STPM Chemistry Form 6 Notes – Terminology
and Concepts: Matter
Matter – anything that occupies space and has mass.
Fundamental Particles of Atoms (Historical Point of View)
John Dalton (1808) – atomic theory
1. Atoms – small indivisible particles.
2. Atoms – neither created nor destroyed.
3. Atoms – chemical reactions result from combination / separation of atoms.
J. J. Thomson (1897)
1. Electrons – negatively-charged particles.
2. Atoms – positively-charged sphere.
Ernest Rutherford (1911)
1. Atoms – consists of a positively-charged nucleus with a cloud of electrons surrounding
nucleus.
2. Protons – positively-charged particles.
Niels Bohr (1913)
1. Electrons – surrounding the nucleus (orbit).
James Cadwick (1932)
1. Neutrons – electrically neutral subatomic particles.
2. Neutrons – mass almost the same with a proton.
3. Nucleus of an atom – consists of protons and neutrons.
Modern Atomic Model
1. Nucleus of an atom – consists of protons and neutrons.
2. Electrons – moving around the nucleus (orbits / electron shells/ quantum shells)
Atoms
Atom – smallest particle of an element.
Relative atomic mass (Ar) - (an element) average mass of one atom of the element
relative to 1/12 times the mass of one atom of carbon-12.
= (average mass of one atom of the element) / (1/12 x mass of one atom of C-12)
Or
= 12 x [(average mass of one atom of the element) / (mass of one atom of C-12)]
Cations – positively-charge ions.
Example: H+, K+, NH4+ and Mg2+
Anions – negatively-charge ions.
Example: Br-, OH-, O2- and S2O32Molecule – a group of two or more atoms.
Relative molecular mass (Mr) – (an element or compound) average mass of one
molecule of the substance relative to 1/12 times the mass of one atom of carbon-12.
= (average mass of one molecule of substance) / (1/12 x mass of one atom of C-12)
Or
= 12 x [(average mass of one molecule of substance) / (mass of one atom of C-12)]
Proton number / Atomic number / Number of protons (Z)

Number of protons in the nucleus of an atom.

Number of electrons (neutral atom).
Nucleon number / Mass number / Number of nucleon (A)

total number of protons and neutrons in the nucleus of an atom.
A=Z+N
N = number of neutrons
Isotopes (of the same element)

atoms having the same proton number but different nucleon number.

same number of protons, number of electrons, electronic configuration and chemical
properties.

different nucleon number, relative mass, density and rate of diffusion.
Relative isotopic mass – the ratio of the mass of one atom of the isotope relative to 1/12
times the mass of one atom of carbon-12 isotope.
= (mass of one atom of the isotope) / (1/12 x mass of one atom of C-12)
Or
= 12 x [(mass of one atom of the isotope) / (mass of one atom of C-12)]
Mass spectrometry
i. Vaporisation chamber – sample is vaporised (produce gaseous atoms or molecules).
ii. Ionisation chamber – vapour is bombarded with a stream of high-energy electrons
to form positive ions. X(g) + e –> X+(g) + 2e. (produce positive ions)
iii. Acceleration chamber – positive ions are attracted towards the high negative potential
plated that accelerates the positive ions to a high and constant velocity. (accelerate the
positive ions).
iv. Magnetic Field – accelerated positive ions are deflected into a circular path according to
the m/e ratio. (separate positive ions of different m/e ratio)
v. Ion detector – positive ions with different m/e ratios will be deflected to the ion detector
that can be recorded on a moving chart. (detect the number and m/e ratio of the positive
ions)
vi. Recorder – a flow of current which is amplified and recorded as peaks. (plot the mass
spectrum of the sample)
Important note:

A lighter ion will deflect more than a heavier ion (the same charge)
Example: 35Cl+ will deflect more than 37Cl+

An ion with a higher charge will deflect more than an ion with a lower charge (the same
mass)
Example: 35Cl2+ will deflect more than 35Cl+
Isotopic abundance = fractional abundance = percentage abundance
One mole – the quantity of a substance that contains the same number of particles (atoms,
ions or molecules) as the number of atoms in exactly 12 grams of carbon-12 isotope.
Avogadro constant, L or NA – number of particles (atoms, ions or molecules) present in
a mole of substance (elements or compounds)
= 6.02 x 1023 (unit is mol-1)
Number of moles = number of atoms or molecules / Avogadro constant (mol-1)
Number of particles in a sample = number of moles x Avogadro constant (mol-1)
Mass (g) = number of moles (n) x M (Ar or Mr)
Number or moles (n) = mass (g) / molar mass (g mol-1)
Mass (g) = number of moles x molar mass (g mol-1)
Number of moles = volume of gas (dm3) / 22.4 dm3 at s.t.p. (0˚C and 1 atm or 101 kPa)
Number of moles = volume of gas (dm3) / 24 dm3 at r.t.p. (25˚C and 1 atm or 101 kPa)
Volume of gas (dm3) = number of moles x / 22.4 dm3 at s.t.p.
Volume of gas (dm3) = number of moles x / 24 dm3 at r.t.p.
Number of moles of solute = MV / 1000
(M = concentration in mol dm-3)
(V = volume in cm3)
Concentration of a solution (g dm-3) = mass of solute (g) / volume of the solution
(dm3)
Concentration of a solution (mol dm-3) = number or moles of solute (mol) / volume of
the solution (dm3)
MaVa / MbVb = a/b
M1V1 = M2V2