TOPIC 13: PERIODICITY
13. 1 TRENDS across PERIOD 3
13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides
and oxides of the elements in period 3 in terms of their bonding and structure.
13.1.2 Describe the reactions of chlorine and the chlorides referred to in 13.1.1 with water.
Oxides
Oxides
oxidation state
of period 3
element
empirical ratio
in terms of O :
state at room
temperature
electrical
conductivity
when molten
structure of
oxide
acid-base
character of
oxide
Na2O


MgO
giant ionic
basic

Al2 O3

amphoteric
P4O6
P4O10
SiO2
giant
covalent
SO2 /SO3
Cl2O
Cl2O7
 simple molecular 

acidic

You need to know all the oxides in the table + be able to determine the oxidation state of the period 3
elements in each oxide.
Physical trends: need to be explained by looking at their structures.
 Physical state: melting and boiling points: rise to a maximum at SiO2 and then decline;
 ionic compounds (Na2O, MgO and Al2O3) on the left side have relatively high melting and boiling
points (high lattice energy); solids at room temperature;
 giant covalent molecular structure like silicon dioxide has the highest melting and boiling points
(lattice held together by strongest bond i.e. covalent); solids at room temperature;
 lowest melting and boiling points are for the simple molecular structures (starting from P4O6); they
are gases at room temperature (weaker dipole-dipole attractions).
 Conductivity in molten state: changes from good to semi-conductor to poor; this is the case because
their structure changes from ionic to simple molecular (see table in topic 3);
Chemical trends.
 the type of bonding between element and oxygen changes from ionic to covalent; aluminium oxide is
still ionic but has quite a strong covalent character whilst silicon dioxide is mostly covalent;

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 type of oxide:
o
o
o
basic: Na2O
MgO
amphoteric: Al2 O3
acidic: SiO2
P4 O10
SO3 Cl2 O7
 some elements show variable oxidation states; e.g. P4 O10 and P4 O6 and SO3 and SO2;
 regular trend in empirical formula of highest oxide: ratio of element : oxygen (X : O) increases from 0.5
to 3.5.
 reactions with water:
 ionic oxides are soluble and react with water to form alkaline solutions (hydroxides) with a
decreasing pH when going to the right of the period (basic oxides);
Na2O (s) + H2O (l)
MgO (s) + H2O (l)
 2Na+ (aq) +
 Mg2+ (aq) +
2OH- (aq)
2OH- (aq)
/2NaOH (aq)
/Mg(OH)2 (aq)
 aluminium oxide is amphoteric:
Al2 O3 (s) +
Al2 O3 (s) +

6HCl
 2AlCl3 (aq) +
(aq)
3H2O (l)
 2NaAlO2 (aq) +
(aluminate ion)
2NaOH (aq)
H2O (l)
silicon dioxide does not dissolve in or react with water ( water remains neutral) but it can react with
sodium hydroxide which is why it is considered an acidic oxide;
SiO2 (s) + 2NaOH (aq)
 Na2SiO3 (s) +
H2O (l)
 simple molecular oxides are soluble in water and react to form strong acidic solutions like
phosphoric, sulphuric and chloric acid.
P4 O10 (s) +
SO2 (g) +
Cl2O +
6H2O (l)
H2O (l)
H2O
 4H3PO4 (aq)
 H2SO3 (aq)
 2HClO (aq)
Chlorides
chlorides
NaCl
MgCl2
Al2Cl6
SiCl4
PCl3
PCl5
Cl2
state at room
temperature
bonding
structure
electrical
conductivity when
molten
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Physical trends
 overall decrease in melting and boiling points as the structures change from ionic to simple molecular.
This happens between Mg and Al; metallic chlorides (NaCl, MgCl2, AlCl3 ) are solids, nonmetal/molecular chlorides are gases (covalent bonding between atoms within molecules but Van Der
Waal’s between molecules). No giant covalent chloride as silicon tetrachloride is a gas;
 conductance: good in molten ionic chlorides while poor in the simple molecular chlorides.
Chemical trends
 the type of bonding formed changes from ionic to covalent; aluminium trichloride however is now
covalently bonded instead of ionic;
 no chloride for argon as it does not form any compounds as it is a noble gas with full outer shell.
 there is a regular trend in the empirical formula of the highest chlorides because of the increase in
oxidation state from +1 to +5;
NaCl
MgCl2
Al2 Cl6 (=dimer)
SiCl4
PCl5
 some elements show variable oxidation states; e.g. PCl5 and PCl3
 reactions with water:
 ionic chlorides (NaCl, MgCl2) dissolve readily due to polar water molecules being able to pull
oppositely charged ions out of the lattice; NaCl only forms a neutral solution;
 magnesium chloride and aluminium trichloride form acidic solutions as their small highly charged
cations (high charge density) attract water molecules (to form complexes) some of which give up
their hydrogen ions to other water molecules surrounding the complex.
[Al (H2O)6]3+
+
H2O

