RI 2022 Mole Concept and Stoichiometry

Raffles Institution Year 5 H2 Chemistry 2022 Lecture Notes 1 a - 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-m~sses 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 2 Relative Masses Chemistry (by Peter Cann and Peter Hughes) 2 Cambridge International AS and A Level 3 The Mole and Related Concepts Chemistry Coursebook with CD-ROM (by 4 Empirical and Molecular Formulae Lawrie Ryan and Roger Norris) 5 Stoichiometry 3 A Level Chemistry 4ed (by E. N. Ramsden) 4 Chemistry in Context 6ed (by Graham Hill and 6 Reacting Volumes of Gases John Holman) 7 Concentration of a Solution 5 Chemistry - The Molecular Nature of Mater 8 Acid-Base Titrations and Change (by Silberberg) 6 http:///www.chemguide.co.uk -1-

1 I Atoms and Sub-atomic Particles I _ _,___ ----__:,:_:_:_::.:..._ ________ ..__.J 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. • All atoms are composed of three fundamental particles - protons, neutrons and electrons. 1.2 (a} (b} (c} Sub-atomic particle proton neutron electron Symbol ;p 1n oe 0 - 1 1 Relative mass 1 1 -1840 Relative charge +1 0 -1 Location within the atom in the nucleus in the nucleus around the nucleus Important terms and definitions Term proton number or atomic number nucleon number or mass number nuclide Symbol z A Definition • 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 element. For example, every atom with an atomic number of 6 is a carbon atom; it contains 6 protons in its nucleus. • 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. • Note: Protons and neutrons are collectively known as nucleons because they are both found in the nucleus. • A nuclide is any species of given mass number and atomic number. • Examples: 1H 1 9 4 Be 12 C 6 16 0 8 • The nuclide of an element is represented by nucleon number (or mass number) -~~-; A X ~z ~:1--- b I - sym o of the element proton number ---1➔ (or atomic number) • Note: Total number of protons and neutrons Number of protons Number of neutrons Number of electrons (for uncharged species) -2- = = = = A z A-Z number of protons = Z

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 Number of Number of Isotopic abundance in protons neutrons electrons natural hydrogen protium ;H or H 1 0 1 99.984% {hydrogen) deuterium ~HorD 1 1 1 0.015% (heavy hydrogen) tritium ~HorT 1 2 1 very rare -1 part in 1017 (radioactive and unstable) • Example 2: Isotopes of chlorine Name Symbol Number of Number of Number of Isotopic abundance in protons neutrons electrons naturally occurring chlorine chlorine-35 ~Cl 17 18 17 75% chlorine-37 31Cl 17 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 o- @ OH- C H3Q+ D oo- '~, l ~ 0 zo I ~ f (b) Consider the atoms and ions given in the table below. Atom/ Ion Solution (a) p a2- R+ s- r No. of electrons No. of protons No. of neutrons Number of electrons 11 11 10 12 13 o- OH- H3Q+ 2 10 (0 1 9 I I 1 8 ~ Number of neutrons Which of the following is an 15 isotope of P? 17 15 A Q @ s 17 B R D T 16 oo- (b) p Q2- R+ s- 10 No. of electrons 11 11 10 12 9 No. of protons t( q \ ( l ( 9 No. of neutrons 15 17 15 17 -3- p 13 I~ 16

2 I Relative Masses I 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: R I . . . mass of 1 atom of the isotope e atIve IsotopIc mass = 1 12 x mass of 1 atom of carbon-12 • Examples: Relative isotopic mass of 21 Ne = 20.994:::: 21.0 Relative isotopic mass of 35C/ = 34.97 :::: 35.0 • Why is the relative isotopic mass of an isotope very close to a whole number? Note: No units 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: (weighted) average mass of 1 atom of the element Relative atomic mass = 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 13 14 15 16 17 18 Al Si p s Cl Ar aluminium silicon phosphorus sulfur chlorine argon 27.0 28.1 31 .0 32.1 35.5 39.9 • The 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. -4-

---11■• Worked Example 2-/// Calculate the relative atomic mass of chlorine from the given data. Isotope Relative isotopic mass Percentage abundance 35C/ 34.97 75.53 31Cl 36.95 24.47 Solution (75.53)(34.97) + (24.47)(36.95) _ Ar of Cl= ~--'---'-