Topic 13.2: First-row d-block elements

13.2.1 List the characteristics properties of transition elements.

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.

13.2.4 Define the term ligand

13.2.5 Describe and explain the formation of complexes of d-block elements

13.2.6 Explain why some complexes of d-block elements are coloured

13.2.7 State examples of the catalytic action of transition elements and their compounds.

13.2.8 Outline the economic significance of catalysts in the Contact and Haber processes

Topic 13.1: Trends across period 3

13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure

13.1.2 Describe the reactions of chlorine and the chlorides referred to in 13.1.1 with water

Topic 13: Periodicity

Topic 13 of the IB HL Chemistry syllabus is the Periodicity. IBO recommends to spend 4 hours on this topic.

This topic has 2 sub-chapters: "Trends across period 3", and "First-row d-block elements". Each are separated with numerical values in order of mentioned.

These are advanced HL syllabus statements, it is recommended to bring a Casio Graphical Calculator instead of Texas. Casio Calculators have the periodic table installed already.

Topic 3.3: Chemical properties

3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group

Group 0: Noble Gases

• Colourless
• Monatomic
• Unreactive

Their lack of reactivity can be explained by the inability of their atoms to lose or gain electrons.

Group 1: Alkali Metals

• Good conductors of electricity
• Low Density
• Shiny surface 
• Reactive metals
• Forms ionic compounds with non-metals

They form single charged ions M+. Their low ionization energies give an indication of the ease with which the outer electron is lost.

Reaction with Water

The alkali metals react with water to produce hydrogen and the metal hydroxide.

The reaction becomes more vigorous as we descend the group.


Group 7: Halogens

• Diatomic molecules
• Coloured
• Gradual change from gases (Fluorine and Chlorine) to Liquid (Bromine) to Solids (Iodine and Astatine)
• Reactive non-metals
• Reactivity decrease down the group
• Form ionic compounds with metals or covalent compounds with other non-metals

Reaction with Group 1 metals

Halogens react with Group 1 metals to form ionic halides. The halogen atoms gains one electron from the Group 1 elements to form a halide ion. The electrostatic force of attraction between the alkali metal ions and halides bonds the ions together.

Once the transfer is complete, the ions are pulled together by the mutual attraction of their opposite charges.

This would be the equation

Displacement reactions

The relative reactivity of the elements can also be seen by placing them in direct competition for an extra electron.

A chlorine nucleus has a stronger attraction for an electron than a bromine nucleus because of its smaller atomic radius and so takes the electron from the bromide ion. The chlorine has gained an electron and so forms the chloride ion. The bromide ion loses an electron to form bromine.

Reaction with Silver

The halogens form insoluble salts with silver. Adding a solution containing the halide to a solution containing silver ions produces a precipitate which is useful in identifying the halide

Silver Chloride = White

Silver Bromide = Cream

Silver Iodide = Yellow
Silver Nitrate = Colourless

3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3

The acid-base properties of the oxides are closely linked to their bonding. Metallic elements, which form ionic oxides, are basic; non-metal oxides, which are covalent are acidic. Aluminium oxide, which can be considered as an ionic oxide with some covalent character, shows amphoteric properties reacting with both acids and bases.

Sodium and Magnesium Oxide are bases

Na2O + HCl →2NaCl + H2O
MgO + H2SO4 → MgSO4 + H2O



Aluminium oxide is amphoteric (alkali and acid) so it reacts with both bases and acids


Al2O3 + 3H2SO4 → Al2(SO4)3 + 3H2O
Al2O3 + 2NaOH + 3H2O → 2NaAl(OH)4



Phosphorous Oxide, Sulfur Oxide, Dichlorine Oxide reacts with water to form Acid Solutions

P4O10 + 6H2O → 4H3PO4
P4O6 +6H2O → 4 H3PO3

SO3 + H2O → H2SO4
SO2 + H2O → H2SO3

Cl2O7 + H2O → 2HClO4
Cl2O + H2O → 2HClO


Silcon dioxide doesn't react with water, but reacts with concentrated alkalis to form silicates

SiO2 + 2OH → SiO3 + H2O

Topic 3.2: Physical properties

3.2.1 Define the terms first ionization energy and electronegativity

First Ionization Energy is the minimum energy required to remove an electron from a neutral gaseous atom in its ground state.

Key words:

Minimum Energy, Neutral Gaseous Atom and Ground State

Electro-negativity defined as the relative attraction that an atom has for the shared pair of electrons in a covalent

Key Words:

Relative Attraction, Shared Pair and Covalent


3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electro-negativities and melting points for alkali metals (Li to Cs) and the halogens (F to I)

Atomic Radii

Since the outer electrons are difficult to locate, in practise, the atomic radius is measured as half the distance between two bonded atoms. For this reason noble gases are given no value as they do not bond with other atoms.

