Atomic Structure & Isotopes (OCR A Level Chemistry A): Revision Note

Subatomic Structure of Atoms & Ions

• All matter is composed of atoms, which are the smallest parts of an element that can take place in chemical reactions

• Atoms are mostly made up of empty space

• The mass of an atom is concentrated in the nucleus, because the nucleus contains the heaviest subatomic particles (the neutral neutrons and positive protons)

  • The mass of the electron is negligible

Relative mass & charge of subatomic particles table

• The nucleus is also positively charged due to the protons

• Negatively charged electrons orbit the nucleus of the atom, contributing very little to its overall mass, but creating a ‘cloud’ of negative charge

• The electrostatic attraction between the positive nucleus and negatively charged electrons orbiting around it is what holds an atom together

The mass of the atom is concentrated in the positively charged nucleus which is attracted to the negatively charged electrons orbiting around it

• An atom is neutral and has no overall charge

• Ions on the other hand are formed when atoms either gain or lose electrons, causing them to become charged

• The number of subatomic particles in atoms and ions can be determined given their atomic (proton) number, mass (nucleon) number and charge

Protons

• The atomic number of an atom and ion determines which element it is

• Therefore, all atoms and ions of the same element have the same number of protons (atomic number) in the nucleus

  • E.g. lithium has an atomic number of 3 (three protons) whereas beryllium has atomic number of 4 (4 protons)

• The number of protons equals the atomic (proton) number

• The number of protons of an unknown element can be calculated by using its mass number and number of neutrons:

Mass number = number of protons + number of neutrons

Number of protons = mass number - number of neutrons

Electrons

• An atom is neutral and therefore has the same number of protons and electrons

• Ions have a different number of electrons to their atomic number depending on their charge

  • A positively charged ion, or cation,  has lost electrons and therefore has fewer electrons than protons

  • A negatively charged ion, or anion, has gained electrons and therefore has more electrons than protons

Neutrons

• The mass and atomic numbers can be used to find the number of neutrons in ions and atoms:

Number of neutrons = mass number (A) - number of protons (Z)

Subatomic Structure of Isotopes

• Isotopes are atoms of the same element that contain the same number of protons and electrons but a different number of neutrons

• The symbol for an isotope is the chemical symbol (or word) followed by a dash and then the mass number

  • E.g. carbon-12 and carbon-14 are isotopes of carbon containing 6 and 8 neutrons respectively

The atomic structure and symbols of the three isotopes of hydrogen

• Isotopes have the same chemical properties but different physical properties

Chemical properties

• Isotopes of the same element display the same chemical characteristics

• This is because they have the same number of electrons in their outer shells

• Electrons take part in chemical reactions and therefore determine the chemistry of an atom

Physical properties

• The only difference between isotopes is the number of neutrons

• Since these are neutral subatomic particles, they only add mass to the atom

• As a result of this, isotopes have different physical properties such as small differences in their mass and density

Relative Mass

• The relative mass of an atom uses the carbon-12 isotope as the international standard 

• One atom of carbon-12 has an accepted mass of 1.992646538 x 10^-26 kg

• It is not realistic to work with this value so the mass of a carbon-12 atom is fixed as exactly 12 atomic mass units / 12υ

• The standard mass for atomic mass is 1υ

  • Therefore, the standard mass for comparison is the mass of  of a carbon-12 atom

Relative isotopic mass

• Relative isotopic mass is defined as the mass of an isotope relative to  of a carbon-12 atom

• For A Level Chemistry it is common to work with mass values rounded to one decimal place, for example:

• The accurate relative isotopic mass of nitrogen is 14.00307401 but this is rounded to 14.0

• The accurate relative isotopic mass of oxygen is 15.99491464 but this is rounded to 16.0

Relative atomic mass

• Most elements on the Periodic Table represent a mixture of different isotopes, which is shown as their relative atomic mass (A_r) 

• The relative atomic mass is the weighted mean / average mass of an atom relative to of the mass of a carbon-12 atom