HL IB Chemistry – Periodic Table & Related Topics

Question-and-answer flashcards covering definitions, trends, reactions and explanations from Periodic Table organisation, periodicity, Group 1 & 17 chemistry, oxides, oxidation states, and transition-metal properties including colour, catalysis and variable oxidation states.

What determines the period (row) number of an element?

The highest principal energy level (outermost occupied shell, n) that contains electrons.

What are valence electrons?

The electrons located in the outermost (highest-energy) shell of an atom.

Why are H and He placed atypically on the Periodic Table?

Their unique electron configurations and properties do not fit perfectly into any standard group, so placement is based on closest chemical similarity (He in Group 0; H in its own group).

Which four main blocks make up the Periodic Table?

s-block, p-block, d-block and f-block.

How do you identify an s-block element?

Its outermost electrons are only in an s-subshell.

Across a period, how do atomic radii change and why?

They decrease because nuclear charge increases while electrons are added to the same shell, pulling them closer to the nucleus.

Down a group, how do atomic radii change and why?

They increase due to additional electron shells and greater shielding of outer electrons from the nucleus.

What is the trend in ionic radius with increasing positive charge?

Ionic radius decreases because fewer electrons are held more strongly by the unchanged nuclear charge.

What is the trend in ionic radius with increasing negative charge?

Ionic radius increases because added electrons repel each other while nuclear charge stays the same.

Define first ionisation energy (IE₁).

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1⁺ ions.

How does first ionisation energy vary across a period?

It generally increases due to higher nuclear charge with similar shielding.

How does first ionisation energy vary down a group?

It decreases because outer electrons are farther from the nucleus and more shielded.

Why is there a drop in IE between Be and B?

The 5th electron of B enters the higher-energy 2p subshell, farther from the nucleus than Be’s 2s electrons.

Why is IE lower for O than for N?

Oxygen’s paired 2p electrons experience spin-pair repulsion, making it easier to remove one than from nitrogen’s half-filled p orbitals.

Define electron affinity (EA₁).

The energy released when one mole of gaseous atoms gains one mole of electrons to form one mole of gaseous 1⁻ ions (usually exothermic).

Which group has the most exothermic first electron affinities?

Group 17 (halogens).

State the trend of electronegativity across a period.

Electronegativity increases across a period due to increasing nuclear charge and constant shielding.

State the trend of electronegativity down a group.

Electronegativity decreases because of increased atomic radius and shielding outweighing nuclear charge.

Why are Group 1 metals called alkali metals?

They form alkaline (high-pH) hydroxide solutions when they react with water.

What are the products when sodium reacts with water?

Sodium hydroxide (NaOH) and hydrogen gas (H₂).

How does reactivity change down Group 1 with water?

Reactivity increases due to larger atomic radius and weaker attraction for the outer ns¹ electron.

Write the balanced equation for potassium and water.

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

Why does potassium ignite hydrogen during its reaction with water?

The reaction is so exothermic that the heat produced ignites the evolved H₂, giving a lilac flame.

What is a halogen displacement reaction?

A more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide ion.

Which halogen will displace bromine from KBr solution?

Chlorine, because it is higher (more reactive) in Group 17.

Explain why fluorine is less exothermic in EA than expected.

Its very small atomic radius causes strong e⁻-e⁻ repulsion in the compact 2p subshell, reducing energy released when an extra electron is added.

How do melting points change down the halogens?

They increase because London dispersion forces grow with more electrons and greater molecular mass.

State the oxide trend across Period 3 (Na→Cl) with water.

Basic → amphoteric → acidic; pH falls from strong alkali (Na₂O) to strong acid (SO₃).

Why is Al₂O₃ amphoteric?

It reacts with both acids (as a base) and bases (as an acid) due to intermediate ionic–covalent character.

Give the pH and product(s) when SO₃ reacts with water.

Forms H₂SO₄ giving a strongly acidic solution around pH 1.

Define oxidation state (oxidation number).

The hypothetical charge an atom would have if all bonds were completely ionic.

What is the oxidation number of sulfur in SO₄²⁻?

