Transcript Periodicity for HL

IB CHEMISTRY HL 1
UNIT 3 PERIODICITY
11th
IB t grade opics 3 and 13
3.1 THE PERIODIC TABLE
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3.1.1 Describe the arrangement of elements in the periodic
table in order of increasing atomic number. The history of
the periodic table will not be assessed.
3.1.2 Distinguish between the terms group and period.
3.1.3 Apply the relationship between the electron
arrangement of elements and their position in the periodic
table up to Z = 20.
3.1.4 Apply the relationship between the number of
electrons in the highest occupied energy level for and
element and its position in the periodic table.
The Periodic Table of Elements
3
THE PERIODIC TABLE
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The columns are called groups. The group number gives
the number of electrons in the valence shell.
The rows are called periods and these are labeled 1-7. The
period number gives the number of occupied electron
shells.
In the IB data booklet, the representative groups in the
Periodic Table are numbered from1 to 7 and the last
column is labeled as “0”.
THE PERIODIC TABLE
We can use the electron configuration to split up
the valence electrons into sub-levels.
Example: C is [He]2s22p2.
Note that valence electrons are in the same main
energy level.
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3.2 PHYSICAL PROPERTIES
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3.2.1 Define the first ionization energy and electronegativity.
3.2.2 Describe and explain the trends in atomic radii, ionic
radii, first ionization energies, electronegativities and melting
points for the alkali metals (Li  Cs) and the halogens (F  I).
Explanation for the first four trends should be given in terms
of the balance between the attraction of the nucleus for the
electrons and the repulsion between electrons. Explanations
based on effective nuclear charge are not required.
3.2.3 Describe and explain the trends in atomic radii, ionic
radii, first ionization energies and electronegativities for
elements across period 3.
3.2.4 Compare the relative electronegativity value of two or
more elements based on their positions in the periodic table.
EFFECTIVE NUCLEAR CHARGE
In any atom the nucleus exerts an attractive
force on the electrons.
 Across a period the number of protons in the
nucleus steadily increases. The effective charge
increases with the nuclear charge as there is no
change in the number of inner electrons.
 The effective nuclear charge experienced by an
atom’s outer electrons increases with the group
number of the element.
 It increases across a period but remains
approximately the same down a group.

Effective nuclear charge (Zeff) is the “positive charge” felt
by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Z
Core
Zeff
Radius (pm)
Na
11
10
1
186
Mg
12
10
2
160
Al
13
10
3
143
Si
14
10
4
132
8
Zeff  Z – number of inner or core electrons
ATOMIC RADIUS
 The
electron cloud does not have a sharp
boundary so atomic radius is usually
measured as half the distance between
two neighboring nuclei:
Atomic Radii
metallic radius
covalent radius
10
TRENDS IN ATOMIC RADII
Atomic radii increase down a group.
 Atomic radii decrease across a period.
 Going down a group there are more electron
shells so the atomic and ionic radii increase. The
effective nuclear charge remains about constant.
 Across period attraction between the nucleus
and the outer electrons increases as the nuclear
charge increases so electrons are pulled in more
and atomic and ionic radii decrease.

TRENDS IN ATOMIC AND IONIC RADII
TRENDS IN IONIC RADII
Positive ions are smaller than their parent
atoms. To form a positive ions the outer shell is
lost ex. Na is 2, 8, 1 whereas Na+ is 2, 8.
 Negative ions are larger than their parent atoms.
To form a negative ions electrons are added in
the outer shell ex. Cl is 2, 8, 7 and Cl- is 2, 8, 8.
There is increased electron-electron repulsion in
the outer shell so they move farther apart and
increase the radius of the outer shell.
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TRENDS IN IONIC RADII
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The ionic radii decrease from groups 1 to 4 for POSITIVE
ions. The ions Na+, Mg2+, Al3+ and Si4+ all have the same
electron arrangement: 2, 8. The decrease in ionic radius is
due to the increase in nuclear charge with atomic number
across the period. The increased attraction between the
nucleus and the electrons pulls the outer shell closer to the
nucleus.
The ionic radii decrease from groups 4 to 7 for the
NEGATIVE ions. The ions Si4-, P3-, S2- and Cl- have the
same electron arrangement 2, 8, 8. The decrease in ionic
radius is due to the increase in nuclear charge across the
period.
TRENDS IN IONIC RADII
The positive ions are smaller than the negative
ions, as the former have only two occupied
electron shells and the latter have three. This
explains the big difference between the ionic radii
of the Si4+ and Si4- ions and the discontinuity in
the middle of the table.
 The ionic radii increase down a group as the
number of electron shells increases.

