Section 1 : Physical Chemistry
Amount of Substance
Relative Atomic Mass and Relative Molecular Mass
Relative atomic mass (Ar): the average mass of an atom of an element relative to 1/12 the mass of
an atom of carbon-12. (Unit: none)
The following all have the same numerical value:
• Relative molecular mass (Mr): the average mass of a molecule relative to 1/12 the mass of an atom
of carbon-12 (carbon was chosen as it is a solid at room temp., non-toxic, and easy to separate).
(Unit: none)
• Relative formula mass (RFM) may be used for ionic compounds but Mr is accepted for all
compounds. (Unit: none)
• Molar mass: the mass of 1 mole. (Unit: gmol-1)
The Mole and the Avogadro Constant
Avogadro constant: the no. of particles in one mole = 6.02 x 1023
∴ for solids and pure liquids (not solutions): mass/Ar = moles = no. of particles/Avogadro’s constant
Calculating Masses in Reactions
1) Calculate the no. of moles of the species with enough data to do this.
2) Use stoichometry of the equation to deduce the no. of moles of target species (the one we are being
asked about).
3) Convert moles of target species to the mass.
E.g. What mass of O2 would be needed to react with H2(g) to produce 9g of H2O?
2H2(g) + O2(g) → 2H2O(l)
• moles H2O = mass/Ar = 9/18 = 0.5 moles
• moles O2 = 0.5 moles of H2O = 0.25 moles
• mass O2 = moles x Ar = 0.25 x 32 = 8g
The Ideal Gas Equation
The ideal gas equation is: pV = nRT where…
• p = pressure (1atm = 101325Pa = 101325Nm-2) (Unit: Pa)
• V = volume (1m3 = 1,000dm3 = 1,000,000cm3) (Unit: m3)
• n = no. of moles (k = kilo = x1,000; M = mega = x1,000,000) (Unit: mol)
• R = gas constant = 8.31JK-1 mol-1
• T = temperature (K = °C + 273; °C = K - 273) (Unit: K)
Empirical and Molecular Formula
Empirical formula: the simplest whole no. ratio of atoms of each element in a compound.
Molecular formula: the actual no. of atoms of each element in a compound.
• ∴molecular formulae are simple whole no. multiples of empirical formulae.
Calculating Empirical Formula
1) Convert masses to moles: moles = mass/Ar
2) Divide each molar quantity by the smallest.
3) If one of the numbers are not close enough to round (not close enough to whole by 2 d.p.) multiply
it by a factor accordingly.
Amount of Substance
Relative Atomic Mass and Relative Molecular Mass
Relative atomic mass (Ar): the average mass of an atom of an element relative to 1/12 the mass of
an atom of carbon-12. (Unit: none)
The following all have the same numerical value:
• Relative molecular mass (Mr): the average mass of a molecule relative to 1/12 the mass of an atom
of carbon-12 (carbon was chosen as it is a solid at room temp., non-toxic, and easy to separate).
(Unit: none)
• Relative formula mass (RFM) may be used for ionic compounds but Mr is accepted for all
compounds. (Unit: none)
• Molar mass: the mass of 1 mole. (Unit: gmol-1)
The Mole and the Avogadro Constant
Avogadro constant: the no. of particles in one mole = 6.02 x 1023
∴ for solids and pure liquids (not solutions): mass/Ar = moles = no. of particles/Avogadro’s constant
Calculating Masses in Reactions
1) Calculate the no. of moles of the species with enough data to do this.
2) Use stoichometry of the equation to deduce the no. of moles of target species (the one we are being
asked about).
3) Convert moles of target species to the mass.
E.g. What mass of O2 would be needed to react with H2(g) to produce 9g of H2O?
2H2(g) + O2(g) → 2H2O(l)
• moles H2O = mass/Ar = 9/18 = 0.5 moles
• moles O2 = 0.5 moles of H2O = 0.25 moles
• mass O2 = moles x Ar = 0.25 x 32 = 8g
The Ideal Gas Equation
The ideal gas equation is: pV = nRT where…
• p = pressure (1atm = 101325Pa = 101325Nm-2) (Unit: Pa)
• V = volume (1m3 = 1,000dm3 = 1,000,000cm3) (Unit: m3)
• n = no. of moles (k = kilo = x1,000; M = mega = x1,000,000) (Unit: mol)
• R = gas constant = 8.31JK-1 mol-1
• T = temperature (K = °C + 273; °C = K - 273) (Unit: K)
Empirical and Molecular Formula
Empirical formula: the simplest whole no. ratio of atoms of each element in a compound.
Molecular formula: the actual no. of atoms of each element in a compound.
• ∴molecular formulae are simple whole no. multiples of empirical formulae.
Calculating Empirical Formula
1) Convert masses to moles: moles = mass/Ar
2) Divide each molar quantity by the smallest.
3) If one of the numbers are not close enough to round (not close enough to whole by 2 d.p.) multiply
it by a factor accordingly.
