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 physical properties

When elements are arranged in order of increasing atomic number, their physical properties follow patterns that recur. This repetition is called periodicity. For example, melting point changes across a period as the structures formed by the elements change; a related pattern appears in a later period.

The repetition is linked to the arrangement of electrons. Elements in the same group have similar outer-electron arrangements, which influence how their atoms bond. However, elements in later periods have more occupied electron shells. For example, lithium and sodium each have one outer electron, but lithium has two occupied shells and sodium has three.

Melting points across a period

The elements in the third period show how bonding and structure affect melting point:

• Sodium, magnesium and aluminium have metallic structures. Their atoms are held together by metallic bonds, which must be overcome when the metals melt.
• Silicon has a giant covalent structure: many atoms are joined in a network by strong covalent bonds. Breaking enough of these bonds to melt silicon requires a great deal of energy, so its melting point is high.
• Phosphorus, sulfur and chlorine form simple molecules. Melting overcomes forces between the molecules, rather than breaking the covalent bonds within them. These forces require much less energy to overcome, so their melting points are much lower than silicon’s.
• Argon exists as separate atoms. Only weak forces between the atoms need to be overcome when it melts, so its melting point is very low.

A related change from metallic to giant covalent and then to simple molecular or atomic structures occurs across the second period. The melting-point pattern is not identical in each period: different elements and structures have different bond strengths and forces. The recurring pattern comes from the recurrence of outer-electron arrangements that influence bonding.

Ionic bonding

Ionic bonding begins when one or more electrons are transferred from one atom to another. The atoms become ions: particles with an electrical charge.

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

For example, a sodium atom, $$\ce{Na}$$, can transfer one electron to a chlorine atom, $$\ce{Cl}$$. The sodium atom becomes $$\ce{Na+}$$, while the chlorine atom becomes $$\ce{Cl-}$$.

The positive and negative ions attract each other. This strong electrostatic attraction between oppositely charged ions is an ionic bond. The attraction holds ions together in an ionic structure; the bond is the attraction between the ions, not the transfer of electrons itself.

Covalent bonding and electron-sharing reactions

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 forms the covalent bond between the atoms.

Atoms can share electron pairs in molecules, which are separate groups of bonded atoms. They can also share electron pairs in a giant covalent structure, a continuous network of covalently bonded atoms rather than separate molecules.

Electron sharing in reactions

During a reaction, electrons can be redistributed between atoms. Existing covalent bonds can break, and new covalent bonds can form as electron pairs are shared between different atoms. This changes which atoms are bonded together and produces different substances.

Changes involving electron sharing can occur alongside other reaction mechanisms.

Molecular shape and behaviour

A molecule is a group of atoms bonded together. The atoms occupy particular positions in three-dimensional space, giving the molecule its shape. To picture a molecule, it is therefore not enough to know which atoms it contains: their arrangement also matters.

The atoms and bonds in a molecule affect how electric charge is distributed within it. The molecule’s shape determines where regions of charge lie relative to one another. Depending on their arrangement, the effects of these regions may reinforce or cancel one another. This influences how the molecule attracts or is attracted to other particles.

Shape also affects how closely molecules can approach and fit alongside one another. This can change the strength of interactions between them and so influence physical properties, such as melting and boiling points. In chemical reactions, shape can affect how particles approach one another and whether the relevant parts can come into contact.

Shape does not determine behaviour on its own. A molecule’s properties and reactions depend on its shape as well as the atoms and bonds it contains.

Giant structures and their properties

A giant structure is a large, repeating arrangement of bonded particles, rather than a collection of separate small molecules. Its properties depend on which particles it contains, how they are arranged and how they are bonded.

In giant ionic structures, the particles are positive and negative ions. In metallic structures, positive metal ions are held together by delocalised electrons. In giant covalent structures, atoms are joined by covalent bonds throughout the structure.

Giant structures: bonding and melting point

Strong bonding throughout a giant structure can give it a high melting point. A large amount of energy is needed to overcome enough of the bonding for the particles to move out of their fixed arrangement.

Giant structures: electrical conductivity

A substance conducts electricity when it contains charged particles that can move through it.

