Elements, compounds and mixtures Revision Notes

Ionic Compounds

Ionic compounds are formed when metals react with non-metals, involving the transfer of electrons to form charged ions. These ions are held together by strong electrostatic forces in a giant ionic lattice structure. Due to these strong forces, a large amount of energy is required to overcome them, resulting in high melting and boiling points. Ionic compounds are typically hard and brittle. They do not conduct electricity when solid because the ions are fixed in the lattice. However, when molten or dissolved in water, the ions become mobile and can carry charge, allowing the substance to conduct electricity.

Simple Molecular Structures

Simple molecular substances, such as oxygen (O₂) or water (H₂O), consist of individual molecules. Within each molecule, atoms are held together by strong covalent bonds. However, the forces of attraction between adjacent molecules (intermolecular forces) are very weak. Only a small amount of energy is needed to overcome these weak intermolecular forces, which explains why simple molecular substances have low melting and boiling points. They do not conduct electricity because there are no free-moving ions or delocalised electrons to carry charge.

Giant Covalent Structures

Giant covalent structures, like diamond or silicon dioxide, are characterised by a vast network of atoms where all atoms are held together by strong covalent bonds. This extensive bonding means that a very large amount of energy is required to break these strong covalent bonds throughout the structure, leading to extremely high melting and boiling points. These materials are often very hard. Generally, they do not conduct electricity because all outer shell electrons are involved in bonding and are not free to move, except for graphite, which has delocalised electrons between its layers.

Metals and Polymers

Metals have a structure consisting of a regular arrangement of positive metal ions surrounded by a 'sea' of delocalised electrons. The strong electrostatic attraction between the positive ions and the delocalised electrons constitutes metallic bonding. This strong bonding results in high melting and boiling points and makes metals good conductors of heat and electricity, as the delocalised electrons can move freely. Polymers are large molecules made from many small repeating units (monomers) joined by strong covalent bonds. Thermosoftening polymers have weak intermolecular forces between chains, allowing them to melt and be reshaped when heated. Thermosetting polymers have strong covalent cross-links between their chains, making them rigid, stronger, and unable to melt when heated.

Understanding the Properties of Mixtures

The first step in choosing a purification technique is to identify the physical and chemical properties of the substances in the mixture. Key properties include: solubility (whether a substance dissolves in a solvent), boiling point (the temperature at which a liquid turns into a gas), melting point (the temperature at which a solid turns into a liquid), particle size, and density. For instance, if one component is soluble and another is not, filtration can be used. If components have significantly different boiling points, distillation or evaporation might be suitable. Knowing these properties helps in predicting how substances will behave during separation.

Common Purification Techniques and Their Applications

Several common techniques are employed for purification. Filtration is used to separate an insoluble solid from a liquid. Evaporation separates a soluble solid from a solvent by boiling off the solvent, leaving the solid behind. Crystallisation is a refined form of evaporation where a saturated solution is cooled slowly to form pure solid crystals. Simple distillation separates a liquid from a dissolved solid or two liquids with significantly different boiling points, by vaporising and then condensing the more volatile component. Chromatography (though more advanced) can separate components based on their different affinities for stationary and mobile phases.

Matching Technique to Mixture Characteristics

To select the correct technique, match the differences in properties to the separation method. If you need to recover a pure solid from a solution, and the solvent is not needed, evaporation is simple. If both the solvent and the solid are valuable, simple distillation is preferred. For separating sand from water, filtration is ideal due to the insolubility of sand. When separating oil from water, decantation or a separating funnel is used because oil and water are immiscible and have different densities. The goal is always to exploit the most distinct property difference between the components.

Scientific Definition of Purity

In chemistry, a pure substance is defined as a material made up of only one type of element or one type of compound. This means all the particles (atoms or molecules) in the substance are chemically identical. For example, pure water consists only of H₂O molecules, and pure oxygen gas consists only of O₂ molecules. There are no other substances mixed in. This strict definition is crucial for scientific experiments and industrial processes where precise chemical properties are required.

Everyday Definition of Purity

The everyday understanding of 'pure' is much broader and less stringent than the scientific definition. In common language, 'pure' often means natural, unadulterated, or free from contamination. For instance, 'pure fruit juice' might mean juice without added sugars or preservatives, even though it contains water, natural sugars, acids, and vitamins – a mixture of many compounds. 'Pure cotton' implies the fabric is 100% cotton fibres, without other textile blends, but cotton itself is a complex natural polymer.

Distinguishing by Physical Properties

The most significant way to distinguish a scientifically pure substance from an impure mixture is by examining its physical properties, specifically its melting and boiling points. A pure substance has a sharp, fixed melting point and boiling point. For example, pure water melts at exactly 0°C and boils at exactly 100°C (at standard atmospheric pressure). The presence of impurities lowers the melting point and raises the boiling point, and crucially, causes these phase changes to occur over a range of temperatures, rather than at a single, precise temperature. This phenomenon is why boiling point determination is a common method to test for purity in a laboratory.

Impact of Impurities

Impurities disrupt the regular arrangement of particles in a substance. When a substance is impure, the different types of particles present require varying amounts of energy to overcome the intermolecular forces holding them together. This leads to a melting point depression (lower melting point) and a boiling point elevation (higher boiling point) compared to the pure substance. Furthermore, instead of a sharp, single-temperature phase change, mixtures will melt and boil over an extended temperature range. This broadening of the melting/boiling range is a direct indicator of impurity.

Metals: Position and Atomic Structure

Metals are predominantly found on the left and in the central block of the Periodic Table. Their atomic structure is characterised by having a small number of electrons in their outermost shell, typically 1, 2, or 3. Due to the relatively weak attraction between the nucleus and these valence electrons, metals readily lose them during chemical reactions. This tendency to lose electrons results in the formation of positive ions and contributes to metallic bonding, which involves a 'sea' of delocalised electrons. This electron arrangement is responsible for the characteristic properties of metals, such as high electrical and thermal conductivity, malleability, and ductility.

Non-metals: Position and Atomic Structure

Non-metals are located on the right side of the staircase line in the Periodic Table. Their atomic structure typically involves having a larger number of electrons in their outermost shell, usually 4, 5, 6, 7, or 8 (except for hydrogen and helium). Unlike metals, non-metals tend to gain or share electrons to achieve a stable, full outer shell. When they gain electrons, non-metals form negative ions. This behaviour leads to properties such as being poor conductors of electricity and heat, and generally having lower melting and boiling points compared to metals. Elements in Group 0 (noble gases) already have a full outer shell, making them very unreactive.

Predicting Properties from Position

The position of an element in the Periodic Table allows for predictions about its metallic or non-metallic character. Elements in Groups 1 and 2 are strong metals, and elements in Groups 13 are also metals. As you move right across a period, elements become less metallic and more non-metallic. Elements in Groups 14 to 17 are generally non-metals. The number of outer electrons, indicated by the group number (for main group elements), is a key factor in determining whether an atom will lose or gain electrons, thus dictating its chemical behaviour as a metal or non-metal. The period number indicates the number of electron shells.

Electron Configuration and Stability

All atoms strive to achieve a stable electron configuration, typically a full outermost electron shell. Metals achieve this by losing their few valence electrons, forming positive ions. Non-metals, on the other hand, achieve stability by gaining or sharing electrons to complete their outer shell, forming negative ions or covalent bonds. This fundamental difference in electron behaviour, directly related to the number of valence electrons, underpins the distinction between metals and non-metals and their placement in the Periodic Table.