Key concepts in chemistry Revision Notes

Properties of Metals

Metals are typically characterised by several key physical properties. They are generally shiny (lustrous) in appearance and exist as solids at room temperature (with the notable exception of mercury). Metals have high melting points, meaning they require considerable heat to turn into a liquid. They also tend to have high densities, packing a lot of mass into a given volume. A defining feature is their excellent electrical conductivity, which is due to delocalised electrons that are free to move throughout the metallic structure.

Properties of Non-metals

Non-metals generally exhibit properties that are opposite to those of metals. Many non-metals are gases or liquids at room temperature, while solid non-metals are often brittle. They typically have low boiling points, indicating they transition to a gaseous state at relatively low temperatures. In terms of conductivity, non-metals are generally poor conductors of electricity (acting as insulators). This is because their electrons are held tightly within their atomic structures and are not free to move and carry charge.

Calculating Relative Formula Mass (Mr)

To find the relative formula mass (Mr) of a compound, add up the relative atomic masses (Ar) of every atom shown in its chemical formula. For example, in H₂O, there are two hydrogen atoms and one oxygen atom. If Ar(H) = 1 and Ar(O) = 16, then Mr(H₂O) = (2 × 1) + 16 = 18. For compounds with brackets, multiply the Ar of atoms inside the bracket by the subscript outside the bracket.

Calculating Percentage by Mass of an Element

The percentage by mass of an element in a compound shows what proportion of the compound's total mass is made up by that element. Use the formula: (Total Ar of element in compound / Mr of compound) × 100%. For instance, to find the percentage by mass of oxygen in H₂O: (16 / 18) × 100% ≈ 88.9%. This is a crucial skill for yield calculations and determining empirical formulas.

What is an Empirical Formula?

The empirical formula shows the simplest whole-number ratio of atoms of each element in a compound. For example, the molecular formula of glucose is C6H12O6, but its empirical formula is CH2O, as the ratio 6:12:6 simplifies to 1:2:1.

Calculating Empirical Formula from Reacting Masses

1. Convert the mass of each element (in grams) to moles using the formula: moles = mass / atomic mass.
2. Divide the number of moles of each element by the smallest number of moles calculated. This gives a mole ratio.
3. If the ratios are not whole numbers, multiply all ratios by the smallest integer that converts them into whole numbers. Rounding should only occur if the values are very close to a whole number or a simple fraction (e.g., 0.5, 0.33, 0.67).
4. Write the formula using these whole numbers as subscripts after each element symbol.

Calculating Empirical Formula from Percentage Composition

1. Assume a 100 g sample, so percentages directly become masses in grams. For example, 75% carbon means 75 g of carbon.
2. Follow steps 1-4 from 'Calculating Empirical Formula from Reacting Masses' using these masses.

The Law of Conservation of Mass

This fundamental law of chemistry states that during a chemical reaction, the total mass of all substances present before the reaction (reactants) is exactly equal to the total mass of all substances present after the reaction (products). Mass cannot be created or destroyed; it is simply rearranged into different chemical forms.

Closed Systems (e.g., Sealed Flask)

In a closed system, such as a reaction in a sealed flask, no matter can enter or leave. When a chemical reaction occurs, like a precipitation reaction, the overall mass of the system remains constant. For example, if magnesium reacts with oxygen in a sealed container, the mass of the container and its contents will not change, even though new magnesium oxide is formed.

Non-Enclosed Systems (e.g., Open Flask)

In an open system, matter can transfer between the system and its surroundings. If a reaction produces a gas that escapes into the atmosphere, the measured mass of the flask and its remaining contents will decrease. For instance, magnesium reacting with acid in an open flask releases hydrogen gas, causing the total mass of the flask to appear to decrease. Conversely, if a reaction consumes a gas from the atmosphere, such as a metal rusting by reacting with oxygen, the measured mass of the solid product will increase because it has incorporated mass from the air.