IB CHEMISTRY HL 1 UNIT 3 PERIODICITY

1
IB CHEMISTRY HL 1UNIT 3 PERIODICITY

• 11th
• IB t grade opics 3 and 13

2
3.1 The Periodic Table

• 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.

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The Periodic Table of Elements
4
The Periodic Table

• 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.

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The Periodic Table

• We can use the electron configuration to split up
  the valence electrons into sub-levels.
• Example C is He2s22p2.
• Note that valence electrons are in the same main
  energy level.

6
3.2 Physical Properties

• 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.

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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
  atoms outer electrons increases with the group
  number of the element.
• It increases across a period but remains
  approximately the same down a group.

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Effective nuclear charge (Zeff) is the positive
charge felt by an electron.
Zeff Z - s
0 lt s lt Z (s shielding constant)
Zeff ? Z number of inner or core electrons
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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

10
Atomic Radii
covalent radius
metallic radius
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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.

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Trends in Atomic and Ionic Radii
13
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

• 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.

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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.

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Cation is always smaller than atom from which it
is formed. Anion is always larger than atom from
which it is formed.
17
Isoelectronic have the same number of electrons,
and hence the same ground-state electron
configuration
Na Ne
Al3 Ne
F- 1s22s22p6 or Ne
O2- 1s22s22p6 or Ne
N3- 1s22s22p6 or Ne
Na, Al3, F-, O2-, and N3- are all isoelectronic
with Ne
18
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.

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General Trends in First Ionization Energies
20
Trends in First Ionization Energies
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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.

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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.

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Electronegativity

• Maximum value is 4.0 which Fluorine has.
• Minimum value is 0.7 which Francium has.

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The Electronegativities of Common Elements
26
Melting Points
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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

• 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.

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Summary of Trends across Period 3
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3.3 Chemical Properties

• 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-).

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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

• Physical properties
• Good conductors of electricity due to delocalized
  valence electrons
• Low densities
• Soft
• Grey shiny surfaces when freshly cut

• Chemical properties
• Very reactive due to single valence electron that
  is lost easily
• Always form 1 ions and combine easily with
  non-metals such as oxygen and halogens.
• Ex. 2Na(s) Cl2(g) ? 2NaCl(s)

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The Alkali Metals, group 1

• 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.

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Alkali Metals Stored Under Oil
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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.

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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

• 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 non-metals. Reactivity decreases
  down the group.
• They form ionic compounds with metals or
  covalent compounds with non-metals.

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The Halogens, all toxic!
41
The Halogens

• At room temperature, F (pale yellow) and Cl
  (yellow-green) are gases, Br is a red-brown
  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.

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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.

44
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
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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 wouldnt work for F-.

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AgI , AgBr, AgCl, AgF
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SUMMARY OF Ag Hal-
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13.1 Trends across period 3

• 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.

50
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)

51
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)

52
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.
• SiO2(s) 2OH-(aq) ? SiO32-(aq) H2O(l)

53
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)

54
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)

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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
• Iodine vapor
• Color of precipitate when BaCl2 and AgNO3 react
• Color of precipitate from 3 when left in the
  sunlight

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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
• Give the colors of the following
• Iodine vapor purple
• Color of precipitate when BaCl2 and AgNO3
  react white
• Color of precipitate from 3 when left in the
  sunlight black

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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.

58
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.

59
Bonding, Melting and Boiling Points

• Na and Mg form ionic oxides so they are solids at
  room temperature and have high mps and bps.
• SiO2 forms a giant covalent lattice so the mp and
  bp are very high.
• The elements on the right form covalent molecules
  so mps and bps are lower and they exist as
  gases, liquids or low melting solids.

60
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.

61
Summary of Oxides of Period 3
62
13.1 Chlorides of 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. 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.

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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.

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Chlorides of Period 3
Formula of chloride NaCl (s) MgCl2 (s) AlCl3(s) / Al2Cl6(g) SiCl4(l) PCl5(s) / PCl3(l) S2Cl2(l) Cl2(g)
Oxidation number 1 2 3 4 5/3 1 0
Electrical conductivity in molten state High High Poor None None None None
Structure Giant ionic Giant ionic Molecular covalent Molecular covalent Molecular covalent Molecular covalent Molecular covalent
65
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.

66
Aluminum chloride

• Despite being a metal, aluminums compounds often
  behave more like non-metals.
• This is due to the small size and high charge of
  its ion.
• AlCl3 sublimes at 178C to form Al2Cl6 molecules.
• AlCl3 dissociates into ions when added to water
• AlCl3(s) ? Al3(aq) 3Cl-(aq)

67
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)63

68
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)63(aq) ?? Al(H2O)5OH2(aq) H(aq)
• Further proton loss can occur
• Al(H2O)5OH2(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)

69
Silicon Chloride

• Unlike the oxides the Si doesnt form giant
  covalent structures as Cl usually only forms one
  bond.
• The chlorides of non-metals have low mps 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)

70
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)

71
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 (well see
  this again in the unit on redox)

72
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.

73
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.

