of lime in agriculture.
