The Periodic Table: Classification Of Elements

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The Periodic Table: Structure and Organisation

The Periodic Table lists all known elements in order of increasing atomic number (1 to 118).
Elements are arranged so that atoms with the same number of shells are placed together, and atoms with similar outer-shell electronic configurations are grouped together.
Periods are rows; each period corresponds to the number of electron shells occupied (e.g. Period 2 elements Li to Ne have 2 shells).
Groups are columns; each group contains elements with the same number of valence electrons, giving similar chemical properties.
Group numbers can be shown as 1–18 (IUPAC format).
Hydrogen is placed in its own group because its properties do not match any one group closely.
Helium is placed in Group 18 because it shares the unreactive nature of the noble gases, despite having only 2 electrons.

The Periodic Table

Metals, Non-Metals and Metalloids

Metals are found on the left and centre of the table (s- and d-blocks, and part of the p-block); they are good conductors of heat and electricity, and are malleable/ductile.
Non-metals are found on the right-hand side (p-block); they are poor conductors, brittle when solid, and often form acidic oxides.
Metalloids are found along a diagonal boundary between metals and non-metals and exhibit properties of both.
Metalloids include elements like boron, silicon, arsenic, antimony.

Metals and non-metals in the Periodic Table

Blocks of the Periodic Table

s-block: Groups 1 and 2 plus helium; outer electrons are in s orbitals.
p-block: Groups 13 to 18; outer electrons are in p orbitals.
d-block: Transition metals (Groups 3 to 12); outer electrons are in d orbitals.
f-block: Lanthanides and actinides; outer electrons are in f orbitals.
These blocks are useful for predicting reactivity, oxidation states, and typical bonding behaviour.
The Periodic Table in the IB Chemistry Data Booklet does not label groups, blocks, or element types, so you must be able to locate periods, groups, and block positions yourself.

Electron Configurations and the Periodic Table

The period number indicates the main energy level (shell) the valence electrons occupy and how many occupied energy levels the atom has.
The group number indicates the number of valence electrons (e.g. Group 4 means 4 valence electrons).
The block (s, p, d, f) is determined by the type of orbital the last electrons occupy.
In the p-block, the element’s position from left to right tells you how many electrons are in the p subshell (e.g. the 2nd element in the p-block has 2 electrons in the p-subshell).
The 4s orbital is lower in energy than the 3d subshell, so it is filled first (e.g. potassium: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹, shorthand [Ar] 4s¹).
When writing full electron configurations, they are often written in numerical order, so 3d is written before 4s (e.g. gallium: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p¹, shorthand [Ar] 3d¹⁰ 4s² 4p¹).
For ions, electrons are lost from the outermost shell first (e.g. Mg²⁺: 1s² 2s² 2p⁶, shorthand [Ne]).

Electron shells and the Periodic Table

Periodic Trends: Atomic and Ionic Radius

Atomic radius is the distance from the nucleus to the outermost electron shell; it is often taken as half the mean distance between two adjacent atoms.
Across a period, atomic radius generally decreases: nuclear charge increases, electrons are added to the same shell, shielding is roughly constant, so outer electrons are pulled closer.
Down a group, atomic radius generally increases: more principal quantum shells are added, outer electrons are further from the nucleus, and shielding increases.
There is a sharp increase in atomic radius between a noble gas and the next alkali metal because the alkali metal has one extra principal quantum shell.
Cations are smaller than their parent atoms because electrons are lost (often an outer shell), reducing electron–electron repulsion and allowing the nucleus to pull remaining electrons closer.
Anions are larger than their parent atoms because extra electrons increase electron–electron repulsion, spreading the electron cloud out further.
Ionic radius across a period: for cations it decreases; after anions form it increases sharply due to added electron repulsion, then decreases again as nuclear charge increases.
Ionic radius down a group increases because more electron shells are added, increasing shielding and distance from the nucleus.

Periodic Trends: Ionisation Energy, Electron Affinity and Electronegativity

First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous ions (e.g. Ca(g) → Ca⁺(g) + e⁻, ΔH = +590 kJ mol⁻¹).
Across a period, ionisation energy generally increases: nuclear charge increases, shielding is roughly constant, so outer electrons are more strongly attracted.
Down a group, ionisation energy generally decreases: outer electrons are further from the nucleus and experience increased shielding.
There is a rapid decrease in ionisation energy between the last element of one period and the first of the next due to increased distance and shielding.
Electron affinity is the energy released when one mole of electrons is gained by one mole of gaseous atoms to form one mole of gaseous ions (e.g. Cl(g) + e⁻ → Cl⁻(g), ΔH = –349 kJ mol⁻¹).
First electron affinity is usually exothermic for non-metals, but some elements (e.g. Group 2, Group 12, Group 15) have positive (endothermic) first electron affinities due to extra repulsion when adding an electron to a stable or half-filled subshell.
Electron affinity across a period becomes more negative/exothermic; down a group it becomes less negative/exothermic.
Electronegativity is the ability of an atom to attract a pair of electrons towards itself in a covalent bond; it increases across a period and decreases down a group.

