Reading the table as a map

The periodic table is more than a list of elements. Read it like a map: an element’s position tells you something about its chemistry. A period is a horizontal row of elements in the periodic table. A group is a vertical column of elements in the periodic table. Groups are numbered $1$ to $18$, and in assessments you should use the periodic table in the data booklet as your reference.

A block is a region of the periodic table in which the highest-energy electrons are being placed into the same type of sublevel. The four blocks are the s-block, p-block, d-block and f-block, matching the s, p, d and f sublevels. The s-block sits on the left, the p-block on the right, and the d-block in the centre. The f-block is usually shown separately below the main table to save space.

Metals, non-metals and metalloids

A metal is an element that typically forms giant metallic structures with delocalized electrons and is found on the left and centre of the periodic table. A non-metal is an element that typically forms covalent molecular or network structures, or monatomic gases, and is found on the upper right of the periodic table. Hydrogen is the awkward exception: it is placed above group $1$ in many tables because it has one electron, but it is not an alkali metal.

A metalloid is an element with properties intermediate between metals and non-metals, often able to show both metallic and covalent behaviour. Metalloids sit along the stepped boundary between the metals and non-metals, including elements such as boron, silicon and germanium. Don’t treat the boundary as a brick wall; metallic and non-metallic character form a continuum.

Period number and outer energy level

The principal quantum number,

, is

For main-group elements, the period number shows the outer occupied energy level. So, for period 3 elements, the valence electrons are in the third energy level, occupying 3s and/or 3p sublevels.

Valence electrons are electrons in the outer occupied energy level of an atom that are available for bonding. Elements in the same group share the same pattern of valence electrons. That’s why a group is often treated as a chemical family, not just a column in the table.

You should know these named classifications:

For groups 1 and 2, the group number gives the number of valence electrons. For groups 13–18, use the final digit instead. Group 16 elements therefore have six valence electrons; group 18 elements have eight, except helium, which has two.

Deducing electron configurations up to argon and krypton

The atomic number,

, is

You need to deduce electron configurations up to $Z = 36$ from an element’s position in the periodic table, and also work backwards from a configuration to a position.

Keep the maximum occupancies ready: s holds 2 electrons, p holds 6, d holds 10 and f holds 14. Up to $Z = 36$, the filling order you need is:

$1s\ 2s\ 2p\ 3s\ 3p\ 4s\ 3d\ 4p$

For example, bromine is in period 4 and group 17. Its electrons fill to argon, then $4s^2$, then $3d^{10}$, then $4p^5$, giving:

$Br: [Ar]\ 4s^2\ 3d^{10}\ 4p^5$

The reverse matters just as much. If an atom has configuration $[Ne]\ 3s^2\ 3p^4$, its outer level is 3, so it is in period 3. It has six valence electrons, so it is in group 16: sulfur.

Classification and discovery

The organization of elements has a strong nature-of-science story. Early periodic tables placed known elements so that similar properties came round again. Gaps were not just empty boxes; they acted as predictions. When an undiscovered element was expected to fit a gap, chemists could estimate its likely properties from the elements around it. That is what a good classification can do: it does more than tidy existing knowledge, it points to what to look for next.

Classification is also a model, not a law of nature. Hydrogen’s position and the exact membership of group 3 show that a useful classification can still have disputed edges.

What periodicity means

Periodicity is the repeating pattern in physical and chemical properties of elements when they are arranged by atomic number. The patterns appear because electron configurations repeat in a regular way. When you explain a trend, keep coming back to three things: nuclear charge, number of occupied energy levels, and shielding.

Shielding is the reduction in attraction between the nucleus and an outer electron caused by inner electrons repelling that outer electron. Effective nuclear charge is the net positive attraction experienced by an electron after shielding has been taken into account.

Atomic radius

Atomic radius is a measure of the size of an atom, usually taken as half the distance between the nuclei of two bonded atoms of the same element. Across a period, atomic radius decreases. Each successive atom has one more proton, so the nuclear charge increases. Electrons, however, are added to the same main energy level, so shielding by inner shells changes very little. The nucleus pulls the outer electrons closer.

Down a group, atomic radius increases. Each step down adds a new occupied energy level. The outer electrons sit further from the nucleus and are more shielded, so the increased distance and shielding outweigh the increased nuclear charge.

Ionic radius

Ionic radius is a measure of the size of an ion. A cation is a positively charged ion formed when an atom loses electrons. Cations are smaller than their parent atoms because an electron shell may be lost, and the same nucleus attracts the remaining electrons more strongly.

