Elements, compounds and mixtures Revision Notes

The Periodic Table: Metals vs. Non-metals

The Periodic Table organises elements based on their atomic number and electron configuration. Elements are fundamentally categorised into metals and non-metals, with a dividing line of metalloids. Metals are found on the left and central parts of the Periodic Table, including Group 1, Group 2, and the transition metals. Non-metals occupy the upper-right section of the Periodic Table, including Groups 14 through to Group 18, and hydrogen. This spatial arrangement reflects underlying differences in their atomic structures and chemical behaviours.

Atomic Structure of Metals

Metals are characterised by having a small number of electrons in their outermost electron shell, typically one, two, or three. These valence electrons are relatively loosely held by the nucleus. Because of this, atoms of metallic elements have a strong tendency to lose these outer electrons during chemical reactions. By losing electrons, metals form positive ions (cations) and achieve a stable electron configuration, often resembling that of the preceding noble gas. This electronic behaviour is the basis for metallic bonding and gives metals their characteristic properties, such as good electrical and thermal conductivity, malleability, and ductility.

Atomic Structure of Non-metals

Non-metals, on the other hand, generally have a larger number of electrons in their outermost electron shell, typically four, five, six, or seven. Due to the stronger attraction between the nucleus and these valence electrons, and the desire to complete their outer shell, non-metal atoms tend to gain electrons from other atoms or share electrons to form covalent bonds during chemical reactions. When non-metals gain electrons, they form negative ions (anions). When they share electrons, they form molecules. This behaviour explains why non-metals are typically poor conductors of electricity and heat, are often brittle, and have varied physical states at room temperature.

What is Relative Formula Mass (Mr)?

The relative formula mass (Mr) is the sum of the relative atomic masses (Ar) of all the atoms in the chemical formula of a substance. It's a relative value, meaning it doesn't have units when expressed as Mr. For example, if you have a compound like carbon dioxide (CO₂), you would add the Ar of one carbon atom to the Ar of two oxygen atoms. The Ar values for elements are typically found on the periodic table.

How to Calculate Mr for Simple Compounds

To calculate the Mr of a simple compound, identify the chemical formula. Then, look up the relative atomic mass (Ar) for each element in the formula. Multiply the Ar of each element by the number of times it appears in the formula. Finally, sum these values. For example, to calculate Mr for methane (CH₄): Ar(C) = 12, Ar(H) = 1. So, Mr(CH₄) = (1 × 12) + (4 × 1) = 12 + 4 = 16. It's crucial to be careful with subscripts in the chemical formula.

Calculating Mr for Compounds with Brackets

When a chemical formula contains brackets, the number outside the brackets multiplies everything inside them. For example, in calcium hydroxide, Ca(OH)₂, there is one calcium atom, two oxygen atoms, and two hydrogen atoms. To calculate Mr for Ca(OH)₂: Ar(Ca) = 40, Ar(O) = 16, Ar(H) = 1. So, Mr(Ca(OH)₂) = (1 × 40) + (2 × 16) + (2 × 1) = 40 + 32 + 2 = 74. Always distribute the number outside the bracket to each atom inside before summing.

Mr in Balanced Chemical Equations

Relative formula masses can be calculated for all reactants and products in a balanced chemical equation. Although the Mr values themselves don't change due to the balancing numbers, these values are essential for further stoichiometric calculations involving moles and masses. For instance, in 2 H₂ + O₂ → 2 H₂O, you would still calculate Mr(H₂) = 2, Mr(O₂) = 32, and Mr(H₂O) = 18. The coefficients in the balanced equation (e.g., the '2' in 2 H₂) tell you the ratio of moles, not how to calculate the individual Mr.

States of Matter and Particle Arrangement

Solids have particles arranged in fixed positions, held by strong forces. They vibrate but do not move past each other. Liquids have particles that are close together but can move past each other, held by weaker forces than in solids. Gases have particles that are far apart and move randomly and rapidly, with very weak or negligible forces between them. Understanding particle arrangement helps to explain the energy needed for state changes.

Energy Transfer During Melting and Boiling

When a solid melts, energy is absorbed from the surroundings. This energy is used to overcome the forces holding the particles in their fixed positions, allowing them to move about more freely as a liquid. The temperature of the substance remains constant during melting, as all absorbed energy is used for this change of state. Similarly, during boiling, energy is absorbed to completely overcome the forces between liquid particles, allowing them to escape as a gas. The temperature also remains constant during boiling despite continuous energy input.

Role of Intermolecular Forces and Chemical Bonds

The strength of the forces between particles dictates the energy needed for a change of state. Strong 'chemical bonds' (ionic or strong covalent) require a lot of energy to break, leading to very high melting and boiling points. Substances with simple molecular structures have 'intermolecular forces' between their molecules, which are much weaker than covalent or ionic bonds. Less energy is needed to overcome these weaker forces, resulting in lower melting and boiling points. For example, covalent compounds like oxygen are gases at room temperature due to very weak intermolecular forces.

Explaining Different Melting/Boiling Points

Different substances have different melting and boiling points because the strength of the forces holding their particles together varies. If a substance has strong forces (e.g., ionic compound, giant covalent structure), a large amount of energy must be supplied to overcome these forces for melting or boiling to occur, resulting in high temperatures for these changes. If a substance has weak forces (e.g., simple molecular substance), less energy is needed, leading to lower melting and boiling points.

Physical Properties of Metals

Metals are typically characterised by several key physical properties. They are usually shiny, meaning they have a metallic lustre. Most metals are solid at room temperature, with the exception of mercury. They are good conductors of heat and electricity, allowing energy to pass through them efficiently. Metals are also malleable, which means they can be hammered or pressed into different shapes without breaking. Furthermore, they are ductile, meaning they can be drawn into thin wires. Generally, metals have high densities, high melting points, and high boiling points.

Physical Properties of Non-metals

Non-metals exhibit physical properties that are largely opposite to those of metals. They are typically dull in appearance and are often brittle, meaning they will shatter or break when an attempt is made to change their shape. Non-metals are poor conductors of both heat and electricity; they act as insulators. Unlike metals, non-metals can exist as solids, liquids (like bromine), or gases (like oxygen and nitrogen) at room temperature. They generally have low densities, low melting points, and low boiling points. Graphite is a unique non-metal that conducts electricity.

Chemical Properties of Metals

Chemically, metals have distinct behaviours. They tend to lose electrons in chemical reactions to form positive ions (cations). This electron loss makes them reactive. When metals react with oxygen, they typically form basic oxides. For example, sodium reacts with oxygen to form sodium oxide, which is a basic oxide. Many metals also react with acids to produce hydrogen gas and a salt, indicating their chemical reactivity.

Chemical Properties of Non-metals

Non-metals show different chemical behaviours compared to metals. They tend to gain or share electrons in chemical reactions. When non-metals react with metals, they typically gain electrons to form negative ions (anions). For instance, chlorine (a non-metal) reacts with sodium (a metal) to form sodium chloride, where chlorine gains an electron. When non-metals react with oxygen, they generally form acidic oxides or neutral oxides, such as carbon dioxide (acidic) or carbon monoxide (neutral).