Group 1 Metals with Water

Group 1 metals are called alkali metals because they form alkaline solutions with high pH when reacted with water.
They are soft and easy to cut (getting softer and denser down the group), have shiny silvery surfaces when freshly cut, conduct heat and electricity, and have low melting points and low densities.
Melting point decreases down the group as atomic radius increases and metallic bonding gets weaker.
Group 1 metals react vigorously with water to form an alkaline metal hydroxide solution and hydrogen gas: group 1 metal + water → group 1 metal hydroxide + hydrogen.
Lithium floats and reacts slowly, releasing bubbles of hydrogen and forming a clear, colourless solution of lithium hydroxide.
Sodium reacts more vigorously, melting into a small silvery ball that moves on the surface, producing hydrogen and a clear, colourless solution of sodium hydroxide.
Potassium reacts more violently, melting into a shiny ball that dashes around, and the heat produced ignites the hydrogen with a lilac flame, forming potassium hydroxide.
Reactivity increases down Group 1 because the outer electron is further from the nucleus and more shielded, so it is more easily lost.

Group 17 Elements with Halide Ions

The halogens are the Group 17 non-metals: fluorine, chlorine, bromine, iodine and astatine; they are poisonous and exist as diatomic molecules (X₂).
All halogens have seven electrons in their outer shell (ns² np⁵) and form halide ions by gaining one electron to complete their outer shell.
Colours and states at room temperature: fluorine is a yellow gas, chlorine a pale green gas, bromine an orange-brown liquid (readily evaporates to a brown gas), iodine a grey-black solid (sublimes to a purple vapour).
Melting and boiling points increase down the group due to increasing molecular mass and stronger van der Waals forces between molecules; the halogens also get darker in colour down the group.
Reactivity decreases down Group 17 because atomic radius and shielding increase, and electron affinity becomes less negative, so the attraction for an incoming electron is weaker.
A more reactive halogen will displace a less reactive halogen from an aqueous solution of its halide; the order of reactivity is chlorine > bromine > iodine.
Example: chlorine displaces bromine from potassium bromide solution (Cl₂ + 2KBr → 2KCl + Br₂), turning the solution orange.
Example: bromine displaces iodine from sodium iodide solution (Br₂ + 2NaI → 2NaBr + I₂), turning the solution brown.

Metallic and Non-Metallic Oxides

Across a period, oxides change from basic through amphoteric to acidic.
Amphoteric means a substance can act as both an acid and a base; aluminium oxide is amphoteric and reacts with both HCl and NaOH.
Oxides on the left of Period 3 (Na₂O, MgO) are basic; Al₂O₃ is amphoteric; oxides on the right (SiO₂, P₄O₁₀, SO₂, SO₃) are acidic.
Basic oxides (Group 1 and 2 metals) form hydroxides in water, giving alkaline solutions: Na₂O + H₂O → 2NaOH (pH ≈ 14); MgO + H₂O → Mg(OH)₂ (pH ≈ 10).
Acidic oxides (non-metals) form oxoacids in water, giving acidic solutions: P₄O₁₀ + 6H₂O → 4H₃PO₄ (pH ≈ 2); SO₂ + H₂O → H₂SO₃ (pH ≈ 1); SO₃ + H₂O → H₂SO₄ (pH ≈ 1).
The trend from basic to acidic oxides is explained by bonding and electronegativity: large electronegativity differences between oxygen and metals on the left give ionic oxides; smaller differences give covalent oxides.
Environmental links: non-metal oxides dissolving in atmospheric water cause acid rain (e.g. SO₂ + H₂O → H₂SO₃; 2NO₂ + H₂O → HNO₃ + HNO₂); CO₂ dissolving in seawater causes ocean acidification (CO₂ + H₂O ⇌ H₂CO₃).

