Elements, compounds and mixtures Revision Notes

The Valence Electrons of Carbon

Carbon is in Group 14 of the periodic table. This means it has four electrons in its outermost electron shell, also known as its valence shell. To achieve a stable electron configuration, which usually means having eight electrons in the outer shell (an octet rule), carbon needs to gain four more electrons. This tendency dictates its bonding behaviour and its fundamental role in organic chemistry.

Formation of Covalent Bonds

Instead of gaining or losing electrons to form ions, which would require a large amount of energy due to its position in the middle of the periodic table, carbon shares its valence electrons with other atoms. This sharing of electrons forms covalent bonds. Each covalent bond involves two atoms sharing a pair of electrons, with one electron often contributed by each atom involved in the bond. This results in both atoms achieving a more stable electron arrangement.

Carbon's Ability to Form Four Covalent Bonds

Since carbon has four valence electrons and needs four more to complete its outer shell, it can form exactly four covalent bonds. These four bonds can be single bonds, double bonds, or triple bonds, or a combination thereof. For example, in methane (CH₄), carbon forms four single covalent bonds with four hydrogen atoms. In carbon dioxide (CO₂), carbon forms two double covalent bonds with two oxygen atoms. This tetravalency (ability to form four bonds) is fundamental to the structure of all organic molecules and their vast diversity.

What are Formulations?

Formulations are mixtures that have been deliberately designed and created to have specific properties or a precise composition for a particular application. Unlike naturally occurring mixtures or pure substances, formulations are engineered by combining different substances in carefully controlled proportions. Each component is chosen to contribute a desired characteristic to the final product. This allows for the creation of materials that perform better or differently than any single component on its own.

Why are Formulations Important?

Many useful materials are formulations because single pure substances often do not possess all the necessary properties for a specific task. For example, a pure metal might be too soft, or a single chemical might be too volatile for a particular use. By combining multiple substances, scientists can achieve an optimised balance of properties like strength, flexibility, colour, conductivity, viscosity, or reactivity. This precise control over properties makes formulations indispensable in various industries, from pharmaceuticals to construction materials.

Examples of Common Formulations

Everyday examples of formulations include paints, fuels, cleaning agents, cosmetics, fertilisers, and medicines. Paints are formulations of pigments, binders, solvents, and additives, each contributing to colour, adhesion, spreadability, and durability. Fuels are blends of hydrocarbons with specific additives to control burn rate and efficiency. Cleaning products combine surfactants, disinfectants, and fragrances. Medicines include an active drug alongside excipients to ensure correct dosage, stability, and absorption. The exact composition of each formulation is crucial for its intended function.

The Purpose of Components in Formulations

Each substance within a formulation serves a specific purpose. For instance, in a medicine, the active ingredient provides the therapeutic effect, while other components (excipients) might act as binders to hold a tablet together, fillers to give it bulk, or coatings to control drug release. In detergents, surfactants reduce surface tension for cleaning, while enzymes can break down specific stains. The careful selection and precise proportioning of these components allow for the creation of highly effective and tailored products designed for diverse needs.

Ionic Bonding: Formation and Structure

Ionic compounds are formed between metals and non-metals. Metal atoms lose electrons to form positive ions, while non-metal atoms gain electrons to form negative ions. These oppositely charged ions are attracted to each other by strong electrostatic forces, forming an ionic bond. These bonds result in a giant ionic lattice structure, where each ion is surrounded by ions of the opposite charge. This arrangement leads to high melting and boiling points because a large amount of energy is required to overcome these strong electrostatic forces.

Covalent Bonding: Simple Molecules and Giant Structures

Covalent bonds occur between non-metal atoms, where electrons are shared to achieve a stable outer shell. Simple molecules, like H₂O or CO₂, have strong covalent bonds within the molecule but weak intermolecular forces between individual molecules. This leads to low melting and boiling points. In contrast, giant covalent structures, such as diamond or silicon dioxide, involve a vast network of covalently bonded atoms. These structures are very strong and have very high melting and boiling points because many strong covalent bonds must be broken.

Polymers and Metallic Bonding

Polymers are long chain molecules formed from many repeating smaller units (monomers) joined by strong covalent bonds. The forces between polymer chains can vary, affecting their properties. Metals consist of positive metal ions arranged in a regular lattice, surrounded by a 'sea' of delocalised (mobile) electrons. The strong electrostatic attraction between the positive ions and the delocalised electrons constitutes metallic bonding. This structure gives metals their characteristic properties, including good electrical and thermal conductivity, and malleability.

Covalent Bonds: Sharing Electrons

In covalent bonding, atoms achieve stable electron configurations by sharing pairs of electrons. A dot and cross diagram for a covalent substance shows the outermost electron shell of each atom, with dots representing electrons from one atom and crosses representing electrons from the other. The shared electrons are placed in the overlapping region of the outer shells. Non-bonding (lone pair) electrons are also shown for each atom. For example, in methane (CH₄), carbon shares one electron with each of the four hydrogen atoms, and each hydrogen atom shares one electron with the carbon atom. Each shared pair consists of one 'dot' electron and one 'cross' electron.

Ionic Bonds: Transferring Electrons

Ionic bonding involves the transfer of electrons from one atom to another, leading to the formation of oppositely charged ions. In a dot and cross diagram for an ionic substance, the atoms are shown before and after the transfer. The atom losing electrons becomes a positively charged cation, and the atom gaining electrons becomes a negatively charged anion. The resulting ions are drawn with their full outer electron shells, enclosed in square brackets, with the ionic charge written outside the brackets. For example, in magnesium oxide (MgO), magnesium (Mg) atoms transfer two electrons to oxygen (O) atoms, forming Mg²⁺ and O²⁻ ions.

Key Steps for Drawing Dot and Cross Diagrams

1. Determine the outer shell electrons for each atom involved using their group number in the periodic table. 2. For covalent substances: Place the atoms in the correct arrangement. Draw the outer shells to overlap where electrons are shared. Use dots for electrons from one atom and crosses for electrons from the other. Ensure all atoms achieve a full outer shell (usually 8 electrons, or 2 for hydrogen). 3. For ionic substances: Identify which atom will lose electrons (usually metals) and which will gain electrons (usually non-metals). Show the transfer of electrons. Draw the resulting ions in square brackets, including their full outer shells and the correct ionic charge outside the brackets. The number of electrons in the outer shell of the ions must reflect a stable configuration.