74
Summary of Oxides and Chlorides
75
Practice Questions

• Which of the following doesnt follow the
  periodicity trend across period 3?
• A. Al2O3 B. Na2O
• C. SO2 D. P4O10
• 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

76
Practice Questions

• Which of the following doesnt follow the
  periodicity trend across period 3?
• A. Al2O3 B. Na2O
• C. SO2 D. P4O10 C
• 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

77
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).

78
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 d-block
  that have a partially filled d-sublevel in one of
  its common oxidation states.
• d-block elements are dense, hard metallic
  elements.

79
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.

80
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

81
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.

82
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.

83
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)

84
Scandium and Zinc

• Sc and Zn dont share all the properties of
  transition elements as they dont 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.

85
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.

86
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.

87
First Row d-block elements
Element Sc Ti V Cr Mn Fe Co Ni Cu Zn
Electronic Structure Ar3d1 4s2 Ar3d24s2 Ar3d3 4s2 Ar3d5 4s1 Ar3d5 4s2 Ar3d6 4s2 Ar3d7 4s2 Ar 3d8 4s2 Ar3d10 4s1 Ar 3d104s2  
Decreasing stability of maximum oxidation state -------gt Decreasing stability of maximum oxidation state -------gt Decreasing stability of maximum oxidation state -------gt Decreasing stability of maximum oxidation state -------gt Decreasing stability of maximum oxidation state -------gt Decreasing stability of maximum oxidation state -------gt Decreasing stability of maximum oxidation state -------gt Decreasing stability of maximum oxidation state -------gt Decreasing stability of maximum oxidation state -------gt Decreasing stability of maximum oxidation state -------gt Decreasing stability of maximum oxidation state -------gt
Increasing stability of 2 oxidation state ---------gt Increasing stability of 2 oxidation state ---------gt Increasing stability of 2 oxidation state ---------gt Increasing stability of 2 oxidation state ---------gt Increasing stability of 2 oxidation state ---------gt Increasing stability of 2 oxidation state ---------gt Increasing stability of 2 oxidation state ---------gt Increasing stability of 2 oxidation state ---------gt Increasing stability of 2 oxidation state ---------gt Increasing stability of 2 oxidation state ---------gt Increasing stability of 2 oxidation state ---------gt
88
Common oxidation states
Sc Ti V Cr Mn Fe Co Ni Cu Zn
7 MnO4-
6 CrO42- Cr2O72- MnO42-
5 VO2 VO3-
4 Ti4 VO2 MnO2
3 Sc3 Ti3 V3 Cr3 Fe3
2 Ti2 V2 Cr2 Mn2 Fe2 Co2 Ni2 Cu2 Zn2
1 Cu
Increasing stability of 2 state -----?------?--------?----------?--------? Increasing stability of 2 state -----?------?--------?----------?--------? Increasing stability of 2 state -----?------?--------?----------?--------? Increasing stability of 2 state -----?------?--------?----------?--------? Increasing stability of 2 state -----?------?--------?----------?--------? Increasing stability of 2 state -----?------?--------?----------?--------? Increasing stability of 2 state -----?------?--------?----------?--------? Increasing stability of 2 state -----?------?--------?----------?--------? Increasing stability of 2 state -----?------?--------?----------?--------? Increasing stability of 2 state -----?------?--------?----------?--------? Increasing stability of 2 state -----?------?--------?----------?--------?
?-----------?-----------?--------?-----Increasing stability of maximum state ?-----------?-----------?--------?-----Increasing stability of maximum state ?-----------?-----------?--------?-----Increasing stability of maximum state ?-----------?-----------?--------?-----Increasing stability of maximum state ?-----------?-----------?--------?-----Increasing stability of maximum state ?-----------?-----------?--------?-----Increasing stability of maximum state ?-----------?-----------?--------?-----Increasing stability of maximum state ?-----------?-----------?--------?-----Increasing stability of maximum state ?-----------?-----------?--------?-----Increasing stability of maximum state ?-----------?-----------?--------?-----Increasing stability of maximum state ?-----------?-----------?--------?-----Increasing stability of maximum state
89
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.

90
Unusual Ion Configurations

• Fe2 Ar3d54s1
• Co2 Ar3d54s2
• Having a half-filled d-block gives stability so
  sometimes the 4s electrons are not all lost first.

91
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)63, Fe(CN)63-, CuCl42- 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.

92
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.

93
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)42 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.

94
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.

95
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.

96
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 complex ion Oxidation state on metal ion Similar to
Cu(NH3)4Cl2 Cu(NH3)42 2 2 CaCl2
K2(CuCl2) (CuCl4)2- 2- 2 K2SO4
97
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.

98
Examples of Complex Ions
Metal ion Water Octahedral Ammonia Octahedral / Square Planar Chloride ion Tetrahedral
Cobalt(II) Pink Co(H2O)62 Straw Co(NH3)62 Blue CoCl42-
Nickel(II) Green Ni(H2O)62 Blue Ni(NH3)62 Yellow-green NiCl42-
Copper(II) Blue Co(H2O)62 Deep blue Cu(NH3)42 Yellow CuCl42-
99
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

100
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

101
Complex Ions

• Cis means the ligands are on the same side.
• Trans means the ligands are on opposite sides.

102
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.

103
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 blue-green 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.

104
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

105
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.

106
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.

107
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)

108
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.

109
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)

110
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.

111
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)

112
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.