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Basics of Chemistry

Electrons and nuclei

The familiar planetary model of the atom was proposed by Rutherford in 1912 following experiments by Geiger and Marsden showing that nearly all the mass of an atom was concentrated in a positively charged nucleus. Negatively charged electrons are attracted to the nucleus by the electrostatic force and were considered by Rutherford to ‘orbit’ it in a similar way to the planets around the Sun.

It was soon realized that a proper description of atoms required the quantum theory; although the planetary model remains a useful analogy from the macroscopic world, many of the physical ideas that work for familiar objects must be abandoned or modified at the microscopic atomic level. The lightest atomic nucleus (that of hydrogen) is 1830 times more massive than an electron.

The size of a nucleus I around 10−15 m (1 FM), a factor of 105 smaller than the apparent size of an atom, as measured by the distances between atoms in molecules and solids. Atomic sizes are determined by the radii of the electronic orbits, the electron itself having apparently no size at all. Chemical bonding between atoms alters the motion of electrons, the nuclei remaining unchanged. Nuclei retain the ‘chemical identity’ of an element, and the occurrence of chemical elements depends on the existence of stable nuclei.

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Nuclear structure

Nuclei contain positively charged protons and uncharged neutrons; these two particles with about the same mass are known as nucleons. The number of protons in the atomic number of an element (Z), and is matched in a neutral atom by the same number of electrons. The total number of nucleons is the mass number and is sometimes specified by a superscript on the symbol of the element.

Thus 1H has a nucleus with one proton and no neutrons, 16O has eight protons and eight neutrons, 208Pb has 82 protons and 126 neutrons. Protons and neutrons are held together by an attractive force of extremely short-range, called the strong interaction. Opposing this is the longer-range electrostatic repulsion between protons. The balance of the two forces controls some important features of nuclear stability.

• Whereas lighter nuclei are generally stable with approximately equal numbers of protons and neutrons, heavier ones
have a progressively higher proportion of neutrons (e.g. compare 16O with 208Pb).
• As Z increases the electrostatic repulsion comes to dominate, and there is a limit to the number of stable nuclei, all
elements beyond Bi (Z=83) being radioactive (see below).

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As with electrons in atoms, it is necessary to use the quantum theory to account for the details of nuclear structure and stability. It is favorable to ‘pair’ nucleons so that nuclei with even numbers of either protons or neutrons (or both) are generally more stable than ones with odd numbers.

The shell model of nuclei, analogous to the orbital picture of atoms (see Topics A2 and A3) also predicts certain magic numbers of protons or neutrons, which give extra stability. These are 16O and 208Pb are examples of nuclei with magic numbers of both protons and neutrons.

Trends in the stability of nuclei are important not only in determining the number of elements and their isotopes (see below) but also in controlling the proportions in which they are made by nuclear reactions in stars. These determine the abundance of elements in the Universe as a whole (see Topic J1).

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Isotopes

Atoms with the same atomic number and different numbers of neutrons are known as isotopes. The chemical properties of an element are determined largely by the charge on the nucleus, and different isotopes of an element have very similar chemical properties. They are not quite identical, however, and slight differences in chemistry and in physical properties allow isotopes to be separated if desired. Some elements have only one stable isotope (e.g. 19F, 27Al, 31P), others may have several (e.g. 1H and 2H, the latter also being called deuterium, 12C, and 13C); the record is held by tin (Sn), which has no fewer than 10.

Natural samples of many elements, therefore, consist of mixtures of isotopes in nearly fixed proportions reflecting the ways in which these were made by nuclear synthesis. The molar mass (also known as relative atomic mass, RAM) of elements is determined by these proportions. For many chemical purposes, the existence of such isotopic mixtures can be ignored, although it is occasionally significant.

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• Slight differences in chemical and physical properties can lead to small variations in the isotopic composition of natural samples. They can be exploited to give geological information (dating and origin of rocks, etc.) and lead to small variations in the molar mass of elements.

• Some spectroscopic techniques (especially nuclear magnetic resonance, NMR, see Topic B7) exploit specific properties of particular nuclei. Two important NMR nuclei are 1H and 13C. The former makes up over 99.9% of natural hydrogen, but 13C is present as only 1.1% of natural carbon. These different abundances are important both for the sensitivity of the technique and the appearance of the spectra.

• Isotopes can be separated and used for specific purposes. Thus the slight differences in chemical behavior between normal hydrogen (1H) and deuterium (2H) can be used to investigate the detailed mechanisms of a chemical reaction involving hydrogen atoms.

In addition to stable isotopes, all elements have unstable radioactive ones (see below). Some of these occur naturally, others can be made artificially in particle accelerators or nuclear reactors. Many radioactive isotopes are used in chemical and biochemical research and for medical diagnostics.
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Radioactivity

Radioactive decay is a process whereby unstable nuclei change into more stable ones by emitting particles of different kinds. Alpha, beta and gamma (α, β, and γ) radiation were originally classified according to its different penetrating power. The processes involved are illustrated in Fig. 1.

• An α particle is a 4He nucleus and is emitted by some heavy nuclei, giving a nucleus with Z two units less and mass
number four units less. For example, 238U (Z=92) undergoes a decay to give (radioactive) 234Th (Z=90).

• A β particle is an electron. Its emission by a nucleus increases Z by one unit but does not change the mass number.
Thus 14C (Z=6) decays to (stable) 14N (Z=7).

• γ radiation consists of high-energy electromagnetic radiation. It often accompanies α and β decay.

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