Cation is always smaller than atom from
which it is formed.
Anion is always larger than atom from
which it is formed.
16
Isoelectronic: have the same number of electrons, and
hence the same ground-state electron configuration
Na+: [Ne]
Al3+: [Ne]
O2-: 1s22s22p6 or [Ne]
F-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
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IONIZATION ENERGY
The first ionization energy of an element is the
energy required to remove one mole of electrons
from one mole of gaseous atoms.
 Ionization energies increase across a period.
Number of protons increases across period 3 so
effective nuclear charge increases and ionization
energy increase with it.
 Ionization energies decrease down a group. Down
a group electrons are further from nucleus so
ionization energy decreases.

General Trends in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy
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TRENDS IN FIRST IONIZATION
ENERGIES
TRENDS IN IONIZATION ENERGIES
There are some small exceptions to the
increasing trend across a period:
 Ionization energy for a p sub-shell is lower than
for an s sub-shell. This is because p orbitals are
slightly higher in energy than s orbitals (in the
same period).
 There is also a decrease from the 5th element to
the sixth as the p sub-shells start to be doubly
filled.
 It is easier to remove the 6th electron as it is
repelled by its partner whereas the 5th electron is
not paired so it takes more energy to remove it.
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TRENDS IN IONIZATION ENERGIES
 Down
a group ionization energy decreases
as the outer electron is further from the
pull of the nucleus.
 Successive ionization energies for one
element increase (but not smoothly) due to
increased effective nuclear charge.
 When electrons are removed from a new
subshell there is a further increase in
ionization energy.
ELECTRONEGATIVITY
Electronegativity is the ability of an atom to
attract electrons in a covalent bond.
 Electronegativity is related to ionization energy
but is specific to BONDING electrons.
 Electronegativity increases from left to right
across a period owing to the increase in nuclear
charge, resulting in an increased attraction
between the nucleus and the bond electrons.
 Electronegativity decreases down a group. The
bond electrons are furthest from the nucleus and
so there is reduced attraction.

ELECTRONEGATIVITY
 Maximum
value is 4.0 which Fluorine has.
 Minimum value is 0.7 which Francium
has.
The Electronegativities of Common Elements
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MELTING POINTS
MELTING POINTS
Melting points of alkali metals (group 1) decrease
down the group as the metallic bonds weaken –
valence electrons are further from the nucleus so
the attraction between the delocalized electrons
and the positive ions decreases.
 Melting points of halogens (group 7) increase
down a group as van der Waals forces increase
with molar mass. The halogens all exist as
diatomic molecules in their standard elemental
form.
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MELTING POINTS
Melting points will increase with stronger
bonding and intermolecular forces. It is a
measure of the difference in forces between the
solid and liquid states.
 Boiling point is a measure of the absolute size of
these forces.
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MELTING POINTS ACROSS A PERIOD
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Across a period the bonding changes from metallic
(strong) to giant covalent (very strong) to van der
Waals forces between molecules (weak).
Melting points generally increase across a period and
reach a maximum at group 4.
The melting points increase and then decrease
accordingly with the changes in strength of bonding.
The bonding changes from metallic (Na, Mg and Al) to
giant covalent (Si) to weak van der Waals forces
between molecules (P4, S8 and Cl2) and single atoms
(Ar). All the period 3 elements are solids at room
temperature except chlorine and argon.
SUMMARY OF TRENDS ACROSS PERIOD 3
3.3 CHEMICAL PROPERTIES
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3.3.1 Discuss the similarities and differences in the
chemical properties of elements in the same group.
The following reactions should be covered: Alkali
metals (Li, Na and K) with water; Alkali metals (Li,
Na and K) with halogens (Cl2, Br2, I2); Halogens
(Cl2, Br2, I2) with halide ions (Cl-, Br-, I-).
CHEMICAL PROPERTIES
Chemical properties of an element are largely
dependent on the number of electrons in the
outer shell.
 This means that groups tend to have similar
chemical properties - they react in a similar way.
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THE NOBLE GASES, GROUP 0
These are the least reactive elements.
 They are monatomic – exist as single atoms.
 They are colorless gases.
 They have complete outer shells of electrons so
have the highest ionization energies for each
period.
 Other elements tend to react to attain the
electron configuration of the noble gases.
 Compounds of xenon, krypton and argon have
been made but it requires special conditions to
create these.
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THE ALKALI METALS, GROUP 1
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Physical properties:
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Good conductors of
electricity due to
delocalized valence
electrons
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Low densities
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Soft
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Grey shiny surfaces
when freshly cut
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Chemical properties:

Very reactive due to single
valence electron that is lost
easily
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Always form 1+ ions and
combine easily with nonmetals such as oxygen and
halogens.