• In a solid ionic structure, ions are fixed in position and cannot carry charge through the substance. When the substance is molten, the ions can move, so it conducts electricity.
• In a metallic structure, delocalised electrons can move through the structure and carry charge, so the metal conducts electricity.

Giant covalent structures: different arrangements

Different arrangements of atoms can give giant covalent structures different properties. Diamond and graphite are both made of carbon atoms, but the atoms are arranged differently.

In diamond, each carbon atom is covalently bonded to four others. Strong bonds extend throughout the structure, making diamond very hard. Diamond does not conduct electricity because it has no mobile charged particles.

In graphite, each carbon atom is covalently bonded to three others in layers. The layers can slide over one another because the forces between them are weak, making graphite soft. Graphite also conducts electricity because it has delocalised electrons that can move along the layers.

Activation energy and reaction rate

For a reaction to occur, reacting particles must get past an energy barrier. This requires energy to overcome the initial resistance to rearranging the particles into products. The minimum energy needed for the reaction to occur is called the activation energy.

Successful collisions

A collision between reacting particles does not always cause a reaction. For a collision to be successful, the particles must:

• collide with at least the activation energy;
• meet in a suitable orientation, so the relevant parts of the particles can interact.

A collision with too little energy cannot overcome the barrier. Even if the particles have enough energy, a collision may not lead to a reaction if they meet in an unsuitable orientation.

Activation energy and reaction rate

The reaction rate depends on how frequently successful collisions occur. More successful collisions each second mean that reactants are converted into products faster.

Reactions can have different activation energies. If particles collide with the same range of energies, a higher activation energy means that fewer collisions have enough energy to overcome the barrier. This can result in fewer successful collisions each second and a slower reaction. A lower barrier can allow more collisions to have sufficient energy, although the particles must still meet in a suitable orientation.

Proton transfer reactions

A proton transfer reaction occurs when a hydrogen ion, $$\ce{H+}$$, moves from one species to another. A hydrogen ion is called a proton because it is a hydrogen atom that has lost its electron.

The species that gives up the proton is an acid. The species that receives the proton is a base. Both species change: the acid has one fewer hydrogen atom, while the base has gained one.

Proton transfer in aqueous neutralisation

In an aqueous solution, a proton can be carried by a water molecule as a hydronium ion, $$\ce{H3O+}$$. A hydroxide ion, $$\ce{OH-}$$, can accept that proton:

The hydronium ion is the acid: it donates a proton and becomes a water molecule. The hydroxide ion is the base: it accepts the proton and also becomes a water molecule.

The transfer changes the species present, but atoms and total charge are conserved. There are four hydrogen atoms and two oxygen atoms on each side. The total charge is zero on each side: the positive and negative charges on the left balance, and the water molecules on the right are neutral.

Electron transfer reactions

In an electron transfer reaction, electrons move from one species, such as an atom or ion, to another. The species that loses electrons is oxidised; the species that gains electrons is reduced.

Oxidation and reduction occur together: the electrons lost by one species are accepted by another. The number of electrons lost must therefore equal the number gained.

Electron transfer and ion formation

Electron transfer can turn atoms into ions. When an atom loses negatively charged electrons, it forms a positive ion. When an atom gains electrons, it forms a negative ion.

For example, a sodium atom can transfer an electron to a chlorine atom. Sodium loses an electron and is oxidised to $$\ce{Na+}$$; chlorine gains that electron and is reduced to $$\ce{Cl-}$$. The oppositely charged ions attract, forming the ionic substance sodium chloride.

Energy conservation in chemical reactions

In a chemical reaction, energy is not created or destroyed. It is transferred between the reacting system and its surroundings. The system is the substances taking part in the reaction; the surroundings are everything outside it.

Bond changes and energy

During a reaction, bonds in the reactants are broken and bonds in the products are formed. Breaking bonds requires energy, whereas forming bonds releases energy. The overall energy transfer depends on the balance between these two amounts.

Energy transfer to or from the surroundings

If bond formation releases more energy than bond breaking requires, the system transfers energy to the surroundings. This is an exothermic reaction.

If bond breaking requires more energy than bond formation releases, the system takes in energy from the surroundings. This is an endothermic reaction.

In either case, the energy gained by one is equal to the energy lost by the other. The total energy of the system and surroundings therefore remains constant.