RI Mole Concept and Stoichiometry Notes 2025

-1- Raffles Institution Year 5 H2 Chemistry 2025 Supplementary Notes – The Mole Concept and Stoichiometry Content • Relative masses of atoms and molecules • The mole, the Avogadro constant • The calculation of empirical and molecular formulae • Reacting masses and volumes (of solutions and gases) Learning Outcomes [The term relative formula mass or Mr will be used for ionic compounds] Candidates should be able to: (a) define the terms relative atomic, isotopic, molecular and formula masses (b) define the term mole in terms of the Avogadro constant (c) calculate the relative atomic mass of an element given the relative abundances of its isotopes (d) define the terms empirical and molecular formula (e) calculate empirical and molecular formulae, using combustion data or composition by mass (f) write and/or construct balanced equations (g) 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] (h) deduce stoichiometric relationships from above calculations such as those in (g) Lecture Outline References 1 Atoms and Sub-atomic Particles 1 Cambridge International AS and A Level Chemistry (by Peter Cann and Peter Hughes) 2 Cambridge International AS and A Level Chemistry Coursebook with CD-ROM (by Lawrie Ryan and Roger Norris) 3 A Level Chemistry 4ed (by E. N. Ramsden) 4 Chemistry in Context 6ed (by Graham Hill and John Holman) 5 Chemistry – The Molecular Nature of Mater and Change (by Silberberg) 6 https://www.chemguide.co.uk 2 Relative Masses 3 The Mole and Related Concepts 4 Empirical and Molecular Formulae 5 Stoichiometry 6 Reacting Volumes of Gases 7 Concentration of a Solution 8 Acid-Base Titrations

-2- 1 Atoms and Sub-atomic Particles 1.1 The sub-atomic particles • An atom is the smallest part of an element which can ever exist, whereas a molecule is the smallest part of an element or a compound which can exist alone under ordinary conditions. • A ll atoms are composed of three fundamental particles  protons, neutrons and electrons. Sub-atomic particle proton neutron electron Symbol 1 1p 1 0 n 0 1e− Relative mass 1 1 1 1840 Relative charge +1 0 –1 Location within the atom in the nucleus in the nucleus around the nucleus 1.2 Important terms and definitions Term Symbol Definition (a) proton number or atomic number Z • The proton number (or atomic number) of an element is the number of protons in the nucleus of an atom of that element. • The atomic number determines the identity of an atom. For example, every atom with an atomic number of 6 is a carbon atom; it contains 6 protons in its nucleus. (b) nucleon number or mass number A • The nucleon number (or mass number) of an element is the total number of protons and neutrons in the nucleus of an atom of that element. • N ote: Protons and neutrons are collectively known as nucleons because they are both found in the nucleus. (c) nuclide A Z X • A nuclide is any species of given mass number and atomic number. • Examples: • The nuclide of an element is represented by • Note: Total number of protons and neutrons = A Number of protons = Z Number of neutrons = A – Z Number of electrons (for uncharged species) = number of protons = Z

-3- 1.3 Isotopes • Isotopes of an element are atoms with the same proton number but different nucleon numbers (i.e. they have the same number of protons but different number of neutrons in the nucleus). • Isotopes have the same number of electrons ⇒ the same chemical properties Isotopes have different numbers of neutrons (i.e. different masses) ⇒ different physical properties • Most elements consist of mixtures of isotopes. The abundance of each isotope in the mixture is called its isotopic abundance (in terms of percentages or fractions). • Example 1: Isotopes of hydrogen Name Symbol Number of protons Number of neutrons Number of electrons Isotopic abundance in natural hydrogen protium (hydrogen) 1 1H or H 1 0 1 99.984% deuterium (heavy hydrogen) 2 1Hor D 1 1 1 0.015% tritium 3 1Hor T 1 2 1 very rare  1 part in 1017 (radioactive and unstable) • Example 2: Isotopes of chlorine Name Symbol Number of protons Number of neutrons Number of electrons Isotopic abundance in naturally occurring chlorine chlorine-35 35 17 Cl 17 18 17 75% chlorine-37 37 17 Cl 17 20 17 25%  Worked Example 1  (a) Which one of the following particles has more electrons than protons and more protons than neutrons? A D− B OH− C H3O+ D OD− (b) Consider the atoms and ions given in the table below. Atom / Ion Number of electrons Number of neutrons P 11 15 Q2– 11 17 R+ 10 15 S– 12 17 T+ 13 16 Which of the following is an isotope of P? A Q C S B R D T Solution (a) D− OH− H3O+ OD− (b) P Q2– R+ S– T+ No. of electrons 2 10 10 10 No. of electrons 11 11 10 12 13 No. of protons 1 9 11 9 No. of protons 11 9 11 11 14 No. of neutrons 1 8 8 9 No. of neutrons 15 17 15 17 16

-4- 2 Relative Masses 2.1 The carbon-12 scale • The masses of atoms are very small, from 10–24 to 10–22 grams. Chemists use a relative atomic mass scale to compare the masses of different atoms. • In 1961, the carbon-12 atom was adopted by the International Union of Pure and Applied Chemistry (IUPAC) as the reference standard for relative atomic masses. • On the carbon-12 scale, atoms of the isotope 12C are assigned a relative atomic mass of 12 and the relative masses of all other atoms are obtained by comparison with the mass of the carbon-12 atom. 2.2 Relative isotopic mass • The relative isotopic mass of a particular isotope of an element is defined as follows: Relative isotopic mass = mass of 1 atom of the isotope 1 12 x mass of 1 atom of carbon-12 Note: No units • Examples: Relative isotopic mass of 21Ne = 20.994 ≈ 21.0 Relative isotopic mass of 35Cl = 34.97 ≈ 35.0 • Why is the relative isotopic mass of an isotope very close to a whole number? Reason: On the carbon-12 scale, the relative masses of the proton and neutron are both very close to one and the electron has a negligible mass. It therefore follows that all relative isotopic masses will be very close to whole numbers which are essentially the corresponding nucleon numbers of the isotopes. • In calculations, the relative isotopic mass of an isotope is often approximated by the nucleon number of that isotope. The two are assumed to be identical in all but the most accurate work. 2.3 Relative atomic mass (Symbol: Ar) • The relative atomic mass (Ar) of an element is defined as follows: Relative atomic mass = (weighted) average mass of 1 atom of the element 1 12 x mass of 1 atom of carbon-12 Note: No units • The Ar values can be found in the Periodic Table given in the Data Booklet. Example: Ar of chlorine = 35.5 • T he relative atomic mass of an element may not be close to a whole number. Why is this so? Reason: This is because naturally occurring elements often consist of a mixture of isotopes and the relative atomic mass of an element takes into consideration the different isotopes and their relative abundances.

-5-  Worked Example 2   Calculate the relative atomic mass of chlorine from the given data. Isotope Relative isotopic mass Percentage abundance 35Cl 37Cl 34.97 36.95 75.53 24.47 Solution Ar of Cl = ( )( ) ( )( ) 75.53 34.97 24.47 36.95 75.53 24.47 + + = 35.5 (3 s.f.)  Worked Examp