Ex. 2Na(s) + Cl2(g) 
2NaCl(s)
THE ALKALI METALS, GROUP 1
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Reactivity increases down the group as the
valence electron is further from the attraction of
the nucleus and ionization energy decreases.
All alkali metals react vigorously with water to
form a metal hydroxide solution (basic) and
hydrogen gas
2Na(s) + 2H2O(l)  2Na+(aq) + 2OH-(aq) + H2(g)
All alkali metals tarnish quickly in air so they
lose their shiny surface. They are stored under oil
to prevent this.
ALKALI METALS STORED UNDER OIL
ALKALI METAL + WATER
Lithium floats and reacts slowly. It releases
hydrogen but keeps its shape.
 Sodium reacts with a vigorous release of
hydrogen. The heat produced is sufficient to melt
the unreacted metal, which moves around on the
surface of the water.
 Potassium reacts even more vigorously to
produce sufficient heat to ignite the hydrogen
produced. It produces a lilac colored flame and
moves excitedly on the water surface.

THE HALOGENS, GROUP 7
F, Cl, Br and I are very reactive non-metals in
group 7.
 All require one electron to complete their valence
shell.
 All exist as diatomic molecules joined by covalent
bonds ex. F2, Cl2, Br2, I2
 Van der Waals forces between the molecules
increase down the group with molar mass.
 They are all quite electronegative with F being
the most electronegative element (smallest
atomic radius).
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THE HALOGENS, GROUP 7
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Physical properties:
They are colored
They show a gradual
change from gases (F2
and Cl2) to liquid (Br2)
to solids (I2 and At2).
Chemical Properties:
Very reactive nonmetals. Reactivity
decreases down the
group.
They form ionic
compounds with
metals or covalent
compounds with nonmetals.
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THE HALOGENS, ALL TOXIC!
THE HALOGENS
 At
room temperature, F (pale yellow) and
Cl (yellow-green) are gases, Br is a redbrown liquid and I is a solid that forms a
black-purple vapor on heating, brown
solution in water and purple solution in
non-polar solvents.
 Gain an electron easily to form Hal Ease of gaining an electron (and
reactivity) decreases down the group as
electrons are further from nucleus.
 Slightly soluble in water as non-polar.
HALOGEN + ALKALI METAL
Halogens react easily with alkali metals to form
ionic halides.
 One electron is transferred from the alkali metal
to the halogen so that the alkali metal forms a 1+
ion and the halogen forms a 1- ion.
 These oppositely charged ions are strongly
attracted to each other and form a strong ionic
bond.
 The most vigorous reaction will occur between
the elements which are furthest apart in the
periodic table: francium at the bottom of group 1
and fluorine at the top of group 7.
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REACTIONS OF HALOGENS
Ex. 2Fr(s) + F2(g)  2FrF(s)
 The relative reactivity of the halogens can be
seen by combining a halogen element with a
metal halide:
 2KBr(s) + Cl2(aq)  2KCl(aq) + Br2(aq)
 Chlorine is more reactive than bromine so it can
displace bromine from the compound. The net
ionic equation could also show this:
 2Br-(aq) + Cl2(aq)  2Cl-(aq) + Br2(aq)
 The reverse reaction would not occur as bromine
is less reactive and cannot displace chlorine.
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HALOGEN + HALIDE
If the Cl2 is reacted with either the Br- or I- ions
then Br2 or I2 will be formed, respectively.
 If this is done in aqueous solution then with both
Br2 and I2 an orange-brown color will appear
from an originally colorless solution.
 The halogens can be distinguished more clearly
in non-polar solvents where they have the
following colors: chlorine is a pale green, bromine
is orange and iodine is violet.
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CL2, BR2 AND I2 IN CYCLOHEXANE
SILVER HALIDES
 Halogens
form insoluble salts with silver
and lead.
 Common test for halide ions is to add
nitric acid followed by aqueous silver
nitrate.
 A precipitate confirms presence of halide:
AgCl is white but darkens in sunlight
 AgBr is cream
 AgI is pale yellow
 AgF is soluble so this test wouldn’t work for F-.

AGI , AGBR, AGCL, AGF
SUMMARY OF Ag+ + Hal-
F-
C l-
B r-
I-
White prec ipitate,
P ale yellow
A queous
turns blac k in
C ream prec ipitate
N o Reac tion
prec ipitate
A g+
s unlight
A g+ A g+ + B r- - > A gB r
+ I - -> A gI
+ C l- -> A gC l
A g+
NR
S olution turns yellowS olution goes yellow
then brown C l2 + 2 Bthen
rblac k prec ipitate
- > B R2 + 2 C lC l2 + 2 I - -> I 2 + 2 C l-
B romine N R
NR
NR
S olution goes yellow
then blac k prec ipitate
B r2 + 2 I - -> I 2 + 2 B r-
I odine
NR
NR
NR
C hlorine N R
NR
13.1 TRENDS ACROSS PERIOD 3
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3.3.2 Discuss the changes in nature from ionic to
covalent and from basic to acidic of the oxides across
period 3. Equations are required for the reactions of
Na2O, MgO, P4O10 and SO3 with water.
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.
Include the following oxides: Na2O, MgO, Al2O3, SiO2,
P4O6 and P4 O10; and the following chlorides: NaCl,
MgCl2, Al2Cl6, SiCl4, PCl3, PCl5 and Cl2.
GROUP 1 AND 2 OXIDES ARE BASIC
Across period 3 the nature of the elements changes.
 Na and Mg form cations so they bond with O2- to
form ionic oxides.
 The oxide ion can bond with H+ ions so they act as
bases dissolving in water to give alkaline solutions.
 Na2O(s) + H2O(l)  2Na+(aq) + 2OH-(aq)
 They will also neutralize acids to produce salt and
water.
 MgO(s) + 2HCl(aq)  Mg2+(aq) + 2Cl-(aq)