Metals form basic oxides, non-metals form acidic oxides

Oxidation States

Oxidation can be defined as addition of oxygen, loss of hydrogen, or loss of electrons; reduction as loss of oxygen, gain of hydrogen, or gain of electrons.
The oxidation number (or state) of an atom is the charge it would have if all bonding were completely ionic.
Oxidation numbers are used to tell if oxidation or reduction has taken place, identify what has been oxidised/reduced, and construct half-equations and balance redox equations.
Oxidation numbers are always written with the sign before the number (e.g. +1, –2).
For simple ions, the oxidation number equals the charge on the ion (e.g. Na⁺ = +1, Mg²⁺ = +2, Cl⁻ = –1, O²⁻ = –2).
In a neutral molecule or compound, the sum of oxidation numbers is zero (e.g. in H₂O, H = +1 and O = –2; in CO₂, C = +4 and O = –2).
In a polyatomic ion, the sum of oxidation numbers equals the ion’s charge (e.g. in SO₄²⁻, S = +6 and O = –2).
In compounds/ions, the more electronegative element is assigned the negative oxidation number (e.g. in F₂O, F = –1 and O = +2).

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1.Which of the following statements describes first ionisation energy?
Easy
- AThe energy required to remove one mole of electrons from one mole of gaseous atoms
- BThe energy required to remove the outermost electron from each atom in one mole of gaseous atoms
- CThe energy required to remove the outermost electron from each atom in one mole of atoms
- DThe energy required to produce one mole of ions from one mole of gaseous atoms
2.Sodium sulfide, Na2S, is a reactive yellow solid, produced when sodium and sulfur react together. How do the ionic radius and atomic radius of sodium compare with those of sulfur?
Easy
- Aionic radius: sodium < sulfur; atomic radius: sodium < sulfur
- Bionic radius: sodium < sulfur; atomic radius: sodium > sulfur
- Cionic radius: sodium > sulfur; atomic radius: sodium > sulfur
- Dionic radius: sodium > sulfur; atomic radius: sodium < sulfur
3.Which property below decreases generally across the second period?
Easy
- AAtomic number
- BAtomic radius
- CElectronegativity
- DIonisation energy
4.The trends in three physical properties of the elements of Period 3 are shown in graphs. Which physical property is not illustrated?
Medium
- AElectrical conductivity
- BAtomic radius
- CMelting point
- DFirst ionisation energy
5.This question refers to isolated gaseous species. The species O2−, Ne and Mg2+ are isoelectronic. In which order do their radii increase?
Medium
- AMg2+ < Ne < O2−
- BNe < O2− < Mg2+
- CMg2+ < O2− < Ne
- DO2− < Ne < Mg2+
6.Which graph correctly shows the electronegativity of the elements Na, Mg, Al and Si, from Period 3, plotted against their first ionisation energies?
Medium
- AA graph showing a positive correlation: as ionisation energy increases, electronegativity increases.
- BA graph showing a negative correlation: as ionisation energy increases, electronegativity decreases.
- CA graph showing no correlation.
- DA graph showing a random scatter.
7.Which period 3 element from sodium to silicon has the largest atomic radius?
Easy
- AMagnesium
- BSodium
- CSilicon
- DAluminium
8.What is the correct classification for the element zirconium, Zr?
Easy
- As block
- Bd block
- Cf block
- Dp block

