Stoichiometric Relationships & The Mole Concept

Stoichiometric Relationships

• Stoichiometry: Describes the relationships between the amounts of reactants and products during chemical reactions.

• Matter is conserved during chemical change.

• Stoichiometry is a form of book-keeping at the atomic level.

• It helps chemists determine the correct amounts of substances for reactions and predict product yields.

• Chemical equations are the universal language of chemistry.

• The mole is the unit of amount.

Introduction to the Particulate Nature of Matter and Chemical Change

• Atoms of different elements combine in fixed ratios to form compounds.

• Compounds have different properties than their component elements.

• Mixtures contain more than one element or compound that are not chemically bonded.

• Mixtures retain individual properties.

• Mixtures can be homogeneous (uniform composition) or heterogeneous (non-uniform composition).

Chemical equations

• Chemical equations show reactants and products using chemical symbols.

• State symbols: (s), (l), (g), and (aq) indicate solid, liquid, gas, and aqueous states, respectively.

• Balancing equations: ensure the same number of atoms of each element on both sides.

• Example: Thermal decomposition of sodium hydrogencarbonate: $2NaHCO<em>3(s) \rightarrow Na</em>2CO<em>3(s) + H</em>2O(g) + CO_2(g)$

Balancing chemical equations

• Decomposition: $CuCO<em>3(s) \rightarrow CuO(s) + CO</em>2(g)$

• Combustion: $2Mg(s) + O_2(g) \rightarrow 2MgO(s)$

• Neutralization: $H<em>2SO</em>4(aq) + 2NaOH(aq) \rightarrow Na<em>2SO</em>4(aq) + 2H_2O(l)$

• Synthesis: $N<em>2(g) + 3H</em>2(g) \rightarrow 2NH<em>3(g)$ Combustion of methane: $CH</em>4(g) + 2O<em>2(g) \rightarrow CO</em>2(g) + 2H_2O(g)$

• Other examples:

  • $2K(s) + 2H<em>2O(l) \rightarrow 2KOH(aq) + H</em>2(g)$

  • $C<em>2H</em>5OH(l) + 3O<em>2(g) \rightarrow 2CO</em>2(g) + 3H_2O(g)$

  • $Cl<em>2(g) + 2KI(aq) \rightarrow 2KCl(aq) + I</em>2(s)$

  • $4CrO<em>3(s) \rightarrow 2Cr</em>2O<em>3(s) + 3O</em>2(g)$

  • $Fe<em>2O</em>3(s) + 3C(s) \rightarrow 3CO(g) + 2Fe(s)$

  • $2C<em>4H</em>{10}(g) + 13O<em>2(g) \rightarrow 8CO</em>2(g) + 10H_2O(g)$

  • $4NH<em>3(g) + 5O</em>2(g) \rightarrow 4NO(g) + 6H_2O(g)$

  • $3Cu(s) + 8HNO<em>3(aq) \rightarrow 3Cu(NO</em>3)<em>2(aq) + 2NO(g) + 4H</em>2O(l)$

  • $H<em>2O</em>2(aq) + N<em>2H</em>4(aq) \rightarrow N<em>2(g) + 4H</em>2O(l) + O_2(g)$

  • $4C<em>2H</em>7N(g) + 15O<em>2(g) \rightarrow 8CO</em>2(g) + 14H<em>2O(g) + 2N</em>2(g)$

Atom Economy

• Formula: $\% \text{ atom economy} = \frac{\text{mass of desired product}}{\text{total mass of products}} \times 100$

• A higher atom economy indicates a more efficient and less wasteful process.

• Relevant to green and sustainable chemistry.

Mixtures

• Mixtures: substances combined without chemical interaction.

• Components retain individual properties.

• Composition is not fixed (e.g., air: 20% oxygen in, 16% oxygen out).

• Homogeneous mixture: uniform composition (e.g., air, salt water, bronze).

• Heterogeneous mixture: non-uniform composition (e.g., water and oil).

Separation Techniques

• Sand and salt: solubility in water, use solution and filtration.

