Quantitative Chemistry (AQA)
1. The Mole Concept
The mole is the unit used in chemistry to measure the amount of a substance. It's a way of counting particles, much like saying "a dozen" means 12 items. A mole corresponds to Avogadro’s constant, which is 6.022 × 10²³ particles (atoms, molecules, ions, etc.).
Example: If you have 1 mole of water (H₂O), you have 6.022 × 10²³ molecules of water.
Molar Mass: The molar mass of a substance is the mass of one mole of that substance. The unit for molar mass is g/mol. This can be calculated by adding the atomic masses of the elements in a compound.
For example, the molar mass of water is:
This means that 1 mole of water weighs 18 grams.
Moles and Molecules: Moles allow you to relate the mass of a substance to the number of particles. For example, if you have 18 grams of water (H₂O), you have exactly 1 mole of water molecules.
2. Balancing Chemical Equations
In chemistry, balancing equations is crucial because it reflects the principle of conservation of mass: matter cannot be created or destroyed, only transformed. A balanced chemical equation shows that the number of atoms of each element is the same on both sides of the equation.
Example:
Unbalanced equation:
Balanced equation:
In the balanced equation, you can see that there is 1 carbon (C) on both sides, 4 hydrogens (H) on both sides, and 4 oxygens (O) on both sides. This satisfies the law of conservation of mass.
3. Conservation of Mass
This is the principle that in a closed system, the total mass of reactants is always equal to the total mass of products in a chemical reaction. Mass is neither created nor destroyed.
Example: If you start with 100 g of reactants, after the chemical reaction, you will end up with exactly 100 g of products (assuming no mass is lost to the surroundings). If a reaction occurs in a container, and you lose mass (for example, if gas escapes), the container will record less mass, but the total mass of the system is still conserved.
4. Calculating Reacting Quantities
In quantitative chemistry, you need to calculate the amounts of reactants or products using moles and the molar ratios from a balanced equation. The key is understanding the relationship between moles and mass, using the equation:
Moles = Mass (g) / Molar Mass (g/mol)
This allows you to calculate how much of one substance will react with another in a chemical equation.
Example:
In the reaction between sodium (Na) and chlorine (Cl₂) to form sodium chloride (NaCl):If you have 46 g of sodium (Na), how many moles of sodium do you have?
Moles of Na =
5. Empirical and Molecular Formulas
Empirical Formula: The empirical formula represents the simplest whole number ratio of elements in a compound. It doesn't tell you the actual number of atoms in a molecule but rather the simplest ratio.
Example: The empirical formula of hydrogen peroxide (H₂O₂) is HO, because the simplest ratio of hydrogen to oxygen is 1:1.
Molecular Formula: This shows the exact number of atoms of each element in a molecule. The molecular formula is a multiple of the empirical formula.
Example: The molecular formula of hydrogen peroxide (H₂O₂) shows that there are 2 hydrogen atoms and 2 oxygen atoms.
6. Concentration of Solutions
Concentration refers to how much solute (the substance being dissolved) is present in a solution. It is measured in moles per cubic decimeter (mol/dm³). The formula to calculate concentration is:
Concentration (mol/dm³) = Moles of Solute / Volume of Solution (dm³)
Example:
If you dissolve 1 mole of salt (NaCl) in 2 dm³ of water, the concentration is:
This is useful for determining how much solute is present in a solution and for performing calculations like titrations.
7. Titration
Titration is a laboratory technique used to determine the concentration of a solution by reacting it with a solution of known concentration. An indicator is often used to signal the end point of the reaction, where the amount of reactant in the unknown solution is exactly neutralized.
Example:
To find the concentration of hydrochloric acid (HCl), you might titrate it with a sodium hydroxide (NaOH) solution of known concentration. When the acid is neutralized, the indicator changes color, and you can calculate the concentration of HCl based on the volume of NaOH used.
8. Percent Yield and Atom Economy
Percent Yield: The percent yield measures the efficiency of a chemical reaction. It compares the actual amount of product obtained (experimental yield) to the theoretical amount of product that should be formed (theoretical yield).
Formula:
Example: If a reaction theoretically produced 10 g of product, but you only obtain 8 g, your percent yield is:
Atom Economy: This concept measures the efficiency of a chemical reaction in terms of how much of the reactants are turned into useful products. High atom economy means more of the reactants are used to produce the desired product, making the process more efficient and sustainable.
Formula:
Example: A reaction with high atom economy is one where the majority of the reactants are converted into the desired product, minimizing waste.
9. Gas Volumes and the Ideal Gas Equation
The ideal gas law relates the properties of gases (pressure, volume, temperature, and moles) in a mathematical equation. Under standard conditions (0°C and 1 atm), 1 mole of any gas occupies 22.4 dm³.
Formula: pV=nRTpV = nRTpV=nRT Where:
p = pressure (Pa)
V = volume (m³)
n = number of moles
R = ideal gas constant (8.31 J/mol·K)
T = temperature (K)
This formula is used to calculate quantities of gases in reactions when conditions deviate from standard temperature and pressure (STP).
10. Limitations of the Ideal Gas Equation
The ideal gas equation assumes that gases are made of particles that have no volume and experience no intermolecular forces. However, real gases have volume, and intermolecular forces do exist. These deviations become noticeable at high pressures and low temperatures.
Example: At very high pressures, gas molecules are closer together, and their volume becomes significant. At low temperatures, intermolecular forces can cause the gas to condense into a liquid.