1 st Stage

Changing ideas about atoms
The idea of atoms has changed hugely over the years. At the moment, scientists believe atoms are very small, have a very small mass and are made of protons, electrons and neutrons. Our current theories were developed by imagination, evidence and advances in technology, with each new idea being built on the ideas of earlier scientists.
Explanations about atoms began about 400 BC, when the Greek philosopher Democritus described materials as being made of small particles. He called these particles ‘atoms’.
However, he had no evidence. It was just an idea. Little more was suggested for more than
2000 years, but in 1803 the British scientist John Dalton used his observations to describe the atom in more detail. His model described an atom as a ‘billiard ball’.
John Dalton published his ideas about atoms in 1803. He thought that all matter was made of tiny particles called atoms, which he imagined as tiny spheres that could not be divided.
Atoms consist of a nucleus containing protons and neutrons, surrounded by electrons in shells. The number of subatomic particles in an atom can be calculated from the atom's atomic number and mass number.

In 1897, 94 years later , J. J. Thomson discovered the electron. Thomson developed the way that the atom was thought of by using a ‘plum pudding’ model to describe atoms. Negative electrons were thought to be embedded in a ball of positive charge, rather like the fruit (the electrons) are part of a pudding (the ball of positive charge).
Changing theories
Sometimes ideas can develop rapidly because of unexpected results.
In 1909 Geiger and Marsden had really surprising results in their experiment with gold leaf and alpha particles. These results led Geiger , Marsden and Rutherford to propose a new idea that an atom has a nucleus. In 1911, Rutherford suggested the atom had a positively charged nucleus and much of the atom was empty space. This was the nuclear model of the atom.
This evidence led Rutherford to suggest a new model for the atom, called the nuclear model . In the nuclear model:
 the mass of an atom is concentrated at its center, the nucleus
 the nucleus is positively charged
The alpha particle scattering experiment

In 1913, Niels Bohr used theoretical calculations that agreed with experimental evidence to adapt the nuclear model. He explained that the electrons orbited the nucleus in definite orbits at specific distances from the nucleus. He explained that a fixed amount of energy (a quantum of energy) is needed for an electron to move from one orbit to the next.
Electrons only exist in these orbits.
Rutherford and Geiger in their lab in Manchester, UK

Further experiments led to the idea that the nucleus contained small particles, called protons. Each proton has a small amount of positive charge.
In 1932 James Chadwick found evidence for the existence of particles in the nucleus with mass but no charge. These particles are called neutrons. This led to another development of the atomic model, which is still used today.
The nuclear model of the atom, showing electrons in shells

Electrons occupy the space around the nucleus in ‘shells’. The space between the nucleus and the electron shells is completely empty. The nucleus contains most of the mass of the atom and the electrons contribute very little. On the other hand, the radius of the atom, where the electrons are orbiting, is much larger than the radius of the nucleus in the center. When we are talking about these differences we are talking about small sizes.
Atoms are very small. A typical atomic radius is about 0.1 nm (1 × 10 -10 m). The radius of a nucleus is less than one ten-thousandth of the radius of an atom (about 1 × 10 -14 m).
 Nucleus and shells
An atom has a central nucleus . This is surrounded by electrons arranged in shells.
The nucleus is tiny compared to the atom as a whole:
• the radius of an atom is about 0.1 nm (1 × 10 -10 m)
• the radius of a nucleus (1 × 10 -14 m) is less than of the radius of an atom
The structure of a carbon atom, not drawn to scale
For comparison, the radius of a typical bacterium is 1 × 10 -6 m and the radius of a human hair is about 1 × 10 -4 m.

Subatomic particles
The nuclei of all atoms contain subatomic particles called protons . The nuclei of most atoms also contain neutrons .
The masses of subatomic particles are very tiny. Instead of writing their actual masses in kilograms, we often use their relative masses . The relative mass of a proton is 1, and a particle with a relative mass smaller than 1 has less mass.
Subatomic particle
Proton
Neutron
Electron
Relative mass
1
1
Very small
Relative charge
+1
0
-1
The mass of an electron is very small compared to a proton or a neutron. Since the nucleus contains protons and neutrons, most of the mass of an atom is concentrated in its nucleus.
Protons and electrons have electrical charges that are equal and opposite. Remember that Protons are Positive, and Neutrons are Neutral.