• Hydrocarbons in crude oil: boiling point, use fractional distillation.

• Iron and sulfur: magnetism, use a magnet.

• Pigments in food colouring: adsorption to solid phase, use paper chromatography.

• Different amino acids: net charge at a fixed pH, use gel electrophoresis.

States of Matter

• Matter exists in different states (solid, liquid, gas) determined by temperature and pressure.

• Kinetic theory of matter: average kinetic energy of particles is related to temperature.

• Liquids and gases are fluids (they flow).

• Diffusion: particles of a substance become evenly distributed due to random movements.

  • Kinetic Energy: $KE = \frac{1}{2}mv^2$

• State symbols: (s), (l), (g), (aq) for solid, liquid, gas, aqueous.

• Balancing equations including state symbols, e.g., $2Na(s) + 2H<em>2O(l) \rightarrow 2NaOH(aq) + H</em>2(g)$

Homogeneous or Heterogeneous

• Sand and water: heterogeneous.

• Smoke: heterogeneous.

• Sugar and water: homogeneous.

• Salt and iron filings: heterogeneous.

• Ethanol and water in wine: homogeneous.

• Steel: homogeneous.

State Changes

• Sublimation: solid to gas.

• Melting: solid to liquid.

• Vaporization: liquid to gas (evaporation and boiling).

• Freezing: liquid to solid.

• Condensation: gas to liquid.

• Deposition: gas to solid.

• Temperature remains constant during state changes as energy is used to break inter-particle forces.

• Evaporation occurs only at the surface and takes place at temperatures below the boiling point

• Boiling occurs at a volume phenomenon, particles leaving throughout the body of the liquid – which is why bubbles occur. Boiling occurs at a specifi c temperature, determined by when the vapour pressure reaches the external pressure.

Kinetic Energy

• $KE = \frac{1}{2}mv^2$

Refrigeration Example

• Propane (C3H8): boiling point -42°C

• Butane (C4H10): boiling point -1°C

• Butane is unsuitable for very cold climates.

Thermal Energy Formula Understanding

• a–b: Solid heated, vibrational energy increases, temperature increases.

• b–c: Melting point reached, energy breaks inter-particle forces, temperature constant.

• c–d: Liquid heated, kinetic energy increases, temperature increases.

• d–e: Boiling point reached, energy breaks all inter-particle forces, temperature constant.

• e–f: Gas heated, kinetic energy increases, temperature increases.

The Mole Concept

• The mole is a fixed number of particles and defines the amount, n, of a substance.

  • Avogadros constant, $L = 6.02 \times 10^{23} mol^{-1}$

• Relative atomic mass (Ar): the weighted average of one atom of an element relative to one-twelfth of an atom of carbon-12.

• Relative formula/molecular mass (Mr): the sum of the weighted average of the masses of the atoms in a formula unit relative to one-twelfth of an atom of carbon-12.

• Molar mass (M): mass of one mole of a substance, expressed in g mol-1.
  Empirical formula: the simplest ratio of atoms in a compound.

• Molecular formula: the actual number of atoms present in a molecule.

• IUPAC guidelines for naming and notation.

Avogadro's Constant and Moles

• Converting between number of particles (N) and number of moles (n): $N = n \times L$

Mass and Moles Formula

• m = mass in grams.

• M = molar mass in g. Mol-1

  Molar Mass Formula:

   $amount = {mass \over molar mass}$$n(mol) = {m(g) \over M(gmol^{-1})}$

Molar Mass Conversion Formula:

    $amount = {mass\over molar mass} n(mol) = {m(g)\over M(gmol^{-1})}$

HOW TO CONVERT FROM PARTICLES TO MASS IN GRAMS FORMAL #

• Converting between the number of particles and the mass in grams $amount = {mass\over molar mass}$

Number of particles: MASS IN GRAMS FORMULA

amount,n X avogadroes contant,L \ divide by molar mass, M
divide by Avagadro contast,L Multiply by molar mass,M