Predicting Chemical Reactions — Groups and Trends
The Periodic Table and Groups
The modern periodic table arranges elements in order of atomic number. Elements in the same group (column) have the same number of outer-shell electrons, which is why they have similar chemical properties. Elements in the same period (row) have the same number of electron shells. Mendeleev’s original table was based on atomic mass and left gaps for undiscovered elements. His predictions about these missing elements were later confirmed, validating the periodic arrangement. The table was subsequently refined using atomic number (not atomic mass) as the fundamental organising principle, which resolved anomalies where elements would have been placed in the wrong group if ordered by mass alone.
Group 1 — The Alkali Metals
Group 1 elements (lithium, sodium, potassium, rubidium, caesium, francium) are called the alkali metals. They all have one electron in their outer shell, which they lose easily to form +1 ions. They are soft metals that can be cut with a knife, have low melting points compared to most metals, and have low densities — lithium, sodium and potassium all float on water. They must be stored under oil to prevent them reacting with moisture and oxygen in the air.
Group 1 metals react with water to produce a metal hydroxide (an alkaline solution, hence “alkali metals”) and hydrogen gas. For example: \(2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)\). Lithium fizzes gently, sodium melts into a ball and fizzes vigorously, and potassium reacts so vigorously it catches fire with a lilac flame. Reactivity increases down the group because the outer electron is further from the nucleus and more shielded by inner electrons, making it easier to lose.
Group 1: reactivity increases down the group because the outer electron is further from the nucleus, more shielded, so easier to lose. Melting points decrease down the group.
Group 7 — The Halogens
Group 7 elements (fluorine, chlorine, bromine, iodine, astatine) are called the halogens. They have seven electrons in their outer shell and need to gain one more electron to achieve a full outer shell, forming −1 ions. They exist as diatomic molecules (\(F_2\), \(Cl_2\), \(Br_2\), \(I_2\)). Halogens react with metals to form ionic compounds called halide salts (e.g. sodium chloride, NaCl) and react with non-metals to form covalent compounds.
Physical properties change down the group as the molecules get larger and intermolecular forces become stronger: fluorine is a pale yellow gas, chlorine is a green gas, bromine is a red-brown liquid, and iodine is a grey solid that sublimes to form a purple vapour. Melting and boiling points increase down the group. Reactivity decreases going down Group 7 because the outer shell is further from the nucleus, so the attraction for an incoming electron is weaker and it is harder to gain that extra electron.
A more reactive halogen can displace a less reactive halogen from a solution of its salt. For example, chlorine displaces bromine from potassium bromide solution: \(Cl_2(aq) + 2KBr(aq) \rightarrow 2KCl(aq) + Br_2(aq)\). The solution changes from colourless to orange-brown as bromine is released. Similarly, chlorine displaces iodine from potassium iodide, turning the solution brown. However, bromine cannot displace chlorine from potassium chloride because bromine is less reactive than chlorine. These displacement reactions provide experimental evidence for the reactivity trend.
Group 0 — The Noble Gases
Group 0 elements (helium, neon, argon, krypton, xenon, radon) are the noble gases. They have full outer electron shells, which makes them very stable and unreactive. They exist as individual atoms (monatomic) rather than molecules because they have no tendency to share, gain or lose electrons. Boiling points increase down the group as the atoms become larger and intermolecular forces increase.
Noble gases are used where unreactive gases are needed: helium for balloons and airships (low density, non-flammable), neon in advertising signs (produces a distinctive red-orange glow when electricity passes through it), and argon in light bulbs and welding (prevents oxidation of the hot filament or metal).
Noble gases are unreactive because they have complete outer electron shells. They do not need to form bonds. Boiling points increase down the group as atomic size and intermolecular forces increase.
Going down any group, atoms have more electron shells. In Group 1, the single outer electron is progressively further from the positive nucleus and more shielded by inner electrons — so it is easier to remove (lose), making the element more reactive. In Group 7, the outer shell is further from the nucleus, so the attraction for an incoming electron is weaker — making it harder to gain an electron, so reactivity decreases. In Group 0, the full outer shell provides stability regardless of size, so all noble gases are unreactive. These trends can be explained entirely by the electronic structure of the atoms. Understanding this pattern means you can predict the properties of any element from its position in the periodic table, even if you have never studied it before.