AS-Level Chemistry Chapter 2 Atoms, Molecules Stoichiometry
Mass of Atoms and Molecules
Concept of relative mass
1) Relative mass is an indication of how heavy is an atom compared to another atom which is used as a standard model.
2) Relative mass is expressed in atomic mass unit (a.m.u).
3) C-12 was chosen to be the standard model because:
ii. it is a solid, easy to handle and easily available
4) C-12 was assigned a mass of exactly 12 a.m.u.. This is known as C-12 scale.
5) For example, an atom which is 3.5 times heavier than a C-12 atom would have a relative mass of (3.5 x 12) = 42 a.m.u.. That means, this atom is 42 times heavier than the mass of (1/12 x the mass of C-12 atom).
Relative isotopic mass
1) Relative isotopic mass is the mass of an isotope measured on a scale in which a carbon-12 atom has a mass of exactly 12 units.
Relative atomic mass, Ar
1) Relative atomic mass, Ar is the weighted average relative masses of all its isotopes measured on a scale in which a carbon-12 atom has a mass of exactly 12 units.
Relative atomic mass, Ar = Average mass of one atom of the element
Mass of one atom of carbon-12 = X = 12
Example: Ratio of Cl-35 to Cl-37 is 3:1. If you have 4 typical atoms of chlorine, total mass is (35 x 3) + (37 x 1) = 142. So, the average mass of the isotopes is 142/4 = 35.5.
This implies that 35.5 is the relative atomic mass of chlorine while 35 is the relative mass of Cl-35 and 37 is the relative mass of Cl-37.
Relative molecular mass, Mr
1) Relative molecular mass, Mr is the weighted average of the masses of the molecules measured on a scale in which a carbon-12 atom has a mass of exactly 12 units.
2) It should only be applied to substances which exist as molecules.
3) It is found by adding up all the relative atomic masses of all the atoms present in the molecule.
4) Examples:
ii. Mr (CHCl3) = 12 + 1 + 3(35.5) = 119.5
Relative formula mass, Mr
1) Relative formula mass, Mr is the weighted average of the masses of the formula units measured on a scale in which a carbon-12 atom has a mass of exactly 12 units.
2) It works for both ionic and covalent compounds.
3) Examples:
ii. Mr (CuSO4 • H2O) = 64 + 32+ 4(16) +5[2(1) + 16] = 249.5
2.2 Mass Spectrometer
What is mass spectrometer?
1) A mass spectrometer is used to determine:
b. relative abundance of isotopes
c. relative atomic mass
d. relative molecular mass
e. structural formula of compounds
Determination of relative atomic mass using mass spectrometer
1) Five steps:
ii. Ionization- gaseous atoms are bombarded with high energy electrons to form positive ions.
iii. Acceleration- the ions are accelerated so that they have the same kinetic energy.
iv. Deflection- ions are deflected by a magnetic field. The amount of deflection depends on:
1) the mass of the ion
2) the amount of positive charge on it
- the higher the charge, the larger the deflection.
- the two factors combine into mass/charge ratio (m/e or m/z).
- the smaller the value of m/e, the larger the deflection
v. Detection
- the data are fed into the computer and the mass spectrum is produced.
2.3 Amount of Substance
The mole and the Avogadro's constant
1) A mole of a substance is the amount of substance that contains the same amount of stated elementary units as there are atoms in 12 g of C-12.
2) The number of atoms is 12 g of C-12 is 6.02 x 10²³ . This number is also known as the Avogadro's constant, L.
3) Examples:
ii. 1 mol of CO2 contains 6.02 x 10²³ CO2 molecules but 3 x (6.02 x10²³) atoms.
iii. 1 mol of NaCl contains 6.02 x 10²³ NaCl units, Na⁺ and Cl⁻ ions.
Moles and mass
Moles and volumes
1) Volume occupied by a gas depends on the amount of gas, temperature and pressure. In other words the volume of a gas is not fixed.
2) Avogadro's law states that for equal volumes of all gases, under the same conditions, contain the same number of moles.
3) Hence, equal number of moles of any gas, under the same conditions, would occupy the same volume. It does not depend on the nature of gas.
4) At room temperature of 20 ℃ and a pressure of 1 atm, one mole of any gas occupies 24 dm³.
5) At standard temperature and pressure (s.t.p), which is 0 ℃ and 1 atm, one mole of any gas occupies 22.4 dm³.
6) i. Complete combustion of hydrocarbon produces water and carbon dioxide.
ii. In incomplete combustion, the possible products are carbon dioxide, carbon monoxide, carbon soot and water.
Moles and concentration of solutions
1) A solution is a homogeneous mixture of two or more substance.
2) The substance presents in small quantity is called the solute while the substance present is larger quantity is called the solvent.
3) Concentration is the amount of solute present in a fixed quantity of solution.
4) Concentration is expressed in terms of g dm⁻³. Concentration in mol dm⁻³ is called molar concentration or molarity.
2.4 Empirical Formula and Molecular Formula
Percentage composition by mass
Concentration / g dm⁻³
Mass of solute / g
Volume of solution / dm³
Percentage composition by mass / % = Ar x No.of mole of that element /Molar mass of compound X 100%
Empirical formula
1) Empirical formula is a chemical formula that shows the simplest ratio of the atoms that combine to form a molecule.
2) Steps to find empirical formula:
ii. Find the number of mole of each element (divide by its Ar).
iii. Find the simplest ratio (divide by the smallest number).
iv. Construct the empirical formula using the simplest ratio.
3) Some facts:
ii. The empirical formula and molecular formula for simple inorganic molecules are often the same.
iii. Organic molecules have different empirical and molecular formula.
Molecular formula
1) Molecular formula is a chemical formula that shows the actual number of atoms that combine to form the compound.
2) In order to deduce the molecular formula of a compound, we need to know:
ii. the empirical formula of the compound.
Principle of conservation of mass
1) Mass is neither created nor destroyed during a chemical reaction. Therefore the total mass of the reactants is equal to the total of the products in a closed system.
2) For example, the total mass of iodine in the reactants is equal to the total mass of iodine in the products.
3) This can be used to solve problems in calculating the empirical formula.
2.5 Stoichiometry and Equations
Stoichiometry
1) Stoichiometry is the proportion of things either reacting or combining.
2) In compounds, it refers to the ratio in which the atoms are combined together. For example, water, H2O has a stoichiometry of 2 hydrogen to 1 oxygen.
3) It also refers to the reacting proportions in a chemical equation. For example: 2H2 + O2 → 2H2O
The stoichiometry shows that 2 moles of hydrogen react with 1 mole of oxygen to form 2 moles of water.
Ionic equations
1) Steps to construct net ionic equations:
ii. Write the complete ionic equation by splitting it into ions(if possible).
iii. Cancel out the spectator ions. (Spectator ions are ions that present in the mixture but do not participate in the reaction.)
iv. Write down the 'leftovers', that is the net ionic equation.
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