+6.

State the fixed oxidation numbers for Group 1 and Group 2 elements.

+1 for Group 1, +2 for Group 2 in compounds.

What is Stock notation?

Use of Roman numerals in a compound’s name to indicate the oxidation state of an element with variable valency (e.g., iron(III) oxide).

Write the name for Cu₂O using Stock notation.

Copper(I) oxide.

Which two first-row transition metals have anomalous electron configurations?

Chromium ([Ar]3d⁵4s¹) and copper ([Ar]3d¹⁰4s¹).

When forming cations, from which subshell are electrons removed first in transition metals?

4s electrons are removed before 3d electrons.

Why do transition metals exhibit variable oxidation states?

The similar energies of 4s and 3d electrons allow different numbers to be removed or shared in bonding.

List three characteristic properties of transition elements.

Variable oxidation states, formation of coloured compounds, and catalytic activity (also: magnetism, complex ion formation, high melting points).

Why are Fe, Co and Ni ferromagnetic?

They have unpaired d-electrons whose magnetic dipoles can align, creating permanent magnetism.

Give one industrial heterogeneous catalyst involving a transition metal.

Iron in the Haber process (N₂ + 3 H₂ ⇌ 2 NH₃).

What is the active metal in catalytic converters that converts CO and NOx in car exhausts?

A mixture of platinum and rhodium on a ceramic support.

What metal ion is present in haemoglobin and what is its coordination number?

Fe²⁺ with coordination number 4 (square-planar porphyrin) plus binding to O₂.

Define ligand.

An ion or molecule that donates a lone pair to form a coordinate (dative) bond with a central metal ion.

Explain briefly why [Cr(NH₃)₆]³⁺ is purple while [Cr(OH)₆]³⁻ is dark green.

Different ligands alter d-orbital splitting (ΔE), changing the wavelength of light absorbed and hence the complementary colour observed.

What is crystal field splitting?

Separation of the five degenerate d-orbitals into two energy levels when ligands create an electrostatic field around a transition metal ion.

How does greater d-orbital splitting affect light absorbed?

A larger ΔE means higher-energy (shorter-wavelength) light is absorbed.

If a solution appears green, which colour of light is predominantly absorbed?

Red (its complementary colour).

Why do transition metal complexes form colours but s-block compounds are usually colourless?

s-block ions lack partially filled d-orbitals, so no d-d electron transitions occur to absorb visible light.

State the relationship linking energy of light to frequency.

E = h f, where h is Planck’s constant.

What wavelength range corresponds to yellow light approximately?

About 575 – 585 nm.

What colour is seen when Ti(III) sulfate absorbs yellow light?

Purple (complementary to yellow).

Give the balanced equation for catalytic decomposition of hydrogen peroxide by MnO₂.

2 H₂O₂(aq) → 2 H₂O(l) + O₂(g).

Why do higher oxidation-state transition-metal ions form more covalent bonds?

High charge density polarises ligands’ electron clouds, increasing covalent character.

Which manganese oxidation state appears in the purple manganate(VII) ion?

+7 in MnO₄⁻.

Across the first transition series, which oxidation states are common to all elements?

+2 and +3.

Why does the +3 state dominate early (Ti–Cr) but +2 later (Fe–Cu)?

Ionisation energies and lattice/solvation energies favour +3 for early metals and +2 for later ones due to increasing nuclear charge and d-electron stability.

Explain the term ‘amphoteric oxide’.

An oxide that can react as both an acid and a base, e.g., Al₂O₃.

Which oxide gives the highest pH when added to water among Na₂O, CaO, CO₂, SO₂?

Na₂O because it forms the strongest alkaline solution (2 NaOH).

State four factors that influence ionisation energy.

Nuclear charge, distance of outer electrons (atomic radius), shielding by inner electrons, and spin-pair repulsion.

What is meant by ‘periodicity’?

Regular repeating trends in properties of elements as atomic number increases across periods.

What causes the sharp increase in atomic radius from a noble gas to the next period’s alkali metal?

Addition of a new electron shell significantly increases atomic size despite low nuclear charge increase.