[Al (H2O)5 ] OH2+
+
H3O+
In addition, AlCl3 also forms HCl when it reacts with water:
AlCl3 +
3H2O

Al (OH)3
+ 3HCl (=fumes) (exothermic reaction)
 simple molecular chlorides all react rapidly with water to form hydrochloric acid and oxyacids in
the case of S, P and Cl.
PCl5
+
3H2O

H3 PO4
+
5HCl (= HCl fumes)(exo)
SiCl4 + 4H2O  Si(OH)4 + 4HCl
Cl2 + H2O  HClO + HCl
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13. 2 First row d-block elements
13.2.1 List the characteristic properties of transition elements.
13.2.2 Explain why Sc and Zn are not considered to be transition elements.
We will focus our study on the first row which contains 10 elements because the 3rd sub-level contains 5
orbitals each accommodating two electrons.
These 10 elements, although they are in the same row and in different groups, show a lot of similarities but
also changes. However, as you move across the row these changes are only very gradual (=transitional)
as opposed to the more pronounced changes when you go across periods 2 and 3.
The reasons why these elements share properties so closely must be found in their electronic structures
and the relative energy levels within their atoms.
The following are characteristic properties of transition elements:




variable oxidation states
complex ion formation
coloured compounds
catalytic properties
To be able to explain these properties, we need to know the electronic configuration of these first row
transition metals. Complete the configuration using the Aufbau rules of the first 10 transition elements.
3d
Sc
[Ar]
Ti
[Ar]
V
[Ar]
Cr
[Ar]
Mn
[Ar]
Fe
[Ar]
Co
[Ar]
Ni
[Ar]
Cu
[Ar]
Zn
[Ar]
4s
 When filling up orbitals:
 Just like in the ‘s’ block elements in period 4, after the 3p orbital the next orbital to be filled is the
4s instead of the 3d.
 the 3d orbitals are being filled up in the first row of the Periodic Table; because of this they all
have similar outer electron arrangements which is why they have similar properties;

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exceptions to the general patterns in the electron arrangement are:
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 Cr: [Ar] 3d5 4s1 instead of [Ar] 3d4 4s2 as the first is the preferred arrangement and
appears to be more stable in terms of energy; the 3d sub-level is half-filled (a half-filled or
filled or empty orbital has greater stability than a partially filled orbital although this does
not always apply) which gives a lower total energy; because the 3d and 4s are so close
together they can be considered as degenerate orbitals.
 Cu: [Ar] 3d10 4s1 instead of [Ar] 3d9 4s2 because filling up of 3d is more stable
arrangement.
 You recall that when filling up orbitals, the 4s orbital is filled before the 3d orbitals. This is because in the
empty atom, 4s orbitals have a lower energy than 3d orbitals. However, once the electrons are actually
in their orbitals, the energy order changes - and in all the chemistry of the transition elements, the 4s
orbital behaves as the outermost, highest energy orbital. So the reversed order of the 3d and 4s orbitals
only applies to building the atom up in the first place. In all other respects, you treat the 4s electrons as
being the outer electrons.
 There is little decrease in atomic radius as you go across the rows of transition metals because as you
go across electrons are being added to the inner 3d orbitals in which the electrons are on average closer
to the nucleus than the 4s electrons. As a result, these 3d electrons shield the outer 4s electrons from
an increased nuclear charge. This cancels out any large decrease in atomic radius caused by an
increased nuclear charge. This also explains the similar first ionization energies of the transition
elements.
Increased nuclear charge is offset by the shielding effect of the 3d electrons.
Electronic configuration of ions
As explained earlier, after Sc, once the electrons are in their orbitals, the 3d electrons are on a slightly lower
energy level so when the transition metal atoms ionise, it are the 4s electrons (they are also shielded from
the nuclear charge by the 3d electrons) which are removed first before the 3d electrons.
Complete the table below
species
Sc3+
Ti4+
V5+
Cr3+
Cr6+
Mn2+
Mn3+
electronic configuration
species
Fe2+
Fe3+
Co2+
Ni2+
Cu+
Cu2+
Zn2+
electronic configuration
Why variable oxidation states (except scandium and zinc): similar ionization energies of 3d
and 4s electrons
13.2.3 Explain the existence of variable oxidation number in ions of transition elements.