AMPHOTERIC ALUMINUM OXIDE
Aluminum oxide does not dissolve in water easily
but it is AMPHOTERIC which means it will react
with (and dissolve in) acids and bases.
 Acting like a base:
 Al2O3(s) + 6H+(aq)  2Al3+(aq) + 3H2O(l)
 Al2O3(s) + 3H2SO4(aq)  Al2(SO4)3(aq) + 3H2O(l)
 Acting like an acid:
 Al2O3 (s) + 3H2O(l) + 2OH-(aq)  2Al(OH)4-(aq)
 Al2O3(s) + 2OH-(aq)  3H2O(l) + 2Al(OH)4-(aq)

ACIDIC OXIDES
The remaining oxides of period 3 (Si – Cl) form
acidic solutions.
 Silicon dioxide has little acid-base activity but it
shows weakly acidic properties by slowly
dissolving in hot concentrated alkalis to form
silicates.
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
SiO2(s) + 2OH-(aq) → SiO32-(aq) + H2O(l)
ACIDIC OXIDES
 Phosphorus
(V) oxide reacts to form a
solution of phosphoric (V) acid, a weak
acid:
 P4O10(s) + 6H2O(l) → 4H+(aq) + 4H2PO4(aq)
 Phosporus (III) oxide reacts with water to
produce phosphoric (III) acid:
 P4O6(s) + H2O(l)  4H3PO3(aq)
ACIDIC OXIDES
Sulfur trioxide reacts with water to make sulfuric
acid:
 SO3(l) + H2O(l)  H2SO4(aq)
 Sulfur dioxide reacts with water to produce
sulfurous acid:
 SO2(g) + H2O(l)  H2SO3(aq)

Cl2O7 reacts with water to produce perchloric
acid:
 Cl2O7(l) + H2O(l)  2HClO4(aq)
 Cl2O reacts with water to produce chlorous acid:
 Cl2O(l) + H2O(l)  2HClO(aq)

LEARNING CHECK
The reactivity increases in what order?
A. Na, K, Li
B. K, Na, Li
C. Li, Na, K
D. Li, K, Na
Give the colors of the following:
2.
Iodine vapor
3.
Color of precipitate when BaCl2 and
AgNO3 react
4.
Color of precipitate from 3 when left in
the sunlight
1.
DO NOW
The reactivity increases in what order?
A. Na, K, Li
B. K, Na, Li
C. Li, Na, K
D. Li, K, Na
C
1.
Give the colors of the following:
2.
Iodine vapor
purple
3.
Color of precipitate when BaCl2 and
AgNO3 react
white
4.
Color of precipitate from 3 when left in
the sunlight
black
THE OXIDES OF PERIOD 3
Across a period the number of valence electrons
increases so there are more electrons that can
form bonds with oxygen.
 Across period 3 each element bonds with an extra
half an oxygen - Na2O, MgO, Al2O3, SiO2, P4O10
(like P2O5), SO3, Cl2O7.

THE OXIDES OF PERIOD 3
The elements on the right of period 3 often form
more than one oxide so they exist in different
oxidation states in these elements.
 Phosphorus can form P4O6 and P4O10 where it has
an oxidation state of +3 and +5, respectively.

BONDING, MELTING AND BOILING POINTS
Na and Mg form ionic oxides so they are solids at
room temperature and have high mp’s and bp’s.
 SiO2 forms a giant covalent lattice so the mp and
bp are very high.
 The elements on the right form covalent
molecules so mp’s and bps are lower and they
exist as gases, liquids or low melting solids.

ELECTRICAL CONDUCTIVITY
The ionic compounds (Na2O and MgO) conduct
electricity when molten (liquid) as the ions can
move through the liquid.
 Aluminum oxide has ionic and covalent
characteristics so it is a poor conductor but has
an extremely high mp.
 The oxides of the non-metals do not conduct
electricity.

