Chemical similarity within a group

A group is a vertical column in the periodic table. Elements in the same group have the same number of electrons in their outer shell.

Outer-shell electrons determine how an atom takes part in chemical reactions: they are the electrons that can be lost, gained or shared. Because elements in a group have the same number of outer-shell electrons, they tend to react in similar ways and have similar chemical properties.

For example, lithium and sodium are both in Group 1 and each has one outer-shell electron. Each can lose that electron to form an ion with a charge of $$+1$$. Knowing that another element is in Group 1 therefore helps predict that it can behave similarly.

Elements in the same group are similar, but their reactions are not necessarily identical.

Atoms and elements

All substances are made from atoms. An atom is the smallest part of an element that can exist as that element.

An element is a pure substance made from only one type of atom. For example, a sample of the element carbon contains only carbon atoms. The atoms in an element may be joined together, but they are still all the same kind of atom.

There are about 100 different elements. Their atoms are the basic building blocks from which other substances are made.

Electron arrangement and the reactivity of metals and non-metals

Atoms react according to the arrangement of their electrons, especially the electrons in their outer shell. These are the electrons that can be lost, gained or shared when atoms react.

An element’s atomic number is the number of protons in each of its atoms. A neutral atom has equal numbers of protons and electrons, so its atomic number also tells us how many electrons it has. Those electrons occupy shells around the nucleus. As atomic number changes from one element to the next, the electron arrangement changes too.

Outer-shell electrons and periodic-table position

An element’s position in the periodic table is linked to its electron arrangement, and therefore to how it tends to react.

• Metals, found mostly on the left, tend to have only a few outer-shell electrons. They tend to lose these electrons when they react.
• Non-metals, found mostly on the right, often have nearly full outer shells. They tend to gain electrons or share electrons with other atoms when they react.

Losing, gaining or sharing electrons can give atoms a more stable electron arrangement, often with a full outer shell. These different tendencies help explain why metals and non-metals show different reaction patterns.

For example, a neutral sodium atom has 11 electrons arranged as 2, 8, 1. It can lose its one outer-shell electron. A neutral chlorine atom has 17 electrons arranged as 2, 8, 7, so it can gain one electron to complete its outer shell. When sodium reacts with chlorine, an electron can transfer from a sodium atom to a chlorine atom.

Periodicity of elemental properties

When elements are arranged in order of increasing atomic number, patterns in their properties recur. This recurrence is called periodicity. The patterns are not exact repetitions: elements can have similar behaviour without having identical properties.

Periodicity and chemical properties

An atom’s outer-shell electrons affect how it reacts. As atomic number increases, outer-shell electron arrangements follow a repeating pattern. Elements in the same group of the periodic table have similar outer-shell arrangements, so they often have similar chemical properties.

For example, the elements in Group 1 each have one outer-shell electron. They react in similar ways, although they do not react equally vigorously. The recurrence of similar outer-shell arrangements helps explain why chemical behaviour follows a periodic pattern.

Periodicity and physical properties

Physical properties, including electrical conductivity and melting point, also show patterns across the periodic table. These properties depend on how atoms are bonded and arranged, which in turn is influenced by atomic structure.

For example, sodium is a metal that conducts electricity, while sulfur is a non-metal that does not. Their different structures help explain this difference: sodium has electrons that can move through its metallic structure, whereas sulfur does not have mobile charged particles.

Melting point also depends on structure and bonding. Silicon has a large covalent structure with strong bonds throughout it, giving it a high melting point. Sulfur consists of small molecules with weaker forces between them, so it has a lower melting point. Such differences form part of the patterns seen across the periodic table, though the values do not repeat identically in each period.

Ionic bonding

Ionic bonding begins when electrons are transferred from atoms of one element to atoms of another. This transfer produces ions: particles with an electrical charge.

An atom that loses electrons has more protons than electrons, so it becomes a positive ion. An atom that gains electrons has more electrons than protons, so it becomes a negative ion.

An ionic bond is the strong electrostatic attraction between oppositely charged ions. The bond is the attraction between the ions, rather than the transfer of electrons itself.

For example, a sodium atom can transfer one electron to a chlorine atom. The sodium atom becomes a positive sodium ion, $$\ce{Na+}$$, and the chlorine atom becomes a negative chloride ion, $$\ce{Cl-}$$. The oppositely charged ions attract each other.

Covalent bonding

A covalent bond forms when two atoms share a pair of electrons. The shared electrons are attracted to the positively charged nuclei of both atoms. These attractions hold the atoms together.

