The Periodic Table: Classification of Elements (DP IB Chemistry)

Hydrogen is placed in a group on its own because it does not share physical and chemical properties with members of the other groups.

The electron configuration provides information about an element's period, group, and block in the Periodic Table.

Group 1 of the Periodic Table is called the alkali metals.

The halogens are Group 17 of the Periodic Table.

Group 18 of the Periodic Table is called the noble gases.

Transition metals are located in d-block of the Periodic Table.

True.

The Period number shows the outer energy levels that is occupied by electrons.

The electronic configuration of the element in Group 2 Period 3 is:

1s²2s²2p⁶3s²

The electronic configuration of the element in Group 17, Period 4 is:

1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁵

This element is a transition element, located in d block of the Periodic Table in Period 4.

False.

The number of valence electrons equals the group number for main group elements, but not for transition elements.

The ionisation energy (IE) of an element is the amount of energy required to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions.

X (g) → X⁺ (g) + e^-

Going across Period 3, atomic radius decreases because:

• The nuclear charge increases

• So, there is a greater attraction between the nucleus and the electrons.

Electron affinity is the amount of energy released when one mole of electrons is gained by one mole of atoms of an element in the gaseous state to form one mole of gaseous ions.

X (g) + e^- → X^- (g)

Going down a group electronegativity decreases because the nuclear charge, atomic radius and shielding by inner electrons increases. So, there is less attraction between the nucleus and outer bonding electrons.

True.

Atomic radii increases going down a group.

The first ionisation energy increases across a period because the nuclear charge increases but there is similar shielding and the distance between the nucleus and outer electron remains reasonably constant so more energy is needed to remove an electron.

The unit of electron affinity is ‎ kJ mol^-1.

False.

Electron affinities generally decrease down a group.

Going across a period, the electronegativity of an element increases.

The ionic radius of Mg²⁺ is smaller than the ionic radius of Na⁺ because the magnesium ion has a higher nuclear charge and fewer electrons so there is a greater force of attraction between them.

The two products formed when sodium reacts with water are sodium hydroxide and hydrogen.

Properties of Group 1 metals are:

• They are soft

• They have relatively low densities

• They have relatively low melting points

• They are very reactive.

The alkali metals share similar chemical properties because they have the same number of electrons in the outer shell.

The word equation for the reaction between lithium and water is:

lithium + water → lithium hydroxide + hydrogen

The balanced symbol equation for the reaction between potassium and water is:

2K (s) + 2H₂O (l) → 2KOH (aq) + H₂ (g)

The product formed when sodium reacts with oxygen is sodium oxide.

Going down Group 1, reactivity increases.

Other observations include:

• Floating

• Moving

• Melting / turning into a ball / sphere

• Getting smaller / disappearing / dissolving

• Leaving a white trail.

Going down Group 1, melting point decreases.

True.

Going down Group 1, density of the elements increases.

Group 1 metals are known as the alkali metals because they form alkaline solutions when they are added to water.

The reactivity of Group 1 metals increases down the group because the outer electron is further from the nucleus and more shielded, which means that there is a weaker force of attraction. This makes it easier to lose the outer electron in reactions.

True.

All the alkali metals react vigorously with the halogens to form an alkali metal halide salt.

Halogens form halide ions by gaining one electron to complete their outer shell.

False.

The reactivity of halogens decreases as you go down Group 17.

A halogen displacement reaction occurs when a more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide.

The melting and boiling points of halogens increase as you go down Group 17.

True.

Chlorine can displace iodine from an aqueous solution of potassium iodide.

The reactivity of halogens decreases down Group 17 because:

• The number of shells of electrons increases so shielding also increases

• This means that the outer electrons are further from the nucleus

• There are weaker electrostatic forces of attraction that attract the extra electron needed

• So, it is harder to gain an electron

The colour and state of chlorine at room temperature is a pale yellow-green gas.

At room temperature, iodine is a grey-black solid.

False.

The halogens are diatomic.

The balanced symbol equation for the reaction between potassium iodide and bromine is:

Br₂ + 2KI → 2KBr + I₂

Iodine cannot displace chlorine from potassium chloride because iodine is less reactive than chlorine.

When chlorine displaces bromine from potassium bromide the solution will turn orange.

True.

When potassium iodide reacts with bromine, a brown solution is formed.

The acid-base character of oxides changes from basic through amphoteric to acidic across a period.

An oxide that can act both as a base and as an acid is known as an amphoteric oxide.

The amphoteric oxide in Period 3 is aluminium oxide.

False.

Metallic oxides generally form hydroxides (bases) when they react with water.

Non-metallic oxides form acids when they react with water.

The balanced symbol equation for the reaction between magnesium oxide and water is:

MgO (s) + H₂O (l) → Mg(OH)₂ (aq)

The pH of the solution formed when NO₂ is added to water is pH 1.

The pH of solutions formed by metal oxides decreases (becomes less basic) across a period.

False.

The oxides of elements become less ionic and more covalent as you go across a period.

The product formed when SO₂ is added to water is sulfurous acid, H₂SO₃ (aq).

False.

Calcium oxide forms a solution with a lower pH than sodium oxide when added to water.

SiO₂ has a giant covalent structure.