더 많은 문제

9.What happens to the pH of water when magnesium oxide is dissolved into it?
Easy
- AThe pH increases, because MgO is acidic
- BThe pH decreases, because MgO is basic
- CThe pH increases, because MgO is basic
- DNothing happens because MgO is neutral
10.The Group II metals have higher melting points than the Group I metals. Which factors result in Group II metals having higher melting points? (select all that apply)
Medium
- AGroup II metals have higher first ionisation energy than the corresponding group I metal
- BThere are smaller interatomic distances in the metallic lattices of the Group II metals
- CMore electrons are available from each Group II metal atom for metallic bonding in the lattice
- DGroup II metals have larger atomic radii than Group I metals
11.Hydrogen is placed in Group 1 of the periodic table because its properties closely match those of the alkali metals.
Easy
True or false?
12.Match each block of the periodic table with the type of orbital that contains the outer electrons.
Medium
- s-block
- p-block
- d-block
- f-block
- s orbitals
- p orbitals
- d orbitals
- f orbitals
13.Place the following elements in order of increasing atomic radius: Mg, Be, C, N
Medium
- Mg
- Be
- C
- N
14.When a student dissolved the oxides of calcium and tellurium in water, the resulting solutions were tested with litmus paper. It would be seen that:
Medium
- ACalcium turns litmus paper blue, but tellurium turns litmus paper red
- BCalcium and tellurium both turn litmus paper red
- CCalcium turns litmus paper red, but tellurium turns litmus paper blue
- DCalcium and tellurium both turn litmus paper blue
15.Which of the following are properties of transition metals? (select all that apply)
Medium
- AThey have variable oxidation states
- BThey can behave as catalysts
- CThey can have magnetic properties
- DThey are all found in the s-block
16.The reactivity of Group 1 metals decreases down the group from lithium to caesium.
Easy
True or false?
17.Which pair of elements reacts most readily?
Medium
- ALi + Br2
- BLi + Cl2
- CK + Br2
- DK + Cl2
18.Which of the following are characteristic properties of transition metals? (select all that apply)
Medium
- AThey have variable oxidation states
- BThey can act as catalysts
- CThey can have magnetic properties
- DThey have low melting points
- EThey are all gases at room temperature
19.Which factors affect the colour of a transition metal ion complex? (select all that apply)
Medium
- AThe identity of the metal ion
- BThe oxidation state of the metal ion
- CThe number of neutrons in the metal ion
- DThe type of ligand attached to the metal ion
- EThe temperature of the solution
20.Which of the following statements about strontium oxide and lithium oxide are true? I. Strontium oxide has pH > 7 II. Lithium oxide is basic III. A solution containing both oxides would turn universal indicator blue
Medium
- AI and II only
- BI and III only
- CII and III only
- DI, II and III
21.Which of the following statements about the trend in metallic radius in Group 1 elements from Na to Rb is correct? I. It increases moving down the group II. It increases moving down due to the addition of electron shells III. It decreases moving down due to increasing nuclear forces
Medium
- AI and II only
- BI and III only
- CII and III only
- DI, II and III
22.An element has the electron configuration 1s² 2s² 2p⁶ 3s² 3p⁴. In which period, group and block is this element found?
Medium
- APeriod 3, Group 16, p-block
- BPeriod 3, Group 14, p-block
- CPeriod 2, Group 16, s-block
- DPeriod 4, Group 16, d-block
23.An element X is in Period 4 and Group 6. Which statement about X is correct?
Medium
- AX has 4 occupied energy levels.
- BX has 6 valence electrons.
- CX is a non-metal.
- DX forms ions with a 2– charge.
24.Which of the following is the abbreviated electron configuration of the nickel(II) ion, Ni²⁺?
Medium
- A[Ar] 3d⁸
- B[Ar] 4s² 3d⁶
- C[Ar] 4s² 3d⁸
- D[Ar] 4s¹ 3d⁷
25.Which graph correctly illustrates a trend found in the halogen group?
Medium
- ABoiling point increases down the group.
- BBoiling point decreases down the group.
- CFirst ionisation energy increases down the group.
- DElectronegativity increases down the group.
26.Which of the following statements about strontium oxide and lithium oxide are true? I. Strontium oxide has pH > 7. II. Lithium oxide is basic. III. A solution containing both oxides would turn universal indicator green.
Medium
- AI and II only
- BI and III only
- CII and III only
- DI, II and III
27.The Period 4 elements gallium (Ga), germanium (Ge), arsenic (As) and selenium (Se) are directly below aluminium, silicon, phosphorus and sulfur. Which Period 4 elements form oxides that dissolve in water to give an acidic solution?
Medium
- AGa and Ge
- BGe and Se
- CAs and Se
- DSe only
28.Non-metallic elements in the p block can have the following two properties: property 1: atoms with no paired electrons in 3p; property 2: have an oxide that can form a strong acid in water. Which properties do phosphorus and sulfur have?
Medium
- Asulfur: 2 only; phosphorus: 1 and 2
- Bsulfur: 1 and 2; phosphorus: 2 only
- Csulfur: 1 and 2; phosphorus: 1 and 2