An anion is a negatively charged ion formed when an atom gains electrons. Anions are larger than their parent atoms because the extra electrons increase electron–electron repulsion and reduce the attraction per electron.

For ions with the same electron configuration, size depends strongly on nuclear charge. An isoelectronic series is a set of atoms or ions with the same number and arrangement of electrons. In an isoelectronic series such as $O^{2-}$, $F^-$, Ne, $Na^+$ and $Mg^{2+}$, radius decreases as proton number increases, because the same electron cloud is pulled in more strongly.

Ionization energy

First ionization energy is the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous $1^+$ ions. It is represented by:

$X(g) \to X^+(g) + e^-$

Across a period, first ionization energy generally increases. Nuclear charge increases, shielding changes little, and atomic radius decreases, so removing an outer electron takes more energy. Down a group, first ionization energy decreases because the outer electron is further from the nucleus and more shielded.

Electron affinity

First electron affinity is the energy change when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous $1^-$ ions. It is represented by:

$X(g) + e^- \to X^-(g)$

Across a period, electron affinity generally becomes more exothermic because increasing nuclear charge attracts an added electron more strongly. Down a group, the pattern is less tidy than the atomic radius and ionization energy trends. Increased shielding would suggest a less exothermic value, but electron–electron repulsions and sublevel arrangements also matter. In exams, if data are provided, use the data rather than forcing a perfect trend that isn’t there.

Electronegativity

Electronegativity is a measure of the ability of an atom in a bond to attract the bonding electrons towards itself. Electronegativity increases across a period because nuclear charge increases and atomic radius decreases. It decreases down a group because the bonding electrons are further from the nucleus and more shielded.

Summary of periodic trends across periods and down groups.

These trends are connected. Small atoms with high effective nuclear charge tend to have high ionization energies and high electronegativities. Large atoms with well-shielded outer electrons tend to have lower ionization energies and more metallic behaviour.

Group 1: increasing metallic character

Metallic character means how strongly an element shows metallic behaviour, especially losing electrons to form positive ions. Group 1 elements, the alkali metals, each have one valence electron. As you go down the group, first ionization energy decreases because the outer electron is farther from the nucleus and has more shielding. The atom loses that electron more easily, so metallic character and reactivity increase.

When a group 1 metal reacts with water, it forms a metal hydroxide and hydrogen gas. Using M as a general symbol for a group 1 metal, the balanced equation is:

$2M(s) + 2H_2O(l) \to 2MOH(aq) + H_2(g)$

For example:

$2K(s) + 2H_2O(l) \to 2KOH(aq) + H_2(g)$

The solution becomes alkaline because the reaction produces hydroxide ions. The reactions get faster and more vigorous down the group: lithium reacts steadily, sodium reacts more vigorously, and potassium reacts more violently. Rubidium and caesium are usually shown by video or simulation rather than in an open classroom.

Group 17: decreasing non-metallic character

Non-metallic character means how strongly an element shows non-metal behaviour, especially gaining or attracting electrons. Group 17 elements, the halogens, each have seven valence electrons. Down the group, electronegativity decreases because atomic radius and shielding increase. Halogens become less effective at gaining an electron, so their reactivity decreases down the group.

A more reactive halogen displaces the halide ion of a less reactive halogen. A halide ion is a $1^-$ ion formed from a halogen atom by gaining one electron. For example, chlorine displaces bromide ions:

$Cl_2(aq) + 2Br^-(aq) \to 2Cl^-(aq) + Br_2(aq)$

Bromine does not displace chloride ions, because bromine is less reactive than chlorine:

$Br_2(aq) + Cl^-(aq) \to \text{no reaction}$

The order to remember is:

$F_2 > Cl_2 > Br_2 > I_2$

So fluorine displaces chloride, bromide and iodide ions. Chlorine displaces bromide and iodide ions. Bromine displaces iodide ions only, and iodine displaces none of the common halides above it.

Why simulations are used

Teachers often use simulations and online reaction videos for group 1 and group 17 reactivity because the real reactions can be hazardous, fast, corrosive or difficult to control fairly. Students can compare trends side by side, slow down observations, repeat conditions, and see reactions that would be unsafe in a school laboratory. The trade-off is simple: a simulation is a model. It helps with patterns, but it cannot replace careful experimental evidence when practical work is safe and appropriate.

Oxides across the periodic table

An oxide is a compound containing oxygen bonded to one or more other elements. Across a period, oxides shift from basic metal oxides, through amphoteric oxides, to acidic non-metal oxides. It’s a clear chemical clue that metallic and non-metallic behaviour sits on a continuum, rather than fitting neatly into two separate boxes.