The formula of the two products formed when NO₂ is added to water are HNO₃ and HNO₂.

The oxidation state is a number assigned to an atom to show the number of electrons transferred in forming a bond. It is the charge that atom would have if a compound was composed of ions.

False.

Oxidation states can be positive, negative, or zero.

The oxidation number of an atom in its elemental state is 0.

The oxidation number of hydrogen in most compounds is +1.

The oxidation number of oxygen in most compounds is -2.

The oxidation states of Group 2 elements is +2.

The oxidation number of sulfur in SO₄^2- is +6.

The oxidation state of oxygen in hydrogen peroxide, H₂O₂, is -1.

False.

In either a compound or an ion, the more electronegative element is given the negative oxidation number.

The oxidation state of an element in a mono-atomic ions is always the same as the ‎ ‎ ‎ charge.

The oxidation state of hydrogen in, sodium hydride, NaH, is -1.

The oxidation state of Zn is 0 because it is an element.

True.

The oxidation state of Fe²⁺ is +2.

The oxidation state of aluminium in Al₂Cl₆ is +3.

The oxidation state of Mn in KMnO₄ is +7.

There is a rapid decrease in ionisation energy between the last element in one period and the first element in the next period.

The decrease in ionisation energy between beryllium and boron is caused by the fifth electron in boron entering the 2p subshell, which is further from the nucleus than the 2s subshell of beryllium, so less energy is needed to remove it.

The decrease in ionisation energy between nitrogen and oxygen is because:

• Nitrogen has the electronic configuration of 1s² 2s² 2p³

• Oxygen has the electronic configuration of 1s² 2s² 2p⁴

• In oxygen there is spin-pair repulsion in the 2p subshell

Spin-pair repulsion is the repulsion between electrons of opposite spin in the same orbital, which can affect ionisation energies.

Across a period, the ionisation energy generally increases.

True.

The discontinuity in the trend of ionisation energy across a period provides evidence for energy sublevels.

There is a decrease in first ionisation energy between the last element in one period and the first element in the next period because the distance between the nucleus and outer electron increases and there is more shielding by inner electrons. These two factors outweigh the increased nuclear charge.

False.

The first ionisation energy of sodium is lower than the first ionisation energy of neon because sodium is in the next period.

Two characteristic properties of transition elements include:

• Variable oxidation states

• High melting points

• Have magnetic properties

• Behave as catalysts

• Form coloured compounds

• Form complex ions with ligands.

The variable oxidation states in transition elements are caused by the presence of electrons in both 4s and 3d orbitals that can be lost.

False.

Only some transition elements, such as iron, cobalt, and nickel, have strong magnetic properties.

Haemoglobin, which contains iron, is an example of a biological catalyst containing a transition element.

A complex ion consists of a central metal atom or ion surrounded by a number of molecules or ions called ligands.

A ligand is a molecule or ion that bonds to a central metal atom in a complex.

Transition elements lose electrons to form positively charged ions.

True.

Roman numerals are used to indicate the oxidation state of transition metal ions.

Transition metals form complex ions because they have variable oxidation states.

A heterogeneous catalyst is a catalyst that is in a different physical state / phase from the reactants.

False.

Iron in the Haber process is an example of a heterogeneous catalyst because it is in a different physical state to the reactants - ammonia and hydrogen gas.

Some transition elements have magnetic properties because of the unpaired electron in the d-orbital being aligned in an external field.

False.

While 4s is generally filled before 3d, there are exceptions like chromium and copper.

The electron configuration of chromium is [Ar] 3d⁵ 4s¹.

The Aufbau Principle states that electrons occupy the lowest energy orbitals available before filling higher energy orbitals.

The electronic configuration of copper Cu is [Ar] 3d¹⁰ 4s¹ not [Ar] 3d⁹ 4s² because having a full 3d subshell is more energetically stable.

The electronic configuration of Mn(III) ions is 1s²2s²2p⁶3s²3p⁶3d^4.

False.

When transition elements form cations, the 4s electrons are removed first.

Transition elements have variable oxidation states because the ionisation energies of successive transition elements is relatively small so more than one electron can be lost easily.

The electronic configuration of Ni²⁺ is [Ar] 3d⁸.

The electronic configuration of Cu²⁺ is ‎ [Ar]3d⁹.

The most common oxidation states for iron ions are +2 and +3.

The splitting of d-orbitals is caused by the electrostatic repulsion between the electrons in the d-orbitals and the electrons in the ligands.

Wavelength and frequency of light are inversely proportional, related by the equation c = fλ, where c is the speed of light.

False.

A larger splitting of d-orbitals results in absorption of shorter wavelength light.

The magnitude of d-orbital splitting is determined by:

• The nature of the metal ion

• The oxidation state of the metal

• The type of ligands

True.

All d-orbitals have the same energy in an isolated transition metal ion.

Transition element complexes are coloured due to the absorption of light when an electron is promoted between the orbitals in the split d-sublevels.

A complementary colour is the colour that, when combined with a given colour, produces white light. It is directly opposite on the colour wheel.

The units for wavelength are metres, m.

The units for frequency are s^-1.

Non-degenerate orbitals are ones which are not equal in energy in the d-subshell.