Structure of atoms: Class 9 Chemistry

Key points

An atom is the smallest particle of an element that keeps the properties of that element. Early Greek thinkers and later John Dalton described atoms as tiny solid particles. Experiments in the 19th and 20th centuries showed that the atom is made of smaller particles, so Dalton's idea of an indivisible atom had to be changed.

The three main subatomic particles are:

• Electron: negative charge (−1.6 × 10⁻¹⁹ C), mass about 1/1836 of a proton. Discovered by J. J. Thomson using cathode rays in a discharge tube.
• Proton: positive charge (+1.6 × 10⁻¹⁹ C), relative mass 1 u. Found by Goldstein using anode rays (canal rays) and named by Rutherford.
• Neutron: no charge, relative mass about 1 u. Discovered by James Chadwick in 1932.

Cathode rays travel in straight lines from the cathode to the anode in a tube at low pressure. They are deflected by electric and magnetic fields towards the positive side, which shows that they carry negative charge. Their nature does not depend on the gas or the metal of the cathode, so electrons are present in all atoms.

Rutherford's model came from the gold foil experiment. Alpha particles were fired at a very thin gold foil. Most passed straight through, a few were deflected through large angles and very few bounced back. This showed that the atom is mostly empty space and that its positive charge and almost all of its mass are packed in a tiny centre called the nucleus. Electrons move around the nucleus.

Rutherford's model had a fault. According to classical physics, an electron moving in a circle should lose energy and spiral into the nucleus, so the atom would collapse. The model also could not explain the line spectrum of hydrogen.

Bohr's model (1913) fixed this with these ideas:

• Electrons move around the nucleus in fixed circular paths called orbits, shells or energy levels.
• While an electron stays in one orbit it neither gains nor loses energy.
• Orbits are numbered n = 1, 2, 3 and so on (K, L, M, N shells) counting outwards. Energy increases with n.
• An electron absorbs energy to jump to a higher orbit and releases energy, as light, when it falls to a lower one.

The atomic number (Z) is the number of protons in the nucleus. It is the identity of an element. In a neutral atom the number of electrons equals Z. The mass number (A) is the total number of protons and neutrons, which are together called nucleons. The number of neutrons is A − Z. An atom is written with A as a superscript and Z as a subscript before the symbol, for example ²³₁₁Na has 11 protons, 11 electrons and 12 neutrons.

Isotopes are atoms of the same element with the same atomic number but different mass numbers, because they have different numbers of neutrons. Hydrogen has three isotopes: protium (no neutron), deuterium (1 neutron) and tritium (2 neutrons). Isotopes have the same chemical properties because they have the same electrons, but slightly different physical properties such as density.

Some uses of isotopes:

• Cobalt-60 gives gamma rays used to treat cancer and to sterilise medical tools.
• Iodine-131 is used to treat and study thyroid disease.
• Carbon-14 is used to find the age of fossils and old wooden objects.
• Uranium-235 is the fuel of nuclear power stations.

Relative atomic mass is the average mass of the isotopes of an element, taking account of how common each is, compared with 1/12 of the mass of a carbon-12 atom. It is measured in atomic mass units (u). This is why chlorine has a relative atomic mass of 35.5 and not a whole number.

Electrons fill shells in order of increasing energy. A shell can hold at most 2n² electrons: 2 in K, 8 in L, 18 in M and 32 in N. For the first 20 elements the outer shell holds at most 8 electrons. Examples: Na is 2,8,1; Cl is 2,8,7; Ca is 2,8,8,2. Electrons in the outermost shell are valence electrons and decide chemical behaviour. Atoms with a full outer shell, such as helium and neon, are very stable and unreactive.

Study tip: Practise writing electronic configurations for elements 1 to 20 until they are automatic. In descriptive answers, name the scientist with each discovery (Thomson, Rutherford, Chadwick, Bohr) and state the observation first, then the conclusion. Show the working in every isotope calculation.

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Common mistakes

• Confusing atomic number with mass number. The atomic number counts protons only.
• Forgetting to subtract electrons for a positive ion or to add them for a negative ion.
• Putting more than 8 electrons in the outer shell for the first 20 elements, or filling M before L is full.
• Saying isotopes have different chemical properties. They differ only in neutrons.
• Writing the relative atomic mass of an element as the mass number of one isotope.

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