A basic oxide is an oxide that reacts with acids or water to produce a base. Group 1 metal oxides react with water to form soluble metal hydroxides:

$M_2O(s) + H_2O(l) \to 2MOH(aq)$

For example:

$Na_2O(s) + H_2O(l) \to 2NaOH(aq)$

Group 2 metal oxides react with water to form metal hydroxides:

$MO(s) + H_2O(l) \to M(OH)_2(aq\text{ or }s)$

For example:

$CaO(s) + H_2O(l) \to Ca(OH)_2(aq)$

Some group 2 hydroxides dissolve only slightly, so the state symbols can change depending on the metal and the conditions.

An acidic oxide is an oxide that reacts with water to produce an acid, or reacts with bases to form salts. Carbon and sulfur form non-metal oxides that dissolve in water to form acids:

$CO_2(g) + H_2O(l) \rightleftharpoons H_2CO_3(aq)$

$SO_2(g) + H_2O(l) \to H_2SO_3(aq)$

$SO_3(g) + H_2O(l) \to H_2SO_4(aq)$

An amphoteric oxide is an oxide that can react both as an acid and as a base. Aluminium oxide is the classic example: it reacts with acids and with strong bases. That is why the period 3 pattern is usually written as basic $Na_2O$ and $MgO$, amphoteric $Al_2O_3$, then acidic $SiO_2$, $P_4O_{10}$, $SO_2$ and $SO_3$.

Period 3 oxides change from basic to amphoteric to acidic across the period.

Bonding link: why metal and non-metal oxides behave differently

Metal oxides usually contain ionic lattices with oxide ions, $O^{2-}$. The oxide ion is strongly basic because it can accept protons from water or acids, forming hydroxide ions or water. Non-metal oxides are usually covalent molecules. When they react with water, they form molecules that release $H^+$ in aqueous solution, so their solutions are acidic. The bonding matters here: it explains the difference in their chemical properties.

Acid rain and ocean acidification

Acid rain is precipitation with a pH below the normal value for rainwater, caused by acidic gases dissolving and reacting in atmospheric water. Pure water at 298 K is neutral at pH 7, but normal rain is already slightly acidic because carbon dioxide dissolves to form carbonic acid:

$CO_2(g) + H_2O(l) \rightleftharpoons H_2CO_3(aq)$

Sulfur dioxide and nitrogen oxides make rain more acidic. Sulfur dioxide may come from volcanic activity and from burning sulfur-containing fossil fuels. It dissolves and is oxidized in air to form stronger acids, including sulfuric acid:

$SO_2(g) + H_2O(l) \to H_2SO_3(aq)$

$2H_2SO_3(aq) + O_2(g) \to 2H_2SO_4(aq)$

Nitrogen dioxide also reacts with water to form acids:

$2NO_2(g) + H_2O(l) \to HNO_3(aq) + HNO_2(aq)$

Acid rain damages limestone buildings, corrodes metals, acidifies lakes and harms plants. Ocean acidification is the decrease in ocean pH caused mainly by increasing atmospheric $CO_2$ dissolving in seawater and forming carbonic acid. It is especially serious for organisms such as corals and shellfish because carbonate chemistry is disturbed, making calcification more difficult.

What oxidation state means

An oxidation state is a signed number given to an atom in a species. It shows the charge that atom would have if all the bonds were treated as fully ionic. In this course, oxidation number means the same thing. Write oxidation states with the sign first, then an Arabic numeral: $+2$, $-1$, $+7$, and so on.

This is bookkeeping. It does not mean every compound is ionic. In water, oxygen is more electronegative than hydrogen, so for oxidation-state calculations the bonding electrons are assigned to oxygen. Oxygen is assigned $-2$ and each hydrogen is assigned $+1$.

Rules for deducing oxidation states

Use the rules in a sensible order: fixed rules first, then the sum rule.

1. An atom in a free element has oxidation state 0. This includes single atoms such as Ne, metals such as Mg, and elemental molecules such as $H_2$, $O_2$, $Cl_2$ and $S_8$.
2. The sum of oxidation states in a neutral compound is 0.
3. The sum of oxidation states in a polyatomic ion equals the charge on the ion.
4. Fluorine has oxidation state $-1$ in compounds.
5. Group 1 metals have oxidation state $+1$; group 2 metals have oxidation state $+2$.
6. Oxygen usually has oxidation state $-2$. In peroxides such as $H_2O_2$ or $Na_2O_2$, oxygen is $-1$. In $OF_2$, oxygen is $+2$ because fluorine is more electronegative.
7. Hydrogen is usually $+1$ when bonded to non-metals, but $-1$ in metal hydrides such as $NaH$ and $CaH_2$.

