Atomic Structure

1.1 Describe how the Dalton model of an atom has changed over time because of the discovery of subatomic particles
John Dalton (1803)
  • Scientists knew that elements combined with each other in specific proportions to form compounds.
  • Dalton claimed that the reason for this was because elements are made of atoms.
  • He published his own three-part atomic theory:
    • All substances are made of atoms. Atoms are small particles that cannot be created, divided, or destroyed.
    • Atoms of the same element are exactly alike, and atoms of different elements are different.
    • Atoms join with other atoms to make new substances.
  • Much of Dalton’s theory was correct, but some of it was later proven incorrect and revised as scientists learned more about atoms.
J.J. Thomson (1897)
  • Thomson used a cathode-ray tube to conduct an experiment which showed that there are small particles inside atoms.
  • This discovery identified an error in Dalton’s atomic theory. Atoms can be divided into smaller parts.
  • Because the beam moved away from the negatively charged plate and toward the positively charged plate, Thomson knew that the particles must have a negative charge.
  • He called these particles corpuscles. We now call these particles electrons.
  • Electrons – The negatively charged particles found in all atoms.
  • Thomson changed the atomic theory to include the presence of electrons. He knew there must be positive charges present to balance the negative charges of the electrons, but he didn’t know where.
  • Thomson proposed a model of an atom called the “plum-pudding” model, in which negative electrons are scattered throughout soft blobs of positively charged material.
Ernest Rutherford (1909)
  • Rutherford conducted an experiment in which he shot a beam of positively charged particles into a sheet of gold foil.
  • Rutherford predicted that if atoms were soft, as the plum-pudding model suggested, the particles would pass through the gold and continue in a straight line.
  • Most of the particles did continue in a straight line. However some of the particles were deflected to the sides a bit, and a few bounced straight back.
  • Rutherford realized that the plum- pudding model did not explain his observations. He changed the atomic theory and developed a new model of the atom.
  • Rutherford’s model says that most of the atom’s mass is found in a region in the center called the nucleus.
  • Nucleus – The tiny, extremely dense, positively charged region in the center of an atom.
  • Rutherford calculated that the nucleus was 100,000 times smaller than the diameter of the atom.
  • In Rutherford’s model the atom is mostly empty space, and the electrons travel in random paths around the nucleus.
1.2 Describe the structure of an atom as a nucleus containing protons and neutrons, surrounded by electrons in shells
1.3 Recall the relative charge and relative mass of protons, neutrons and electrons
1.4 Explain why atoms contain equal number of protons and electrons
  • Atoms are neutral, therefore amount of protons = amount of electrons, so that the charges cancel
1.5 Describe the nucleus of an atom as a very small compared to the overall size of the atom
1.6 Recall that most of the mass of an atom is concentrated in the nucleus
 1.7 Recall the meaning of the term mass number of an atom
  • Mass (nucleon) Number = number of protons + neutrons
 1.8 Describe atoms of a given element as having the same number of protons in the nucleus and that this number is unique to that element
 1.9 Describe isotopes as different atoms of the same element containing the same number of protons but different numbers of neutrons in their nuclei
 1.10 Calculate the numbers of protons, neutrons and electrons in atoms given the atomic number and mass number
  • Atomic (proton) Number = number of protons (= number of electrons if it’s an atom, because atoms are neutral)
 1.11 Explain how the existence of isotopes results in relative atomic masses of some elements not being whole numbers

Relative Atomic mass

  • Atoms are very small, having a radius of about 0.1nm (1 x 10^-10m).
  • The radius of a nucleus is less than 1/10 000 of that of the atom (about 1 x 10^-14 m).
  • Different types of atoms have different masses. This mass is too small to measure using a conventional scale, therefore we compare their masses to each other. A carbon atom having a mass number 12, i.e. (12C) is taken as standard for this comparison and its relative atomic mass is 12.

NB: Atoms of the same element can have different numbers of neutrons; these atoms are called isotopes of that element.

  • The relative atomic mass of an element is an average value that takes account of the abundance of the isotopes of the element.
  • It is written as Ar or A.M.
  • Some of the elements exist in nature as a mixture of their isotopes in specific proportions. The R.A.M of such elements is the average mass of the different proportions of each isotope in the mixture.
  • A.M. of some elements:
hydrogen 1
oxygen 16
copper 63.5
iron 55.8
  • It’s important to note that atomic number and mass number are always whole numbers because they are based on the number of sub-atomic particles, while the R.A.M. can have fractions because it is the average mass of different isotopes.
1.12 (HT only) Calculate the relative atomic mass of an element from the relative masses and abundances of its isotopes

A sample of chlorine gas is a mixture of 2 isotopes, chlorine-35 and chlorine-37. These isotopes occur in specific proportions in the sample i.e. 75% chlorine-35 and 25% chlorine-37. Calculate the R.A.M. of chlorine in the sample.

The average mass, or R.A.M. of chlorine can be calculated using the following equation:

R.A.M. =   (mass of isotope-A  x  % of isotope-A) + (mass of isotope-B  x  % of isotope-B)
= (35 x 75) + (37 x 25)
= 3550
R.A.M. = 35.5