SUMMARY OF OXIDES OF PERIOD 3
Formula
Ratio of
Atoms
Bonding
Na 2O
2:1
MgO
2:2
Ionic
Ionic
AcidBasic
Base
Character
Other
Na 2O2
oxides
Al2O3
2:3
SiO2
2:4
P2O5
2:5
SO3
2:6
Cl2O7
2:7
Highly
Polar
Polar
Polar
Polar
polar
covalent covalent Covalent covalent
covalent
Basic Amphoteric Weakly Acidic
Acidic
Acidic
acidic
P2O3
SO2
ClO2 &
Cl2O
13.1 CHLORIDES OF PERIOD 3
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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.
Include the following oxides: Na2O, MgO, Al2O3, SiO2,
P4O6 and P4 O10; and the following chlorides: NaCl,
MgCl2, Al2Cl6, SiCl4, PCl3, PCl5 and Cl2.
13.1.2 Describe the reactions of chlorine and the
chlorides referred to in 13.1.1 with water.
THE CHLORIDES OF PERIOD 3
Across period 3 the elements bond to one more
chlorine - NaCl, MgCl2, AlCl3, SiCl4 and PCl5.
 On the right of the period the elements can exist
in different oxidation states ex. PCl3 also exists.

CHLORIDES OF PERIOD 3
Formula
of
chloride
NaCl
(s)
MgCl2 (s) AlCl3(s) / SiCl4(l)
Al2Cl6(g)
PCl5(s)
/
PCl3(l)
S2Cl2(l Cl2(g)
)
Oxidation
number
+1
+2
+3
+4
+5/+3
+1
0
High
Poor
None
None
None
None
Electrical
High
conductivit
y in molten
state
Structure
Giant ionic
Molecular covalent
GROUP 1 AND 2 CHLORIDES
The ionic compounds, NaCl and MgCl2, are ionic
crystalline solids with high melting points.
 NaCl dissolves in water to form a neutral
solution:
 NaCl(s)  Na+(aq) + Cl-(aq)
 MgCl2 dissolves to form a slightly acidic solution:
 MgCl2(s)  Mg2+(aq) + 2Cl-(aq)
 The resulting solutions can conduct electricity
due to the free moving ions.

ALUMINUM CHLORIDE
 Despite
being a metal, aluminum’s
compounds often behave more like nonmetals.
 This is due to the small size and high
charge of its ion.
 AlCl3 sublimes at 178°C to form Al2Cl6
molecules.
 AlCl3 dissociates into ions when added to
water:
 AlCl3(s)  Al3+(aq) + 3Cl-(aq)
ALUMINUM CHLORIDE
The aluminum ion is small and has a high charge
(3+) thus it has a high charge density.
 This means it attracts water molecules when in
solution and forms the complex ion: [Al(H2O)6]3+

ALUMINUM CHLORIDE
The ion is said to be hydrated and behaves as an
acid be releasing H+ from one of the H2O
molecules:
 [Al(H2O)6]3+(aq)  [Al(H2O)5OH]2+(aq) +
H+(aq)
 Further proton loss can occur:
 [Al(H2O)5OH]2+(aq)  [Al(H2O)4OH2+(aq) +
H+(aq)
 The solution is acidic enough to react with a
weak base and produce CO2(g):
 2AlCl3(aq) + 3Na2CO3(s)  3CO2(g) + Al2O3(s) +
6NaCl(aq)

SILICON CHLORIDE
Unlike the oxides the Si doesn’t form giant
covalent structures as Cl usually only forms one
bond.
 The chlorides of non-metals have low mp’s as
there are weak intermolecular forces between the
molecules.
 They react with water to form an acidic solution
containing H+, Cl-, O2- or an oxyacid of the
element (hydrolysis reaction):
 SiCl4(l) + 2H2O(l)  SiO2(s) + 4HCl(aq)

PHOSPHORUS CHLORIDES
PCl3 produces phosporous acid and hydrochloric
acid:
 PCl3(l) + 3H2O(l)  H3PO3(aq) + 3HCl

PCl5 produces phosphoric acid and hydrochloric
acid:
 PCl5(s) + 4H2O(l)  H3PO4(aq) + 5HCl(aq)

CHLORINE AND WATER
 In
water, Cl2 reacts slowly in a reversible
reaction to make a mixture of HCl and
HOCl acids:
Cl2(aq) + H2O(l)  HCl(aq) + HOCl(aq)
 This is disproportionation reaction where
Cl2 is reduced to HCl and oxidized to
HOCl (we’ll see this again in the unit on
redox)
CHLORINE AND WATER

The test for Cl2 uses this reaction: it turns litmus
paper from blue to red due to the HCl and then
colorless due to the bleaching power of HOCl.
THE HALOGENS
Chloric acid and ClO- are used as bleaches (ex.
For paper)
 They are also toxic to microbes so are used as
disinfectants and in water treatment.
 Halogens form ionic bonds with metals to make
salts containing a halide ion. These salts are
usually white and soluble in water ex. NaCl.