On descending a group, the atomic radius increase. This is because the outer electrons are getting further from the nucleus. This applies for both alkali metals and halogens.

Ionic Radii

It is defined by the distance between nucleus and outer most electrons of a positive metal cation or negative non-metal anion.

• Both metal cation and non-metal anion increase in size on descending the group
• Metal cations tend to be smaller than their atom as they have lost their outer most electrons
• Non-metal anions tend to be larger than their atom as they have gained electrons into their outer energy levels

Ionization Energy

It is defined by the minimum energy required to remove an electron from a neutral gaseous atom in its ground state.

• Ionization energy decreases on descending a group
• The outer most electrons are easier to remove as they are a greater distance from the nucleus shielding from the positive nucleus charge by the core electrons

Electronegativity

Electronegativity is defined as the relative attraction that an atom has for the shared pair of electrons in a covalent bond.

• On descending a group, electronegativity decreases. While the atomic radius and the nuclear charge increases, the level of shielding increases and the effective nuclear charge decrease.
• The most electronegative elements are in the top right of the periodic table. (F, O and N). Which are involved in hydrogen bonding.

Melting points

• The variation in melting points is related to the strength of the inter-molecular forces holding the atoms together
• Melting point of metals decreases on descending the group. As the atoms become larger, the metallic bonds holding them together gets weaker. The valence electrons are further from the nucleus.
• Melting point of non-metals increases on descending the group. This is due to the increase in the strength of the Van der Waal's forces between the molecules. The strength of these forces is determined by the total number of electrons in the atom.
• Across a period, the metals in groups 1, 2 and 3 display an increase in melting point. This is due to an increase in strength of the metallic bonds as the number of valence electrons increases
• The Group 4 element will have the highest melting point as it will be a covalently bonded macro-molecule. This contains only strong covalent bonds in the solid phase.
• The non-metal elements in Groups 5, 6 and 7 will have much lower melting point. Theses are also covalently bonded but are simple molecular structures with weaker Van der Waal's forces holding them together.
• The Group 0 noble gases will have the lowest melting point due to it being monoatomic. It will therefore have the weakest of Van der Waal's forces.

3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electro-negativities for elements across period 3

Atomic radii

Across a period, the atomic radius decreases. This is because of the increase in effective nuclear charge and no increase shielding, therefore the electrons get pulled closer towards the nucleus

Ionic radii

Across a period, the metal cation and the non-metal anion get smaller. This is because of an increase in protons (Effective Nuclear Charge) whilst all the ions have the same electronic configuration

First Ionization Energies

Ionization energy generally increases across a period. This is due to the increase effective nuclear charge as the number of protons increases with no increase in shielding.

Electronegativity

Across a period, electronegativity increases. This is the result of an increased number of protons and thus an increased effective nuclear charge with no greater level of shielding


3.2.4 Compare the relative electro-negativity values of two or more elements based on their positions in the periodic table

By looking at the difference in electro-negativity values

We can tell whether a molecule has covalent or ionic bonds

Topic 3.1: The periodic table

3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number

The elements in a periodic table is placed in order of increasing atomic number (Z), which we know is a fundamental property of the element - number of protons in the nucleus of its atom. To read the periodic table, simply start from the top left and read across. You will find that the atomic number increases.

3.1.2 Distinguish between the terms group and period

Group is the columns or the vertical elements. Period is the row or the horizontal elements.

3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z=20.

Groups state how many valence electrons there are

Periods state how many electron shells there are.

Valence electron is an electron in the outer shell of an atom

Example,

Carbon has an electron arrangement of 2, 6

It is in the 2nd period and 6th group, hence its position on the periodic table

Calcium has an electron arrangement of 2, 8, 8, 2

It is in the 4th period and 2nd group, hence its position as an alkali earth metal.


3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table

An atom with one valence electron will be in group 1

2, 1 would be in group 1

An atom with two valence electron will be in group 2

2, 2 would be in group 2

...etc

An atom with a full outer shell will be in group 0

2, 0 would be in group 0

Topic 3: Periodicity

Topic 3 of the IB HL Chemistry syllabus is the Periodicity. IBO recommends to spend 6 hours on this topic.

This topic has 3 sub-chapters: "The periodic table", "Physical properties" and "Chemical Properties". Each are separated with numerical values in order of mentioned.

These are advanced basic syllabus statements, it is recommended to bring a Casio Graphical Calculator instead of Texas. Casio Calculators have the periodic table installed already.