oxidation state of a transition element in a compound or molecular ion = number of electrons
released/sharing in either covalent (sometimes the bonds formed have a greater covalent character) or
ionic bonds;
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



all transition elements , except Sc (+3 only) and Zn (+2 only), have more than one stable oxidation
state.
transition metals have variable oxidation states because the 3d and 4s electrons have similar
ionization energies (only a gradual increase) because the five inner d orbitals are at a similar energy
level as the single 4s orbital.
common oxidation state: as first and second ionization energies are very similar in all transition
elements, all first row d-block elements have + 2 as an oxidation state which corresponds with the loss
of the 4s electrons. Exceptions are Cr and Cu which lose the single 4s electron and one 3d electron to
have +2 as their oxidation state.
The only gradual increase in successive ionization energies is because there is less repulsion every
time an electron is removed.
Some d-block elements are not considered transition elements
A transition element is a d-block element which forms at least one stable ion
that has a partially filled d-sub-level.
As Zn2+ and Sc3+ do not have partially filled d-orbitals (with Zn2+ the 3d orbital is filled and with Sc3+ the 3d
orbital is empty), they are not considered transition metals. This is also the reason why they do not form
coloured compounds or show little catalytic activity.– see later.
You should be familiar with the following oxidation states (in brackets are their configurations):
Cu
Cr
Mn
Fe
+1
+3
+4
+3
(=[Ar] 3d10 )
(=[Ar] 3d3 )
(=[Ar] 3d3 )
(=[Ar] 3d5 )
+2
+6
+7
+2
(=[Ar] 3d 9 )
(=[Ar] )
(=[Ar] )
(=[Ar] 3d6 )

the higher oxidation states of V, Cr and Mn involve the formation of the [Ar] core; they show these high
oxidation states when bonded COVALENTLY with oxygen or fluorine (e.g. V2O5 , CrO3 , Mn2O7 ) or as
part of a complex ion with oxygen (CrO42-, MnO4-). Losing 7 electrons would require an amount of
energy which is equal to the total amount of all 7 successive energies. They do this with oxygen or
fluorine as they are very electronegative and are able to pull away a larger number of electrons;

The other elements to the right of Mn in the first row do not show such high oxidation states as in their
atoms the higher nuclear charge starts to count and starts pulling more strongly onto the 3d electrons
ensuring that really only the 4s electrons are easily removed. These molecules or oxyions containing Cr
or Mn are excellent oxidising agents as they are easily reduced.

Due to their high charge density chromium and manganese polarise anions (like O2- ) and the ionic bond
becomes more covalent so compounds like CrO3 and Mn2O7 are simple molecular and even acidic
solids.
Formation of complexes or complex ions (result of the high charge density)
13.2.4 Define the term ligand.
13.2.5 Describe and explain the formation of complexes of d-block elements.
A complex or complex ion has a metal ion at its centre around which there are a number of other molecules
or ions. These complexes/complex ions are usually formed when transition metals are dissolved in water; or
become hydrated. However, the transition metal ions also form complexes in other circumstances.
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When the complex is charged it is called a complex ion; the charge on the ion is delocalised over the entire
complex as indicated by the square brackets.
Transition metals can form complexes because their ions:


have a high charge density: they have quite a large nuclear charge but are relatively small;
the 3d electrons are not so effective (as 2s or 2p electrons) at shielding the effect of the ionic charge
which really comes from the nucleus.
This allows the transition metal ions to have a great polarising power and they can attract lone pairs from
other atoms to form complexes.
A complex is a compound which is formed when a donor or ligand, which has been attracted by the charge
of the transition metal ion, donates an electron pair (= dative bond) into an empty low energy orbital (3d, 4s
or 4p) of the metal ion. Such a compound is called a co-ordination compound.
Molecules like water, ammonia and negative molecular ions can all act as ligands as they have at least one
lone pair.
A ligand = a molecule or negative ion which contains a non-bonding electron
pair which it uses to form a dative bond with the central ion in a complex.
Co-ordination number
The number of ligands that are attached to such a metal ion is referred to as the co-ordination number.
This number is determined by the vacant orbitals (3d, 4s and 4p) which hybridise.
Common co-ordination numbers are:
 Cu = 4
 Fe = 6
 Ag = 2
Examples of complex ions
[Fe(H2O)6]3+
[Fe(CN)6]3-
[CuCl4]2-
[Cu(NH3)4]2+
[Ag(NH3)2]+
Shapes of complex ions
Complexes also have shapes which we can predict using the VSEPR theory. The shape depends on the
number of ligands:



if co-ordination number = 6
if co-ordination number = 4
if co-ordination number = 2
then shape = octahedral
then shape = tetrahedral (or square planar = less common)
then shape = linear
Examples of octahedral ions from http://www.chemguide.co.uk/inorganic/complexions/shapes.html
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Coloured compounds: a result of the complex formation.
13.2.6 Explain why some complexes of d-block elements are coloured.
A compound that contains electronic energy levels which are close together could absorb radiation in the
visible light spectrum and therefore display a colour.
This must is the case in d-block metals as their compounds are frequently coloured both in solid state
(hydrated) and in solution.
In d-block metal atoms the five 3 d orbitals have the same energy but in most of their compounds (usually
complexes), these 5 equivalent energy levels are split in two or three different sets of energy level. This
splitting is caused by the ligands as the electron clouds around the ligand repel the d orbitals; as a result
they push the orbitals closest to them to energy levels higher than the other energy levels which are not
near the ligand. As a result there are d sublevels of different energy – non-equivalent - and this process is
called field-splitting.
Although they are different in energy, these split energy levels are still close together. This allows an
electron in a lower 3d orbital to absorb radiation (from the visible light spectrum) and be promoted
(=transition) to a higher 3d orbital. The amount of energy needed for this transition (or the energy
difference between the split levels) corresponds to the frequencies of radiation within the visible light
spectrum. For this to happen, the ion must have partially-filed 3 d orbitals.
Compounds containing transition ions with empty d-orbitals (eg Ti4+ ) or full d orbitals (eg Zn2+ ) are
colourless as no transition between split levels can occur (in the case of full orbitals there are no spaces).
Remember that Zn and Sc are not considered transition metal; another reason is that they cannot form
coloured compounds. This is because in the only ion they form, there are either no 3d electrons e.g. in Sc3+
or there are no vacant orbitals eg Zn2+.
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The field splitting in a Cu2+ when it forms a complex with water is shown in the diagram below (from
http://www.chemguide.co.uk/inorganic/complexions/colour.html)
The colour of a transition metal complex depends on:

the number of d electrons in the metal ion; this effects the amount of repulsion and therefore the amount
of filed splitting. Transition metals can form differently coloured compounds as they take on different
oxidation states in these compounds which means they have a different number of 3d electrons
resulting in different field splitting.

number of ligands around the ion or coordination number; also effects how much field splitting goes on.

