Chapter 1 “Moles and Equations”, Candidates should be able to:
(a) define the terms relative atomic, isotopic, molecular and formula
masses, based on the 12C scale
Ar , relative atomic mass – the weighted average mass of naturally occurring atoms of an element, taking into account the proportions of naturally occurring isotopes, measured on a scale where an atom of carbon-12 has exactly 12 units.
relative isotopic mass – the mass of a particular isotope of an element on a scale where an atom of the carbon-12 isotope has a mass of exactly 12 units.
Mr , relative molecular mass – the mass of a molecule measured on a scale on which an atom of the carbon-12 isotope has a mass of exactly 12 units.
relative formula mass – the mass of one formula unit of a compound measured on a scale on which an atom of the carbon-12 isotope has a mass of exactly 12 units.
Isotope – atoms of an element with the same number of protons but different numbers of neutrons and therefore mass.
Nucleon number – the total number of neutrons plus protons in an atom, see top left corner of symbol.
(b) define the term mole in terms of the Avogadro constant
Avogadro constant – the number of atoms (or ions, molecules or electrons) in a mole of atoms (or ions, molecules): its numerical value is 6.02 x 10 23
Mol, mole – the unit of amount of substance. Is is the amount of substance that has the same number of particles (atoms, ions, or molecules) as there are atoms in exactly 12g of the carbon-12 isotope.
Given grams of compound X (1 mole)/(grams from periodic table) = mole
Given number of moles X ( grams from periodic table)/(1 mol) = grams
(c) *analyse mass spectra in terms of isotopic abundances and molecular fragments [knowledge of the working of the mass spectrometer is not required]
Mass spectrometer – an instrument for finding the relevant isotopic abundance of elements and to help identify unknown organic compounds
(d) calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum
Ar = 1. Multiply each mass by percent found
2. add all together
3. divide by 100
(e) define the terms empirical and molecular formulae
Empirical formula – the formula that tells us the simplest ratio of the different atoms present in a molecule.pg 245 Modern Chemistry
1. Determine moles of each element
2. Divide smallest number into mole of each
3. Round to the nearest whole number {too far to round is (x.1 ~ x.9)}
Molecular formula – the formula that tells us the actual numbers of each type of atom in a molecule. Pg. 248 Modern Chemistry
1. X = Molecular Formula Mass / Empirical Formula Mass
2. Molecular Formula = X (Empirical Formula)
(f) calculate empirical and molecular formulae, using combustion data or composition by mass
(g) *write and/or construct balanced equations. Pages 226, 228, 230 Modern Chemistry
(h) perform calculations, including use of the mole concept, involving:
(i) reacting masses (from formulae and equations)
(ii) volumes of gases (e.g. in the burning of hydrocarbons)
(iii) volumes and concentrations of solutions
When performing calculations, candidates’ answers should reflect the number of significant figures given or asked for in the question.
When rounding up or down, candidates should ensure that significant figures are neither lost unnecessarily nor used beyond what is justified (see also Practical Assessment, Paper 3, Display of calculation and reasoning)
(i) deduce stoichiometric relationships from calculations such as those in (h)
stoichiometry – shows the relationship of moles of the reactants and products. If one or more reactants is in excess, then you only need to know the limiting reactant. Page 308 Modern Chemistry
(Grams X) * (Moles X)/(Grams X) * (Moles Y)/(Moles X) * (Grams Y)/(Moles Y) = Grams Y
Special Notes:
1L = 1000 mL = 1000 cm3 = 1 dm3. You should notice that it is written dm-3 to avoid having dm3 in the denominator.
% composition by mass = (Atomic mass*number of moles of element)/(molar mass of compound)* 100
Molarity = or mol dm-3
Titrations:
Calculating solution concentration:
1. Calculate moles of acid
2. Calculate moles of base
3. Concentration = = (number of moles of solute (mol)) / (volume of solution (dm3))
Stoichiometry by titration:
1. Find moles of each
2. Determine the ratio
3. Write the equation
Gas Volumes: At room temperature and pressure one mole of gas occupies 24.0 dm3
Total volume = (RTP)24.0 dm3 * moles
Moles = (Total volume)/RTP
Alkane Formula = Cn H2n+2, only single bonds
methane, CH4 - one carbon and four hydrogen
ethane, C2H6 - two carbon and six hydrogen
propane, C3H8 - three carbon and 8 hydrogen
butane, C4H10 - four carbon and 10 hydrogen
pentane, C5H12 - five carbon and 12 hydrogen
hexane, C6H14 - six carbon and 14 hydrogen
Alkene = Cn H2n, double bonds present
Alkyne = CnH2n-2 , triple bonds present
(a) define the terms relative atomic, isotopic, molecular and formula
masses, based on the 12C scale
Ar , relative atomic mass – the weighted average mass of naturally occurring atoms of an element, taking into account the proportions of naturally occurring isotopes, measured on a scale where an atom of carbon-12 has exactly 12 units.
relative isotopic mass – the mass of a particular isotope of an element on a scale where an atom of the carbon-12 isotope has a mass of exactly 12 units.
