Atomic Structure and the Periodic Table

Atoms are the basic building blocks of all substances. An atom is the smallest part of an element that can exist while retaining the properties of that element. Each element is represented by a chemical symbol, such as O for oxygen and Na for sodium. Elements contain only one type of atom, while compounds contain atoms of two or more different elements chemically combined in fixed proportions. Compounds are formed by chemical reactions and can only be separated into their constituent elements by chemical reactions.

Chemical reactions involve the formation of new substances and normally involve a detectable energy change. Reactions can be represented using word equations, symbol equations and chemical formulae. Students should know the names and symbols of the first 20 elements in the periodic table as well as the elements in Groups 1 and 7 required by the specification. Chemical equations should be balanced because atoms are neither created nor destroyed during a chemical reaction.

A mixture contains two or more elements or compounds that are not chemically bonded together. Each substance in a mixture retains its own chemical properties, and no new substance is formed simply by mixing them. Mixtures can therefore be separated using physical techniques such as filtration, crystallisation, simple distillation, fractional distillation and chromatography. The appropriate separation method depends on the physical properties of the substances present.

The accepted model of the atom developed as new experimental evidence became available. Early scientists considered atoms to be tiny indivisible spheres. After the electron was discovered, the plum pudding model proposed that an atom was a sphere of positive charge containing negative electrons. Results from the alpha-particle scattering experiment showed that most of the atom is empty space and that its positive charge and most of its mass are concentrated in a very small central nucleus, leading to the nuclear model.

Further developments refined the nuclear model of the atom. Niels Bohr proposed that electrons orbit the nucleus at specific distances or energy levels, and his theoretical calculations agreed with experimental observations. Later evidence showed that the positive charge of the nucleus could be divided into smaller particles called protons. James Chadwick later provided evidence for neutrons in the nucleus, completing the basic modern model of atomic structure used at GCSE level.

Atoms contain three main subatomic particles: protons, neutrons and electrons. A proton has a relative charge of +1 and a relative mass of 1, a neutron has no electrical charge and a relative mass of 1, and an electron has a relative charge of -1 and a very small relative mass. In a neutral atom, the number of electrons equals the number of protons, so the overall electrical charge of the atom is zero.

The atomic number of an element is the number of protons in the nucleus of each atom of that element. All atoms of the same element have the same number of protons, so the atomic number uniquely identifies the element. The mass number is the total number of protons and neutrons in the nucleus. The number of neutrons can therefore be calculated by subtracting the atomic number from the mass number.

Atoms are extremely small, with a typical radius of about 0.1 nm, or approximately 1 × 10^-10 m. The nucleus is much smaller, with a radius less than 1/10,000 of the radius of the atom, approximately 1 × 10^-14 m. Almost all of the mass of an atom is concentrated in its nucleus because protons and neutrons have much greater masses than electrons.

Isotopes are atoms of the same element that contain the same number of protons but different numbers of neutrons. Because they have the same atomic number, isotopes are still atoms of the same element, but they have different mass numbers. Nuclear symbols can be used to show the mass number and atomic number of an isotope. From these values, the numbers of protons, neutrons and electrons in an atom or ion can be calculated.

Relative atomic mass is the weighted mean mass of the atoms of an element compared with one-twelfth of the mass of a carbon-12 atom. It takes into account both the masses of the different isotopes of an element and their relative abundances. To calculate relative atomic mass from percentage abundances, multiply each isotope mass by its percentage abundance, add the results and divide the total by 100.

Electrons occupy energy levels or shells around the nucleus, beginning with the lowest available energy level. For the first 20 elements, the first shell holds up to two electrons, the second up to eight and the third is filled according to the GCSE model. Electronic structures can be written using numbers, such as 2,8,1 for sodium, or shown using shell diagrams. The electronic structure of an atom determines many of its chemical properties.

The periodic table arranges elements in order of increasing atomic number. Elements with similar chemical properties are placed in the same vertical groups because they have the same number of electrons in their outer shell. The position of an element in the periodic table is therefore linked to its electronic structure. This allows chemists to use periodic trends to predict possible reactions and properties of elements.

The modern periodic table developed from earlier attempts to classify elements. Before the discovery of protons, neutrons and electrons, elements were largely arranged according to atomic weight. Dmitri Mendeleev improved the table by leaving gaps for undiscovered elements and sometimes changing the order of elements when their chemical properties indicated a better position. Later discoveries matched many of his predictions, and knowledge of isotopes helped explain why strict atomic-weight order was sometimes inappropriate.

Metals are generally found on the left and towards the bottom of the periodic table, while non-metals are found mainly towards the right and top. Metals tend to form positive ions, whereas non-metals do not normally form positive ions. Group 0 elements are noble gases with full outer electron shells and are therefore very unreactive. Their boiling points increase down the group as relative atomic mass increases.

Group 1 elements are alkali metals with one electron in their outer shell, and their reactivity increases down the group. Students should know the reactions of the first three Group 1 metals with water, oxygen and chlorine. Group 7 elements are halogens with seven outer-shell electrons; their reactivity decreases down the group, while melting and boiling points increase. A more reactive halogen can displace a less reactive halogen from a solution of its salt. Transition metals differ from Group 1 metals by generally being denser, stronger, harder and less reactive, having higher melting points, forming ions with different charges, producing coloured compounds and often acting as catalysts.