the nature of the ligand: each ligand has its own effect on the relative energies of the d electrons, e.g.
NH3 has a greater effect than water because it has a lower electronegativity and therefore attracts its
lone pair less strongly allowing it to repel more other electrons.
Catalysis
13.2.7 State examples of the catalytic action of transition elements and their compounds.
Catalysts provide pathways with lower activation energies allowing a chemical reaction to go faster.
Transition metals are good catalysts because:

they have the availability of 3d and 4s electrons which allows them to change easily between
different oxidation states (homogenous catalysis)

have empty orbitals which can be used to make temporary bonds (heterogenous catalysts).
Heterogenous catalysis
In heterogenous catalysis, the catalyst is in a different phase from the reactants e.g. usually in a solid state
with the reactants in liquid or gaseous state.
The transition metal catalysts provide reaction sites onto which the reactant gases are adsorbed (makes
weak temporary bonds); this lowers the activation energy of the reaction
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The transition metals are able to form such weak temporary bonds with a large range of ions and molecules
which have lone pairs or  electrons. These lone pairs and  electrons can occupy available space in the
empty orbitals in the transition metal atom.
Most transition atoms can make more than one temporary bond (more than one site!) which are broken as
soon as the products are made.
Examples of heterogenous catalysts:

Ni in the conversion of alkenes into alkanes.

Fe in the Haber process: nitrogen and hydrogen dissociate into atomic hydrogen and nitrogen on the
surface of the iron; the adsorption weakens the covalent bonds within the molecules producing more
reactive species.

MnO2 in the decomposition of hydrogen peroxide.

V2O5 in the Contac process.

Palladium and platinum in a catalytic converter.

Co in vitamin B12
Homogenous catalysis
A homogenous catalyst is in the same phase as the reactant. The catalyst does take part in the reaction as
it forms an intermediate which is regenerated in a later step of the reaction.
Examples:


vanadium oxide in the Contact process;
oxidation of iodide ions by peroxodisulphate ions; below is the equation of the reaction without a
catalyst:
S2O8 2- (aq) + 2I- (aq)
 2SO42- (aq) + I2 (s)
(1)
However, if iron (III) ions are added, the reaction is greatly speeded up as the following two steps occur:
Iron (III) ions oxidise the iodide ions producing iron (II) ions
2Fe 3+ (aq)
+
2I- (aq)

2Fe2+ (aq)
+
I2 (s)
(2)
the iron(II) ions are then converted back into iron (III) ions in the subsequent step:
2Fe 2+ (aq)
+
S2O8 2- (aq)