Mr , relative molecular mass – the mass of a molecule measured on a scale on which an atom of the carbon-12 isotope has a mass of exactly 12 units.
relative formula mass – the mass of one formula unit of a compound measured on a scale on which an atom of the carbon-12 isotope has a mass of exactly 12 units.
Isotope – atoms of an element with the same number of protons but different numbers of neutrons and therefore mass.
Nucleon number – the total number of neutrons plus protons in an atom, see top left corner of symbol.
(b) define the term mole in terms of the Avogadro constant
Avogadro constant – the number of atoms (or ions, molecules or electrons) in a mole of atoms (or ions, molecules): its numerical value is 6.02 x 10 23
Mol, mole – the unit of amount of substance. Is is the amount of substance that has the same number of particles (atoms, ions, or molecules) as there are atoms in exactly 12g of the carbon-12 isotope.
Given grams of compound X (1 mole)/(grams from periodic table) = mole
Given number of moles X ( grams from periodic table)/(1 mol) = grams
(c) *analyse mass spectra in terms of isotopic abundances and molecular fragments [knowledge of the working of the mass spectrometer is not required]
Mass spectrometer – an instrument for finding the relevant isotopic abundance of elements and to help identify unknown organic compounds
(d) calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum
Ar = 1. Multiply each mass by percent found
2. add all together
3. divide by 100
(e) define the terms empirical and molecular formulae
Empirical formula – the formula that tells us the simplest ratio of the different atoms present in a molecule.pg 245 Modern Chemistry
1. Determine moles of each element
2. Divide smallest number into mole of each
3. Round to the nearest whole number {too far to round is (x.1 ~ x.9)}
Molecular formula – the formula that tells us the actual numbers of each type of atom in a molecule. Pg. 248 Modern Chemistry
1. X = Molecular Formula Mass / Empirical Formula Mass
2. Molecular Formula = X (Empirical Formula)
(f) calculate empirical and molecular formulae, using combustion data or composition by mass
(g) *write and/or construct balanced equations. Pages 226, 228, 230 Modern Chemistry
(h) perform calculations, including use of the mole concept, involving:
(i) reacting masses (from formulae and equations)
(ii) volumes of gases (e.g. in the burning of hydrocarbons)
(iii) volumes and concentrations of solutions
When performing calculations, candidates’ answers should reflect the number of significant figures given or asked for in the question.
When rounding up or down, candidates should ensure that significant figures are neither lost unnecessarily nor used beyond what is justified (see also Practical Assessment, Paper 3, Display of calculation and reasoning)
(i) deduce stoichiometric relationships from calculations such as those in (h)
stoichiometry – shows the relationship of moles of the reactants and products. If one or more reactants is in excess, then you only need to know the limiting reactant. Page 308 Modern Chemistry
(Grams X) * (Moles X)/(Grams X) * (Moles Y)/(Moles X) * (Grams Y)/(Moles Y) = Grams Y
Special Notes:
1L = 1000 mL = 1000 cm3 = 1 dm3. You should notice that it is written dm-3 to avoid having dm3 in the denominator.
% composition by mass = (Atomic mass*number of moles of element)/(molar mass of compound)* 100
Molarity = or mol dm-3
Titrations:
Calculating solution concentration:
1. Calculate moles of acid
2. Calculate moles of base
3. Concentration = = (number of moles of solute (mol)) / (volume of solution (dm3))
Stoichiometry by titration:
1. Find moles of each
2. Determine the ratio
3. Write the equation
Gas Volumes: At room temperature and pressure one mole of gas occupies 24.0 dm3
Total volume = (RTP)24.0 dm3 * moles
Moles = (Total volume)/RTP
Alkane Formula = Cn H2n+2, only single bonds
methane, CH4 - one carbon and four hydrogen
ethane, C2H6 - two carbon and six hydrogen
propane, C3H8 - three carbon and 8 hydrogen
butane, C4H10 - four carbon and 10 hydrogen
pentane, C5H12 - five carbon and 12 hydrogen
hexane, C6H14 - six carbon and 14 hydrogen
Alkene = Cn H2n, double bonds present
Alkyne = CnH2n-2 , triple bonds present