IB CHEMISTRY HL 1UNIT 3 PERIODICITY

11th

IB t grade opics 3 and 13

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.

3

The Periodic Table of Elements

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”.

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.

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.

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.

8

Effective nuclear charge (Zeff) is the “positive charge” felt by an electron.

Na

Mg

Al

Si

11

12

13

14

10

10

10

10

1

2

3

4

186

160

143

132

ZeffCoreZ Radius (pm)

Zeff = Z - 0 < < Z ( = shielding constant)

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:

10

Atomic Radii

metallic radius covalent radius

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.

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.

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.

16

Cation is always smaller than atom from which it is formed.Anion is always larger than atom from which it is formed.

17

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

Isoelectronic: have the same number of electrons, and hence the same ground-state electron configuration

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.

19

General Trends in First Ionization Energies

Increasing First Ionization Energy

Incr

ea

sing

Firs

t Io

niz

atio

n E

ner

gy

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.

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.

25

The Electronegativities of Common Elements

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.

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.

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.

SUMMARY OF TRENDS ACROSS PERIOD 3

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-).

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.

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.

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)

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.

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).

THE HALOGENS, GROUP 7

Physical properties:They are coloredThey 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.

THE HALOGENS, ALL TOXIC!

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.

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.

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.

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.

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- Cl- Br- I-

Aqueous Ag+

No Reaction

White precipitate, turns black in sunlight Ag+ + Cl- -> AgCl

Cream precipitate Ag+ + Br- -> AgBr

Pale yellow precipitate Ag+ + I- -> AgI

Chlorine NR NRSolution turns yellow then brown Cl2 + 2Br- -> BR2 + 2Cl-

Solution goes yellow then black precipitate Cl2 + 2I- -> I2 + 2Cl-

Bromine NR NR NRSolution goes yellow then black precipitate Br2 + 2I- -> I2 + 2Br-

Iodine NR NR NR NR

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.

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.

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

1. The reactivity increases in what order?

A. Na, K, Li B. K, Na, LiC. Li, Na, K D. Li, K, NaGive the colors of the following:2. Iodine vapor3. Color of precipitate when BaCl2 and

AgNO3 react4. Color of precipitate from 3 when left

in the sunlight

DO NOW

1. The reactivity increases in what order?

A. Na, K, Li B. K, Na, LiC. Li, Na, K D. Li, K, Na C

Give the colors of the following:2. Iodine vapor purple3. Color of precipitate when BaCl2 and

AgNO3 react white4. 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 Na2O MgO Al2O3 SiO2 P2O5 SO3 Cl2O7 Ratio of Atoms

2:1 2:2 2:3 2:4 2:5 2:6 2:7

Bonding Ionic Ionic Highly polar covalent

Polar covalent

Polar covalent

Polar Covalent

Polar covalent

Acid-Base Character

Basic Basic Amphoteric Weakly acidic

Acidic Acidic Acidic

Other oxides

Na2O2 P2O3 SO2 ClO2 & Cl2O

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.

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) / 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 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 non-metals.

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

Element Na Mg Al Si P S Cl Bonding Metaliic Giant

covalent Molecular Covalent

Chloride Formula NaCl MgCl2 AlCl3 SiCl4 PCl5,

PCl3 Complex Cl2

Bonding Ionic Intermediate Molecular covalent

Oxide Formula Na2O MgO Al2O3 SiO2 P4O10,,

P4O6 SO3, SO2

Cl2O7, Cl2O

Bonding Ionic Intermediate Giant covalent

Molecular covalent

Acid/base properties

Sol. Basic

Insol. Basic

Amphoteric Insol. Acidic

Soluble acidic

PRACTICE QUESTIONS

1. Which of the following doesn’t follow the periodicity trend across period 3?

A. Al2O3 B. Na2OC. SO2 D. P4O10

2. Which of the following would cause a reaction (could be more than one)?

A. Chlorine and sodium bromideB. Bromine and potassium fluorideC. Bromine and calcium iodideD. Iodine and magnesium bromide

PRACTICE QUESTIONS

1. Which of the following doesn’t follow the periodicity trend across period 3?

A. Al2O3 B. Na2OC. SO2 D. P4O10 C

2. Which of the following would cause a reaction (it might be more than one)?

A. Chlorine and sodium bromideB. Bromine and potassium fluorideC. Bromine and calcium iodideD. Iodine and magnesium bromide A &

C

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 d-block that have a partially filled d-sublevel 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

Electronic Structure

[Ar]3d1 4s2

[Ar]3d24s

2

[Ar]3d3 4s2

[Ar]3d5 4s1

[Ar]3d5 4s2

[Ar]3d6 4s2

[Ar]3d7 4s2

[Ar] 3d8 4s2

[Ar]3d10 4s1

[Ar] 3d104

s2 

Decreasing stability of maximum oxidation state ------->

Increasing stability of +2 oxidation state --------->

COMMON OXIDATION STATESSc 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 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 complex ion

Oxidation state on metal 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 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.

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.