2Fe3+ (aq)
+
2SO42- (aq)
(3)
When combining equations (2) and (3) the original equation (1) is obtained. The iron (III) ions have
provided a lower energy path by which an electron transfer can take place; they could do this because they
can change their oxidation states twice and emerge unchanged.
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IB past paper questions
PAPER 1
1. (M07)
2. (M06) Which electrons are lost by an atom of iron when it forms the Fe3+
n?
A. One s orbital electron and two d orbital electrons
B. Two s orbital electrons and one d orbital electron
C. Three s orbital electrons
D. Three d orbital electrons
2. (N05) Which particles can act as ligands in complex ion formation?
I. Cl-
II. NH3
A. I and II only
III. H2O
B. I and III only
C. II and III only
D. I, II and III
3. (N05) Which factors lead to an element having a low value of first ionization energy?
I. large atomic radius
II. high number of occupied energy levels
III. high nuclear charge
A. I and II only
B. I and III only
C. II and III only
D. I, II and III
4. (N06) Which properties are typical of d-block elements?
I. complex ion formation
II. catalytic behaviour
III. colourless compounds
A. I and II only
B. I and III only
C. II and III only
D. I, II and III
5. Which of the following salts form coloured solutions when dissolved in water?
I.
A.
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ScCl3
I and II only
II.
B.
FeCl3
II and III only
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III.
C.
NiCl2
III and IV only
IV.
D.
ZnCl2
I, II, III and IV
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6. (M99) Based on melting points, the dividing line between ionic and covalent chlorides of the
elements Mg to S lies between
A. Mg and Al
B. Al and Si
C. Si and P
D. P and S
7. (M99) Which ion is colourless?
A. [Cr(H2O)6]3+
B. [Fe(CN)6]4-
C. [ Cu(NH3)4]2+
D. [Zn(H2O)4]2+
8. (M00) Which of the following chlorides give a neutral solution when added to water?
I.
A. I only
NaCl
II Al2Cl6
B. I and II only
III PCl3
C. II and III only
D. I, I and III
9. (M03) What is the electron configuration for an atom with Z = 22?
A. 1s2 2s2 2p6 3s2 3p6 3d4
B. 1s2 2s2 2p6 3s2 3p6 4s2 4p2
C. 1s2 2s2 2p6 3s2 3p6 3d2 4p2
10.
D. 1s2 2s2 2p6 3s2 3p6 4s2 3d2
Which is an essential feature of a ligand?
A.
a negative charge
B.
an odd number of electrons
C.
the presence of two or more atoms
D.
the presence of a non-bonding pair of electrons
11. (N00) Which aqueous complex ion will NOT be coloured?
A. Ni2+
B. Fe2+
C. Sc3+
D. Cr3+
12. (M01) The electronic configuration of chromium (Cr) is
A. 1s2 2s2 2p6 3s2 3p6 3d4 4s2
B. 1s2 2s2 2p6 3s2 3p6 3d5 4s1
C. 1s2 2s2 2p6 3s2 3p6 3d6
D. 1s2 2s2 2p6 3s2 3p6 3d14s2
13. Which could not act as a ligand in a complex ion of a d-block element?
A. Cl-
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B. NCl3
C. PCl3
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D. PCl5
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14. Which one of the following electron arrangements could be that of a d-block element in its ground
state?
A. 1s2 2s2 2p6 3s2 3p6 3d4
B. 1s2 2s2 2p6 3s2 3p6 3d1 4s1
C. 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p2
D. 1s2 2s2 2p6 3s2 3p6 3d4 4s2
15. The compounds Na2O, Al2O3 and SO2 respectively are
16.
A.
acidic, amphoteric and basic.
B.
amphoteric, basic and acidic.
C.
basic, acidic and amphoteric.
D.
basic, amphoteric and acidic.
Which general trends are correct for the oxides of the period 3 elements (Na2O to Cl2O)?
A.
17.
I.
Acid character decreases.
II.
Electrical conductivity (in the molten state) decreases.
III.
Bonding changes from ionic to covalent.
I and II only
B.
I and III only
C. II and III only
D.
I, II and III
Which of the following oxides is (are) gas(es) at room temperature?
I.
A.
SiO2
I only
II.
B.
P4O6
III only
III.
C.
SO2
I and II only
D. II and III only
PAPER 2
1. Elements with atomic number 21 to 30 are d-block elements.
(a) Identify which of these elements are not considered to be typical transition elements.
[1]
(b) Complex ions consist of a central metal ion surrounded by ligands. Define the term ligand. [2]
(c) Complete the table below to show the oxidation state of the transition element.
ion
oxidation state
[CuCl4]2−
Cr2O72–
[3]
[Fe(H2O)6]3+
(d) Identify two transition elements used as catalysts in industrial processes, stating the process
in each case.
[2]
(e) Apart from the formation of complex ions and apart from their use as catalysts, state two other
properties of transition elements.
[2]
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2. (N03)
(a) Describe the acid-base character of the oxides of the period 3 elements Na to Ar. For aluminium oxide,
write balanced equations to illustrate its acid-base character.
[2]
(b) Outline the reasoning for the following in terms of electronic configuration:
(i) The first ionization energy of Al is lower than that of Mg.
[2]
(ii) V3+ (aq) is coloured and can behave as a reducing agent, whereas Zn2+ (aq) is not coloured and
does not behave as a reducing agent.
[6]