SUMMARY OF OXIDES AND
CHLORIDES
Ele m ent
Bonding
Na
Mg
Metaliic
Al
Si
Giant
covalent
Chloride
Formula
NaCl
MgCl2
AlCl3
SiCl4
Bonding
Ionic
Oxide
Formula
Na 2O
MgO
Al2O3
SiO2
Bonding
Ionic
Acid/base
properties
Sol.
Basic
Insol.
Basic
Intermediate Giant
covalent
Amphoteric Insol.
Acidic
P
S
Cl
Molecular Covalent
PCl5,
Complex Cl2
PCl3
Intermediate Molecular covalent
P4O10,, SO3,
Cl2O7,
P4O6
SO2
Cl2O
Molecular covalent
Soluble acidic
PRACTICE QUESTIONS
Which of the following doesn’t follow the
periodicity trend across period 3?
A. Al2O3
B. Na2O
C. SO2
D. P4O10
1.
Which of the following would cause a
reaction (could be more than one)?
A. Chlorine and sodium bromide
B. Bromine and potassium fluoride
C. Bromine and calcium iodide
D. Iodine and magnesium bromide
2.
PRACTICE QUESTIONS
Which of the following doesn’t follow the
periodicity trend across period 3?
A. Al2O3
B. Na2O
C. SO2
D. P4O10
C
1.
Which of the following would cause a
reaction (it might be more than one)?
A. Chlorine and sodium bromide
B. Bromine and potassium fluoride
C. Bromine and calcium iodide
D. Iodine and magnesium bromide A & C
2.
13.2 FIRST-ROW D-BLOCK ELEMENTS
13.2.1 List the characteristic properties of
transition elements. Examples should include
variable oxidation number, complex ion
formation, existence of colored compounds and
catalytic properties.
 13.2.2 Explain why Sc and Zn are not considered
to be transition elements.
 13.2.3 Explain the existence of variable oxidation
number in ions of transition elements. Students
should know that all transition elements can
show an oxidation number of +2. In addition,
they should be familiar with the oxidation
numbers of the following: Cr (+3, +6), Mn (+4,
+7), Fe (+3) and Cu (+1).

FIRST ROW D-BLOCK ELEMENTS
 3d
spans from Scandium to Zinc.
 The d-block does not follow the periodic
patterns of the s and p blocks – they all
have similar physical and chemical
properties.
 Transition elements are a subset of the dblock that have a partially filled dsublevel in one of its common oxidation
states.
 d-block elements are dense, hard metallic
elements.
PHYSICAL PROPERTIES

Typical physical properties of transition elements are:
 High electrical and thermal conductivity
 High melting point
 Malleable – easily beaten into shape
 High tensile strength – can hold large loads
 Ductile – easily drawn into wires
 These properties are all explained by the strong
metallic bonding. The 3d and 4s electrons are all
delocalized and form a strong attraction to the
positive ions. The large number of delocalized
electrons accounts for the high electrical
conductivity and higher density than group 1 and 2
metals.
CHEMICAL PROPERTIES

Typical chemical properties of transition
elements are:
Variety of stable oxidation states (just means ions
with different charges)
 Ability to form complex ions
 Formation of colored compounds
 Catalytic activity as either elements or compounds

ELECTRON CONFIGURATIONS
In most of the 3d elements 4s is filled and the
number of electrons in 3d varies from one
element to the next.
 In Cr and Cu there is only 1 electron in 4s so that
there will be more unpaired electrons in 3d - this
increases stability.
 When any of the 3d elements form positive ions
the 4s electrons are removed first.

OXIDATION STATES OF 3D
The metals in 3d can lose different number of
electrons to form different ions.
 These ions are all said to be in different oxidation
states.
 The oxidation state (oxidation number) is the
same as the charge on the ion,
ex. Cr3+ has an oxidation state of +3; Cr2+ has an
oxidation state of +2.

OXIDATION STATES OF 3D




The 3d electrons shield the 4s electrons so the first
ionization energy is relatively constant across the
period giving the elements similar properties.
From left to right effective nuclear charge increases so
the maximum oxidation state is most stable for the
elements on the left of 3d (Sc – Mn).
The maximum oxidation state means all 3d and 4s
electrons are lost.
The +2 state is most stable for elements on the right
(Fe – Zn)
SCANDIUM AND ZINC
Sc and Zn don’t share all the properties of
transition elements as they don’t have a partially
filled d block.
 Zn always forms 2+ ions, it loses the 4s2 electrons
and keeps the 3d full.
 Sc always forms 3+ ions, it loses all its valence
electrons, 4s2 and 3d1.

FIRST ROW D-BLOCK ELEMENTS
s-block metals lose s electrons easily but the
ionization energies for the inner electrons are so
high that these are never lost.
 For this reason they always have the same
oxidation state - a +1 ion has oxidation number
+1.
 Transition metals have slightly higher effective
nuclear charge so first ionization energies are
higher but there is no sudden increase in
successive ionization energies.

FIRST ROW D-BLOCK ELEMENTS
The sudden increase in ionization energies occurs
only once all the 3d and 4s electrons have been
removed.
 The oxidation state of transition elements varies
depending on how strongly oxidizing the
environment is.
 This depends on the presence of a species that
readily gains electrons.