For example, two hydrogen atoms can each contribute one electron to a shared pair, forming a hydrogen molecule, $$\ce{H2}$$. The shared pair is attracted to both hydrogen nuclei, holding the atoms together.

In a chemical reaction, atoms can form covalent bonds by sharing electrons. Covalent bonds can also be rearranged: bonds between some atoms break and new bonds form between atoms, producing different substances.

Molecular shape and behaviour

A molecule is a group of atoms bonded together in a particular three-dimensional arrangement. Its shape describes how those atoms are positioned in space, not just which atoms it contains.

The arrangement affects how other particles can approach and interact with the molecule. For example, the position of an atom within a molecule can affect how easily another particle can reach it during a chemical reaction. Shape can therefore influence the molecule’s chemical behaviour.

Shape can also influence physical behaviour. Differently shaped molecules may pack together differently, affecting the attractions between them and properties such as melting point.

A molecule’s behaviour depends on both which atoms are present and how they are arranged. Even molecules containing the same kinds and numbers of atoms can behave differently if their atoms are arranged differently.

Giant structures and properties

A giant structure contains very large numbers of particles joined in an extended arrangement. It does not consist of separate small molecules. The particles, the way they are arranged and the bonding between them affect the structure’s properties.

Strength and melting point

A force may change a material’s shape by moving its particles relative to one another. How easily this happens depends on the arrangement of the particles and the bonding between them. Strong bonding can resist changes to the structure, but strength also depends on how the particles are arranged.

Melting requires enough energy to overcome the forces holding particles in their fixed arrangement. Where strong bonds extend throughout a giant structure, substantial energy is needed to overcome enough of them for the substance to melt. Such structures therefore tend to have high melting points.

Electrical conductivity

Electrical conduction requires charged particles that can move through the substance and carry charge. The presence of charged particles alone is not enough: if they are fixed in position, they cannot carry charge through the structure. Whether a giant structure conducts therefore depends on both the charged particles it contains and whether those particles are free to move.

Activation energy and reaction rates

Reacting particles must overcome an energy barrier before they can react. When particles collide, they need enough energy to overcome this barrier. The minimum energy needed is called the activation energy.

A collision with less than the activation energy cannot lead to a reaction. A collision with sufficient energy can lead to a reaction, although not every such collision necessarily does.

Activation energy and reaction rate

At the same temperature, collisions do not all have the same energy. If a reaction has a higher activation energy, a smaller proportion of collisions generally have enough energy to overcome the barrier. Fewer collisions can therefore lead to reaction, contributing to a slower reaction rate.

Proton-transfer reactions

A proton-transfer reaction occurs when one species transfers a hydrogen ion, $$\ce{H+}$$, to another species. A hydrogen ion is a proton.

The species that gives away the proton is an acid, or proton donor. The species that receives it is a base, or proton acceptor. Both species change: the acid loses a proton, while the base gains one.

For example, hydrogen chloride transfers a proton to ammonia:

Here, $$\ce{HCl}$$ is the acid because it donates a proton and becomes $$\ce{Cl-}$$. Ammonia, $$\ce{NH3}$$, is the base because it accepts the proton and becomes $$\ce{NH4+}$$.

Electron-transfer reactions

In an electron-transfer reaction, one species loses electrons and another gains them. A species can be an atom, ion or molecule.

Oxidation is the loss of electrons. Reduction is the gain of electrons. The two changes occur together: the electrons lost by one species are gained by another. The combined process is called a redox reaction.

For example, a magnesium atom can lose two electrons to form a magnesium ion, $$\ce{Mg^{2+}}$$. A copper ion, $$\ce{Cu^{2+}}$$, can gain those two electrons to form a copper atom. Magnesium is oxidised because it loses electrons; the copper ion is reduced because it gains them.

Conservation of energy in reactions

During a chemical reaction, energy can be transferred between the reacting substances and their surroundings. It can also change form, but it is neither created nor destroyed. The total energy of the reacting system and its surroundings remains constant: energy gained by one is lost by the other.

Bond changes and energy transfer

A chemical reaction involves breaking bonds in the reactants and forming bonds in the products. Breaking bonds requires energy, while forming bonds releases energy. The overall energy transfer depends on the balance between these two changes:

• In an exothermic reaction, forming bonds releases more energy than is needed to break bonds. Energy is transferred overall from the reacting system to the surroundings.
• In an endothermic reaction, breaking bonds requires more energy than is released when bonds form. Energy is transferred overall from the surroundings to the reacting system.

In both cases, energy is conserved. An exothermic reaction does not create energy, and an endothermic reaction does not destroy it.