(c) (i) Explain how the first ionization energy of K compares with that of Na and of Ar.
[3]
(ii) Explain the differences between the atomic radii and the first ionization energies of Na and Mg. [4]
(iii) Explain the difference between the second ionization energies of Na and Mg.
[2]
(d) Define the term ligand. Cu2+ (aq) reacts with ammonia to form the complex ion [Cu(NH3)4]2+.
[4]
Explain this reaction in terms of an acid-base theory, and outline the bonding in the complex ion formed
between Cu2+ and NH3.
3.
(N00)
(a) The first ionization energies of the elements Na to Ar are given in the data booklet.
[6]
(i) Account for the general increase for ionization energy across the period.
(ii) Explain why the first ionization energy of aluminium is less than that of magnesium.
(iii) Explain why the first ionization energy of sulfur is les than that of phosporus.
(b) List the formulae of the chlorides of Na, Mg, Al, Si and P. Why is there no chloride of argon? Give
the name of the bonding in the chloride of silicon both within and between the molecules.
[5]
(c) Classify the acid-base character of one oxide of each of the elements in the period from Na to S.
Illustrate your answer by writing a balanced chemical equation for the reaction of magnesium oxide
and phosphorus oxide with water.
[5]
(d) (i) Write balanced equations to show how aluminium oxide reacts with hydrochloric acid and with
sodium hydroxide.
[2]
(ii) Write a balanced equation to show what happens when FeCl3 is added to water.
[1]
(e) Describe and explain the redox reactions of Cl2, Br2 and I2 with Cl-, Br- and I- ions.
[5]
4. (a) (i) State the meaning of the term electronegativity and explain why the noble gases are not
assigned electronegativity values.
(ii) State and explain the trend in electronegativity across period 3 from Na to Cl.
(iii) Explain why Cl2 rather than Br2 would react more vigorously with a solution of I−.
Topic13
4 hours
[2]
[2]
[2]
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(b) State the acid-base properties of the following period 3 oxides.
MgO
Al2O3
P4O6
Write equations to demonstrate the acid-base properties of each compound.
[7]
5. (N00) This question concerns the atomic structure of iron.
(a) (i) Show the electron configurations (including electron spins) of iron and its positive ions, by filling
in the boxes below:
[3]
4s
3d
Fe
Fe2+
Fe3+
(ii) What is the oxidation state of iron in [Fe(CN)6]4-?
[1]
(iii) Iron can also exist in an oxidation state of +6. Give the formula of a species containing only
iron and oxygen in which Fe(VI) might exist.
[1]
(b) (i) Define the term ligand.
[1]
(ii) Explain why the two iron complex ions [Fe(H2O)6]2+ and [Fe(CN)6]4- have different colours? [2]
6. (N99) Iron shows the characteristic properties of a transition metal.
(a) Why does zinc not show the characteristic properties of a transition metal?
[1]
(b) Iron can form Fe2+ ions as well as Fe3+ ions. Give an example of another transition element that
can show variable oxidation states, showing clearly two of its oxidation states.
[1]
(c) Give the formula and describe the shape of the complex ion formed between Fe3+ and water.
[2]
(d) Explain why complexes of Sc3+ are colourless whereas complexes containing Fe3+ are coloured.
[2]
7. (N07)
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4 hours
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Mark scheme TOPIC 13
Paper 1
1
2
3
4
5
A
B
A
A
B
6
7
8
9
10
A
D
A
D
D
11
12
13
14
15
C
B
D
D
D
16
17
18
19
20
C
B
Paper 2
1. (N05)
(a) Sc and Zn
(b) A ligand = a molecule or negative ion which contains a non-bonding electron pair which it
uses to form a dative bond with the central ion in a complex.
(c) Cr = +6 Cu + 2 Fe = +3
(d) Ni in hydrogenation of alkenes
Fe in Haber process
(e) Coloured compounds/very high melting and boiling point/variable oxidation states
2. (N03)
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4 hours
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4 hours
17 of 21
3. (N00)
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4 hours
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4. (a) (i) Electronegativity is the ability for an atom to attract a bonding pair of electrons (a shared pair in a
diatomic covalent bond);
Noble gases do not form covalent bonds as they have full outer shells
(ii) Increases when moving to the right because;
nuclear charge increases/atomic radius decreases so the nucleus can attract other electrons
better;
(iii) Cl2 has larger electronegativity than Br2;
Smaller atomic radius/better at attracting electron from I(b) MgO (s) + H2O (l)  Mg2+ (aq) +
2OH- (aq) /Mg(OH)2 (aq);
Al2 O3 (s) +
6HCl (aq)  2AlCl3 (aq) + 3H2O (l) and Al2 O3 (s) +
 2NaAlO2 (aq) + H2O
P4 O10 (s) + 6H2O (l)  4H3PO4 (aq)
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4 hours
2NaOH (aq)
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5. (N00)
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4 hours
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6. (N99)
7. (N07)
Topic13
4 hours
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