FIRST ROW D-BLOCK ELEMENTS
Element
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
[Ar]
[Ar]3 [Ar] [Ar] [Ar]3 [Ar] [Ar] [Ar]3 [Ar] [Ar]3
Electronic
3d104
d1 3d24 3d3 d5 3d5 3d6 d7 3d8 d10
Structure
s2
4s2
s2 4s2 4s1 4s2 4s2 4s2 4s2 4s1
Decreasing stability of maximum oxidation state ------->
Increasing stability of +2 oxidation state --------->
Sc Ti
VOXIDATION
Cr
MnSTATES
Fe
COMMON
Ni
Cu
Zn
MnO4-
+7
CrO42Cr2O72-
+6
MnO42-
VO2+
VO3-
+5
+4
+3
Co
Sc3
Ti4+
VO2+
Ti3+
V3+
Cr3+
Ti2+
V2+
Cr2+
MnO2
Fe3+
+
+2
Mn2+
Fe2+
Co2 Ni2+ Cu2+ Zn2+
+
+1
Cu+
Increasing stability of +2 state -------------------------------------
-----------------------------------Increasing stability of maximum state
FIRST ROW D-BLOCK ELEMENTS
The stability of the half filled 3d level - as seen in
Cr and Cu - also affects the stability of oxidation
states.
 In Mn the +2 state which has a half filled state is
much more stable than +3 or +4 - these are quite
strong oxidants.
 With iron the +3 is most stable as it has the half
filled 3d shell and the +2 state is quite strongly
reducing.
 In Cu +1 exists as it has a full 3d sub-shell and
like Zn2+ its compounds are not colored.

UNUSUAL ION CONFIGURATIONS
Fe2+ : [Ar]3d54s1
 Co2+: [Ar]3d54s2
 Having a half-filled d-block gives stability so
sometimes the 4s electrons are not all lost first.

13.2 FIRST-ROW 3-D BLOCK ELEMENTS
 13.2.4 Define the term ligand.
 13.2.5 Describe and explain the formation of complexes
of d-block elements. Include [Fe(H2O)6]3+, [Fe(CN)6]3-,
[CuCl4]2- and [Ag(NH3)2]+.
 13.2.6 Explain why some complexes of d-block
elements are colored. In complexes, the d sub-level
splits into two sets of orbitals of different energy and
the electronic transitions that take place between
them are responsible for their colors.
COMPLEX IONS
The ions of d-block metals and those in the lower
section of the p-block (like lead) have unfilled
valence d and p orbitals.
 These orbitals can accept a lone pair of electrons
from species, known as ligands, to form a dative
covalent bond between the ligand and the metal
ion.
 Ex. An NH3 molecule can donate its non-bonding
electron pair to a Cu2+ ion.

COMPLEX IONS
 This
behavior where one species donates a
pair of electrons and another accepts is
Lewis acid-base behavior
 Species that have ligands bonded to a
central metal atom are known as
COMPLEX IONS.
Ex. [Cu(NH3)4]2+ forms when excess
ammonia is added to a solution of a copper
(II) salt.
 The charge is the sum of the metal ion
charge and the charges on the ligands.
COMPLEX IONS
LIGANDS are species that can donate a lone pair
of electrons to a metal ion.
 The most common examples are water, ammonia
(NH3), chloride ion and cyanide ion (CN-).
 Most complex ions have either 6, 4 or 2 ligands.
 The number of ligands is the COORDINATION
NUMBER of the metal ion.

COMPLEX IONS
2
ligands form a
linear complex.
4
ligands usually
form a tetrahedral
shape but can be
square planar.
 6 ligands usually
form an octahedral
shape.
COMPLEX IONS
 Complex
ions can have a positive or negative
charge and can form salts with ions of the
opposite charge - they are soluble and the
solution conducts electricity.
 Some complexes are neutral because the
charges cancel - these are insoluble.
Complex
ion
Charge on Oxidation
complex
state on metal
ion
ion
Similar to
[Cu(NH3)4Cl2]
[Cu(NH3)4]2+
2+
+2
CaCl2
[K2(CuCl2)]
(CuCl4)2-
2-
+2
K2SO4
COMPLEX IONS
The formation of complex ions stabilizes certain
oxidation states.
 The formation of a complex ion can also affect the
color of a metal ion in solution.
 For many complexes, ligand replacement can
occur depending on which complex is more stable.

EXAMPLES OF COMPLEX IONS
Metal ion
Water
Octahedral
Ammonia
Octahedral /
Square Planar
Chloride ion
Tetrahedral
Cobalt(II)
Pink
[Co(H2O)6]2+
Straw
[Co(NH3)6]2+
Blue
[CoCl4]2-
Nickel(II)
Green
[Ni(H2O)6]2+
Blue
[Ni(NH3)6]2+
Yellow-green
[NiCl4]2-
Copper(II)
Blue
[Co(H2O)6]2+
Deep blue
[Cu(NH3)4]2+
Yellow
[CuCl4]2-
COMPLEX IONS
Complex ions exhibit ISOMERISM in a similar
way to organic compounds.
 There are 3 types of chromium (III) chloride
hexahydrate that vary as shown below:

COMPLEX IONS
STEREOISOMERISM also occurs in complex
ions.
 Ex. Pt(NH3)2Cl2 has a square planar shape but
may occur in a cis or trans form:

COMPLEX IONS
Cis means the ligands are on the same side.
 Trans means the ligands are on opposite sides.