- Dsulfur: 2 only; phosphorus: 1 only
29.Which of the following statements is true as you move down Group 1?
Medium
- AThe first ionisation energy increases
- BThe atomic radius increases
- CThe melting point increases
- DThe reaction with water becomes less vigorous
30.Which statement is incorrect?
Medium
- ACaesium has a lower first ionisation energy than rubidium
- BRubidium reacts less violently with water than potassium
- CPotassium has a larger atomic radius than sodium
- DSodium melts at a higher temperature than caesium
31.Which of the following statements about the trend in metallic radius in Group 1 elements from Na to Rb is/are correct? (select all that apply)
Medium
- AIt increases moving down the group
- BIt increases moving down due to the addition of electron shells
- CIt decreases moving down due to increasing nuclear forces
- DIt stays the same because the number of valence electrons is constant
32.Which of the following statements about the halogens are true? (select all that apply)
Medium
- AThe boiling points increase down Group 17.
- BThe reactivity of the halogens decreases down Group 17.
- CThe halogens become lighter in colour down the group.
- DChlorine can displace bromine from an aqueous solution of potassium bromide.
- EIodine can displace chlorine from an aqueous solution of sodium chloride.
33.Use the following data for Period 3 elements: Na (electronegativity 0.9), Mg (1.2), Al (1.5), Si (1.8), P (2.1), S (2.5). Oxygen has an electronegativity of 3.4. Which oxide is expected to have the most ionic character?
Hard
- ANa₂O
- BMgO
- CAl₂O₃
- DSO₃
34.The first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of gaseous atoms. Which equation represents the second ionisation energy of magnesium?
Medium
- AMg(g) → Mg⁺(g) + e⁻
- BMg⁺(g) → Mg²⁺(g) + e⁻
- CMg(g) → Mg²⁺(g) + 2e⁻
- DMg²⁺(g) → Mg³⁺(g) + e⁻
35.Which of the following statements about first electron affinity is correct?
Medium
- AIt is always exothermic for all elements.
- BIt is the energy required to remove an electron from a gaseous atom.
- CIt is the energy released when one mole of gaseous atoms gains one mole of electrons.
- DIt becomes more negative down a group.
36.Which of the following oxides would react with both hydrochloric acid and sodium hydroxide solution?
Medium
- ANa₂O
- BMgO
- CAl₂O₃
- DSO₂
37.Which of the following pairs of elements would react most vigorously together?
Medium
- ALi and Br₂
- BLi and Cl₂
- CK and Br₂
- DK and Cl₂
38.The first ionisation energy of an element generally increases across a period and decreases down a group.
Easy
True or false?
39.The second electron affinity of oxygen is exothermic because the second electron is attracted to the O⁻ ion.
Medium
True or false?
40.What is the oxidation number of sulfur in the sulfate ion, SO4²−?
Easy
41.Complete the definition of electronegativity.
Easy
Electronegativity is the ability of an atom to attract a pair of electrons towards itself in a ____ bond.
42.What is the name of the group of elements that are found along the diagonal boundary between metals and non-metals and exhibit properties of both?
Easy
43.Calculate the oxidation number of sulfur in the sulfate ion, SO₄²⁻.
Medium
44.Calculate the oxidation number of oxygen in hydrogen peroxide, H₂O₂.
Medium
45.The first ionisation energy of sodium is 496 kJ mol⁻¹. Calculate the energy required to remove one electron from a single sodium atom, in joules. Use Avogadro's constant, L = 6.02 × 10²³ mol⁻¹. Give your answer in scientific notation to 2 significant figures.
Hard
46.Complete the definition of first ionisation energy.
Medium
The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous ____ to form one mole of gaseous ____.
47.Identify the block of the periodic table for each element based on its electron configuration.
Medium
Element with configuration [Ar] 4s² is in the ____ block. Element with configuration [Ar] 3d⁵ 4s¹ is in the ____ block. Element with configuration [Ar] 4s² 4p³ is in the ____ block.
48.The following table shows the melting points of the halogens. Complete the table by filling in the missing values.
Medium
Fluorine: –220 °C, Chlorine: –101 °C, Bromine: ____ °C, Iodine: ____ °C
49.Explain why the first ionisation energy of magnesium is higher than that of sodium.
Medium
50.Explain why the boiling points of the halogens increase down Group 17 from chlorine to iodine.
Hard
51.Explain why the second electron affinity of oxygen is endothermic while the first electron affinity is exothermic.
Hard
52.What is the oxidation number of sulfur in the sulfate ion, SO4²−?
Easy
- A8
- B6
- C5
- D7
53.Calculate the oxidation number of sulfur in the sulfate ion, SO₄²⁻.
Medium
- A5
- B8
- C7
- D6
54.Calculate the oxidation number of oxygen in hydrogen peroxide, H₂O₂.
Medium
- A-1
- B-2
- C1
- D0
55.The first ionisation energy of sodium is 496 kJ mol⁻¹. Calculate the energy required to remove one electron from a single sodium atom, in joules. Use Avogadro's constant, L = 6.02 × 10²³ mol⁻¹. Give your answer in scientific notation to 2 significant figures.
Hard
- A8.2e-19 J
- B-1 J
- C2 J
- D1 J

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