An element has oxidation state zero because atoms of the same element have the same electronegativity. Take $Cl_2$: the bonding electrons are shared equally between identical atoms, so neither atom is assigned a gain or loss of electrons.

A clean method

For $NaH$, sodium is $+1$, so hydrogen must be $-1$. This is the metal hydride exception students often miss.

For $Na_2O_2$, sodium is $+1$. The two sodium ions give $+2$ overall, so the two oxygen atoms must give $-2$ overall. Each oxygen is $-1$ because this is a peroxide.

Oxyanion names

An oxyanion is a polyatomic anion that contains oxygen and at least one other element. Systematic names often show the oxidation state of the non-oxygen atom using Roman numerals in brackets. For example, $SO_4^{2-}$ contains sulfur in oxidation state $+6$, so it may be named sulfate(VI). $SO_3^{2-}$ contains sulfur in oxidation state $+4$, so it may be named sulfate(IV).

Common names are also accepted for familiar oxyanions:

Oxidation states are especially useful in redox chemistry because increases and decreases in oxidation state show where oxidation and reduction have occurred. This leads into analysing redox reactions later.

Why the graph is not a smooth climb

First ionization energy generally rises across a period. Nuclear charge increases, while shielding by inner shells changes very little. The line is not perfectly smooth, though: there are two key drops, from group 2 to group 13, and from group 15 to group 16. Those breaks show that electrons occupy sublevels and orbitals with different energies.

Group 2 to group 13: s to p

Compare beryllium and boron:

Be: $1s^2 2s^2$

B: $1s^2 2s^2 2p^1$

Boron loses a 2p electron. A 2p electron is higher in energy than a 2s electron, and the 2s electrons shield it slightly more from the nucleus. So even though boron has a greater nuclear charge than beryllium, it takes less energy to remove boron’s outer electron. This drop shows that the p sublevel is at a higher energy than the s sublevel in the same main energy level.

Group 15 to group 16: paired-electron repulsion

Now look at nitrogen and oxygen:

N: $1s^2 2s^2 2p^3$

O: $1s^2 2s^2 2p^4$

In nitrogen, the three 2p electrons each occupy a separate p orbital. In oxygen, one 2p orbital has a pair of electrons. The electron removed from oxygen is one of that pair, and paired electrons repel each other because they occupy the same region of space. This repulsion raises the energy of the electron being removed, so less energy is needed than expected.

Be careful with the wording here. Don’t just write “half-filled sublevels are specially stable” as a slogan. Focus on the energy of the electron removed: higher-energy p electrons, and higher-energy paired electrons, are easier to remove.

What counts as a transition element

A transition element is an element whose atoms have an incomplete d-sublevel, or one that forms at least one stable ion with an incomplete d-sublevel. So, being in the d-block is not enough on its own.

Zinc gives a useful contrast. Zinc atoms have $[\text{Ar}] 4s^2 3d^{10}$, and the common $\text{Zn}^{2+}$ ion has $[\text{Ar}] 3d^{10}$. In both cases the d-sublevel is full, so zinc does not fit the definition. Scandium is less clear-cut: $\text{Sc}^{3+}$ is $[\text{Ar}]$, which has no d electron, but compounds containing $\text{Sc}^{2+}$with a$3d^1$configuration are known. The case for including scandium is that it can form an ion with an incomplete d-sublevel; the case against is that its overwhelmingly common chemistry is$\text{Sc}^{3+}$`, which lacks d electrons.

Characteristic properties

Transition elements commonly have these properties:

• Variable oxidation state: they can form ions with different charges because 4s and 3d electrons are relatively close in energy.
• High melting points: many delocalized electrons and small, highly charged metal ions give strong metallic bonding.
• Magnetic properties: many have unpaired d electrons. You do not need to learn types of magnetism, but you should recognize that unpaired electrons are the key idea.
• Catalytic properties: they often provide surfaces or variable oxidation states, giving alternative reaction pathways.
• Formation of coloured compounds: incomplete d-sublevels allow electronic transitions involving visible light.
• Formation of complex ions with ligands: transition metal ions can accept electron pairs from surrounding ions or molecules.

A catalyst is a species that increases the rate of a reaction by providing an alternative pathway with lower activation energy, without being consumed overall. A heterogeneous catalyst is a catalyst in a different physical phase from the reactants. Transition metals often work well as heterogeneous catalysts because reactant molecules can adsorb onto their surfaces, react, and then desorb.

Catalytic converters in vehicles use transition metals such as platinum, palladium and rhodium to help convert harmful exhaust gases into less harmful products. You don’t need to memorize industrial details here; the key classification point is that incomplete d-sublevels help produce a family of recognizable properties.