COLORED IONS
Usually the d orbitals in an atom have equal
energy.
 When an atom has ionic or polar ligands around
it the d orbitals are often split into 2 groups, one
with higher energy than the other.
 The difference between these levels corresponds
to a frequency of light in the visible region.

COLORED IONS
 If
white light passes through the complex
ion colored light is absorbed, electrons are
excited to the higher d orbitals and the
opposite color is seen.
Example: Most copper (II) compounds
absorb red and yellow so we see bluegreen color.
 If there are no electrons in the d orbitals
like in Sc3+ and Ti4+ then the compounds
are colorless.
 If the d orbitals are full as in Zn2+ then
the compounds are also colorless.
ENERGY OF LIGHT

The difference in energy level between the 2 sets
of d orbitals depends on the following:
Nuclear charge and the identity of the metal ion
 Charge density of the ligand
 Number of d electrons present and hence the
oxidation number of the central ion
 Shape of the complex ion

13.2 FIRST-ROW 3-D BLOCK ELEMENTS
13.2.7 State examples of the catalytic action of
transition elements and their compounds.
 Examples should include:








MnO2 in the decomposition of hydrogen peroxide.
V2O5 in the Contact process.
Fe in the Haber process and in heme.
Ni in the conversion of alkenes to alkanes.
Co in vitamin B12.
Pd and Pt in catalytic converters.
Mechanism of action will not be assessed.
CATALYSTS
A catalyst enables a reaction to happen by
providing an alternative pathway with a lower
activation energy. It is not used up or changed in
the reaction so it does not appear in the chemical
equation.
 Transition metals act as catalysts easily because
they can form complex ions resulting in close
contact with ligands.
 The number of stable oxidation states also means
they can gain and lose electrons easily in redox
reactions.

HETEROGENEOUS CATALYSTS
A heterogeneous catalyst is in a difference state
from the reaction. Ex. a solid catalyst with
gaseous reactants.
 Heterogeneous catalysts are more common than
homogeneous catalysts.
 A heterogeneous catalyst provides an active
surface where the reaction can occur, ex. Solid
MnO2 catalyses the decomposition of hydrogen
peroxide:
 2H2O2(aq) → 2H2O(l) + O2(g)

CATALYTIC CONVERTERS
 Platinum
and palladium are found in the
catalytic converters in car exhaust
systems where they help to reduce the
emission of CO and NO.
 2CO + 2NO → 2CO2 + N2
 Many important industrial catalysts
involve transition elements.
 The economic importance of the chemical
industry rests on the food, clothes,
medicines and other varied products that
it makes.
HABER PROCESS
The chemical industry is a sign of development of
a country as it converts simple cheap raw
materials into more useful and valuable
substances.
 The Haber Process uses iron as a catalyst to
convert the “free” nitrogen from the atmosphere
to make ammonia and then explosives (on which
wars depend), fertilizers (helps grow food crops)
and polymers such as nylon.
 N2(g) + 3H2(g)  2NH3(g)

CONTACT PROCESS
The Contact Process uses vanadium (V) oxide
(V2O5) as a catalyst to convert sulfur dioxide to
sulfur trioxide.
 2SO3(g) + O2(g) → 2SO3(g)
 SO3 is used to make sulfuric acid, the “king of
chemicals” which is used to make fertilizers,
polymers, detergents, paints and pigments.
 Sulfuric acid is also the electrolyte in car
batteries.
 Heterogeneous catalysts are preferred in industry
as they are easier to filter off and remove from the
products.

HOMOGENEOUS CATALYSTS
A homogeneous catalyst is in the same phase
(state) as the reactants.
 Example
Iron (II) catalyzes the slow reaction between
acidified hydrogen peroxide and iodide ions.
H2O2(aq) + 2H+(aq) +2Fe2+(aq) →2H2O(l) +
2Fe3+(aq)

2I-(aq) + 2Fe3+(aq)  I2(s) + 2Fe2+(aq)
TRANSITION METALS IN THE BODY
Fe2+ is found in heme in hemoglobin. O2 is
transported around the blood because the Fe2+
can form a weak bond with O2.
 This bond is easily broken when the oxygen
needs to be released.
 Co3+ forms an octahedral complex in vitamin B12.
One of the ligand sites is available for biological
activity.
 Vitamin B12 is needed for the production of red
blood cells and a healthy nervous system.
 Homogeneous catalysts work well in the body as
they mix with the environment they are in.