The number of protons in an atom of an element is its atomic number . Remember that:
 all atoms of a given element have the same number of protons
 atoms of different elements have different numbers of protons
An atom contains equal numbers of protons and electrons . Since protons and electrons have equal and opposite charges, this means that atoms are have no overall electrical charge.
For example, the atomic number of sodium is 11. Every sodium atom has 11 protons and 11 electrons. It has 11 positive charges and 11 negative charges.
The mass number of an atom is its total number of protons and neutrons.
Atoms of different elements usually have different mass numbers , but they can be the same. For example, the mass number of argon atoms and calcium atoms can both be 40.

superscript subscript
X
= the chemical symbol for the element.
Z
= atomic number z
A
X
A
= Mass Number [the whole number closest to the accurate relative mass
(equal to the sum of the numbers of protons and neutrons in the nucleus)]
= the number of protons in the nucleus.
A-Z
= number of neutrons in the nucleus
The symbol for an atom can be written to show its mass number at the top, and its atomic number at the bottom.
To calculate the numbers of subatomic particles in an atom, use its atomic number and mass number:
• number of protons = atomic number
• number of electrons = atomic number
• number of neutrons = mass number - atomic number
Question / The atomic number of a sodium atom is 11 and its mass number is 23.
Calculate the number of protons, neutrons and electrons it contains.

Atoms of the same element must have the same number of protons , but they can have different numbers of neutrons . Atoms of the same element with different numbers of neutrons are called isotopes . Isotopes of an element have:
• the same atomic number
• different mass numbers
Three isotopes of hydrogen
All hydrogen atoms contain one proton (and one electron ), but they can contain different numbers of neutrons. Hydrogen-1 is the most abundant (most common) isotope of hydrogen.
Isotopes of any particular element have exactly the same chemical properties , but their physical properties vary slightly because they are dependent upon the atomic mass.
A minority of elements, such as the fluorine atom, are mono-isotopic in that their nuclei are unique. Examples of isotopy are shown in the Table.

An isotope is named after the element and the mass number of its atoms. For example, carbon-12 is an isotope of carbon with a mass number of 12.
All three isotopes of hydrogen have identical chemical properties . This is because the number of electrons determines chemical properties, and all three isotopes have one electron in their atoms.
The relative atomic mass of an element is a weighted average of the masses of the atoms of the isotopes. It takes account of the abundance of each of the isotopes of the element.
Relative atomic masses can be found in the periodic table. They have the symbol A r
.
Take care not to confuse mass numbers and relative atomic masses:
• mass numbers are always whole numbers (protons or neutrons cannot be split into parts)
• relative atomic masses are often rounded to the nearest whole number, but are actually not whole numbers
For example, the relative atomic mass of chlorine is 35.5 rather than a whole number. This is because chlorine contains two different isotopes, chlorine-35 and chlorine-37.

The relative atomic mass of an element (RAM) is a weighted average of the masses of the atoms of the isotopes
The three chemically important fundamental particles are used to construct atoms, molecules and infinite arrays which include a variety of crystalline substances and metals.
The interactions between constructional units are summarized in Table 1.3.

The nucleus of an atom contains the heavier subatomic particles – the protons and the neutrons. The electrons, the lightest of the sub-atomic particles, move around the nucleus at great distances from the nucleus relative to their size. They move very fast in electron energy levels very much as the planets orbit the Sun.
It is not possible to give the exact position of an electron in an energy level. However, we can state that electrons can only occupy certain, definite energy levels and that they cannot exist between them. Each of the electron energy levels can hold only a certain number of electrons.
●● First energy level holds up to two electrons.
●● Second energy level holds up to eight electrons.
●● Third energy level holds up to 18 electrons.

Q.1/ The table shows the mass numbers and abundances of naturally occurring copper isotopes.
Calculate the relative atomic mass of copper. Give your answer to 1 decimal place.
Mass number
63
65
Abundance
69%
31%
Q.2/
Q.3/
Q.4/

Q.5/
Q.6/
Q.7/
Q.8/
Q.9/

Q.9/
Q.10/
Q.11/
Q.12/
Q.13/

A.
B.
C.
A.
B.
C.
Q.14/ Choose the correct answer.
A.
B.
C.
A.
B.
C.
A.
B.
C.
A.
B.
C.

A.
B.
C.
A.
B.
C.
A.
B.
C.
A.
B.
C.
Q A Q A
1 6
2
3
7
8
4
5
9
10