Electron configurations of first-row transition metal ions

The first-row transition elements are the period 4 d-block elements from scandium to zinc. In neutral atoms, 4s fills before 3d, so the usual configuration is written as $[Ar] 4s^2 3d^x$. There are two well-known exceptions:

Cr: $[Ar] 4s^1 3d^5$

Cu: $[Ar] 4s^1 3d^{10}$

For ions, remove electrons from 4s before 3d. It can seem the wrong way round, since 4s fills first, but after the 3d sublevel contains electrons, the 4s electrons are lost first.

Examples:

Why variable oxidation states are possible

Successive ionization energies are the energies needed to remove electrons one at a time from the same species. Transition elements can often lose different numbers of electrons because the successive ionization energies for removing 4s electrons and several 3d electrons are fairly close. As a result, several cations may be energetically accessible, especially when forming a compound releases energy through bonding or lattice formation.

Chromium shows this well. It can use its 4s electron and several 3d electrons in bonding, so oxidation states from $+1$ to $+6$ are known. Copper has fewer common oxidation states because its successive ionization energies rise more sharply sooner; $Cu^+$ and $Cu^{2+}$ are far more common than higher states.

The pattern to use is:

1. Write the neutral atom configuration, remembering Cr and Cu.
2. Remove 4s electrons first.
3. Remove 3d electrons next.
4. Link the number of electrons removed to the oxidation state of the ion.

For example, vanadium is $[Ar] 4s^2 3d^3$. If it forms $V^{5+}$, all five 4s and 3d electrons have been removed, giving $[Ar]$. This is why $+5$ is a plausible high oxidation state for vanadium compounds.

Complex ions and ligands

A complex ion is an ion with a central metal ion bonded to surrounding ligands by coordinate bonds. A ligand is an ion or molecule that donates a lone pair of electrons to a central metal ion. A coordinate bond is a covalent bond where both bonding electrons come from the same atom.

Transition metal ions can form complex ions because they accept lone pairs from ligands such as $H_2O$, $NH_3$, $Cl^-$ or $CN^-$. The detailed reactions between ligands and transition metal ions are covered elsewhere. For this topic, the key idea is that ligand bonding changes the energies of the d orbitals.

Why many complexes are coloured

In an isolated transition metal ion, the five d orbitals all have the same energy. Orbitals like this are called degenerate orbitals. Once ligands bond to the metal ion, the ligand electron pairs interact with the d orbitals and split the d-sublevel into two sets with different energies. You are not expected to know the different splitting patterns or how they relate to coordination number.

If the energy gap between the split d orbitals matches the energy of visible light, an electron can absorb that light and move from a lower d orbital to a higher d orbital. The light that isn’t absorbed reaches your eye. So the colour you observe is complementary to the colour absorbed.

A complementary colour is a colour opposite another colour on a colour wheel. If a complex absorbs yellow light, it appears violet. If it absorbs orange-red light, it appears blue-green. Use the colour wheel in the data booklet instead of trying to memorise every pairing.

Wavelength and frequency

The relationship between wavelength and frequency is:

A shorter wavelength means a higher frequency. Since higher-frequency light has higher energy, a larger d-orbital splitting leads to absorption of shorter-wavelength light. Remember that the observed colour is the complementary colour, not the absorbed colour.

For example, if a complex appears red, use the colour wheel to find the complementary colour absorbed, then use its wavelength in $c = \lambda f$ to calculate the frequency. Watch the units: nanometres must be converted to metres before substitution.

What affects colour

The colour of a transition metal complex depends on factors that change the size of the d-orbital splitting, especially:

• the identity of the metal ion;
• the oxidation state, or charge, of the metal ion;
• the identity of the ligand.

Stronger ligand interactions usually give larger splitting. With larger splitting, higher-energy light is absorbed, so the absorbed light has shorter wavelength and higher frequency. The colour seen shifts to the complementary colour.

Colorimetry and concentration

Colorimetry is an analytical technique used to determine concentration by measuring how much light a coloured solution absorbs. Spectrophotometry is a related technique that measures absorption at selected wavelengths, often with more precise wavelength control.

In a typical experiment, you prepare standard solutions of known concentration, measure their absorbance at a suitable wavelength, and plot a calibration curve of absorbance against concentration. You then measure the absorbance of the unknown solution and find its concentration by interpolation from the graph. Good technique includes calibrating the instrument, using suitable dilution, adding uncertainty bars where appropriate, and checking that the best-fit line behaves sensibly, often passing close to the origin for dilute solutions.