Atoms and matter

Having read this chapter you will be able to:

• Appreciate the planetary and Bohr models of the atom.

• Define an element's atomic number and atomic mass.

• Recall the key differences between solids, liquids and gases.

• Understand the role of energy in changing states.

• Recognize the value of phase diagrams for showing state, triple point and critical point.

1.1 The atom

The word 'atom' originates from the Greek atomos meaning indivisible. In 1912, however, a New Zealand physicist, Ernest Rutherford, caused a sensation by revealing the atom is divisible. He had shown that the mass of the atom is concentrated in a tiny positively charged nucleus, surrounded by a tenuous cloud of negatively charged electrons. The new science of atomic physics was born which, for better or worse, heralded the beginning of the atomic age.

Rutherford's model of the atom

Rutherford created what is now the classic model of the atom, that of an 'atomic planetary model' with the electrons (planets) orbiting the nucleus (the sun). Planets are drawn to the sun due to gravitational force, but the attraction for the atom is due to the particles' electrical charges; the positively charged nucleus attracts the negatively charged electrons.

The nucleus of the atom contains nucleons, and is where virtually all of the atom's mass is held. There are two types of nucleons: protons and neutrons, and both have approximately the same mass, which is about 1840 times the mass of an electron. Protons are positively charged while neutrons have no charge. The number of protons defines the atomic number and may be thought of as the 'fingerprint' of an element because it is fixed for a specific element, e.g. hydrogen has atomic number = 1 and carbon has atomic number = 6. The atomic number determines the element's place in the periodic table.

The total number of nucleons is almost exactly equal to the atomic mass. Atomic mass is expressed in units of atomic mass units (not in units of actual mass). There are small differences between the atomic mass and the nucleon number depending on the element in question.

Atoms of the same element can have different numbers of neutrons in their nucleus, and these are known as isotopes of the element. For example, helium (atomic number = 2) has two isotopes: helium-3, and helium-4. Helium-4 is by far the most common isotope and has two protons and two neutrons in its nucleus (see Figure 1.1), so has an atomic mass of approximately 4. Helium-3, which is highly sought after for fusion research, has only one neutron so has an atomic mass of approximately 3. On earth, there are less than two atoms of helium-3 for every 10 000 of helium-4.

Some isotopes are stable while some are highly unstable and emit particles or radiation as they disintegrate. These isotopes are described as radioactive and are discussed in more detail in Chapter 26.

Units. Atomic mass number was originally standardized so that one atomic mass unit was equal to the mass of a proton (a hydrogen nucleus). This convention has now been changed so that an atomic mass unit is equal to 1/12th of the mass of a carbon-12 nucleus.

The Bohr model and energy levels

Just two years after publication of Rutherford's model of the atom, the Danish physicist Niels Bohr incorporated the idea of energy levels into a new atomic model. Bohr's model had strict rules for electrons: they could only exist in defined energy levels, so they could jump from one level to another but their energy levels were fixed. These energy levels are organized into 'shells' around the atom, an idea which forms a cornerstone of quantum physics and led to the development of the laser (see Chapter 24). Sometimes Bohr's model is called the Rutherford–Bohr model as Bohr essentially improved Rutherford's original model.

Chemical bonding

The attraction between atoms is known as a chemical bond, which allows the formation of chemical substances containing two or more atoms. Chemical bonds can be strong interatomic bonds such as covalent bonds or ionic bonds, or (usually) weaker intermolecular bonds such as dipole–dipole interactions or hydrogen bonding (see Section 2.4).

In covalent bonds, two atoms share one or more of their outer shell electrons. The negatively charged electrons occupy the space between the positively charged nuclei and are attracted to both nuclei simultaneously. The electrons can be thought to exist in a 'cloud' between the nuclei, because they are moving rapidly around an equilibrium position between the atoms. This attraction overcomes the repulsion which would otherwise exist between the two nuclei, so a strong bond is formed. Covalent bonds usually form between non-metallic atoms, for example, in organic compounds, as well as in diatomic gases and water molecules.

In an ionic bond, an outer electron is transferred from one atom to another. The electron is more tightly bound in the new atom so is able to exist at a lower energy level than in the donor atom. The result of the transfer is that the electron-accepting atom becomes a negatively charged ion (an anion), while the other becomes a positive ion (a cation), resulting in an electrostatic attraction between them. Ionic bonds usually occur between metallic atoms (forming cations) and nonmetals (forming anions). The cation and anion bond to form a metal salt, a well known example being sodium chloride, (Na+Cl-).

1.2 States of matter

Solids, liquids and gases

The way atoms interact with one another determines the properties of matter. Interatomic and intermolecular bonds both determine the bulk properties of a compound, including whether it exists as a solid, a liquid or a gas at a given temperature. Solids have rigid bonds between their molecules; liquids have looser bonds; gases have minimal bonds. Table 1.1 summarizes the microscopic...