Chemical reactions Revision Notes

What are State Symbols?

State symbols are letters written in parentheses next to the chemical formula of each substance in a chemical equation. They represent the physical state of the substance at the time of the reaction. These symbols are essential for providing a complete picture of a chemical process. The four main state symbols you need to know are: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution. Understanding these symbols helps predict the nature of reactants and products.

Common State Symbols and Their Meanings

The state symbols represent distinct physical states:

• (s): This stands for solid. Solids have a fixed shape and a fixed volume. Examples include iron (Fe(s)), sodium chloride (NaCl(s)), or a precipitate formed during a reaction.
• (l): This stands for liquid. Liquids have a fixed volume but take the shape of their container. Examples include water (H₂O(l)) or molten metals.
• (g): This stands for gas. Gases have neither a fixed shape nor a fixed volume; they will expand to fill any container. Examples include oxygen (O₂(g)), carbon dioxide (CO₂(g)), or hydrogen (H₂(g)).
• (aq): This stands for aqueous solution. An aqueous solution is a substance dissolved in water. Many ionic compounds, acids, and bases are found in this state. For instance, hydrochloric acid is HCl(aq), meaning hydrogen chloride gas is dissolved in water.

Importance of State Symbols

Including state symbols in chemical equations is very important for several reasons. They provide vital information about the reaction conditions and the physical form of the substances involved. For example, knowing if a reactant is a solid or an aqueous solution might influence how it is measured or mixed. If a product is a gas, it might bubble out of the solution, or if it's a solid, it might form a precipitate. This information is key for designing and understanding chemical experiments, hazard assessment, and interpreting laboratory observations. Without them, a chemical equation provides an incomplete picture of the reaction.

What are Exothermic and Endothermic Reactions?

Chemical reactions always involve energy changes. An exothermic reaction releases more energy when new bonds are formed than is absorbed to break existing bonds. This causes heat to be given out to the surroundings, leading to a rise in temperature. Examples include combustion and neutralisation reactions. Conversely, an endothermic reaction absorbs more energy to break existing bonds than is released when new bonds form. This takes heat from the surroundings, causing the temperature of the surroundings to fall. Examples include photosynthesis and some dissolution processes.

Understanding Reaction Profile Diagrams

Reaction profiles are visual tools that show the energy changes throughout a chemical reaction. The vertical axis represents energy (sometimes labelled enthalpy), and the horizontal axis represents the 'reaction pathway' or 'progress of reaction'. Both exothermic and endothermic reactions require an initial input of energy, known as activation energy, to start the reaction. This is shown as a 'hump' or 'peak' on the reaction profile diagram.

Exothermic Reaction Profile Features

For an exothermic reaction, the energy level of the products is lower than the energy level of the reactants. The difference between these two levels represents the overall energy change (ΔH), which is released to the surroundings. The activation energy is the difference in energy between the reactants and the peak of the curve. It is important to label the reactants, products, activation energy, and the overall energy change on the diagram. The curve typically starts at a higher energy for reactants, goes up to a peak (transition state), and then down to a lower energy for products.

Endothermic Reaction Profile Features

In an endothermic reaction profile, the energy level of the products is higher than the energy level of the reactants. The difference between these levels is the overall energy change (ΔH), which is absorbed from the surroundings. Similar to exothermic reactions, the activation energy is the difference between the reactants' energy and the peak of the curve. Diagrams should be clearly labelled with reactants, products, activation energy, and overall energy change. The profile begins at a lower energy for reactants, rises to a peak, and then levels out at a higher energy for products.

Understanding the pH Scale (Higher)

The pH scale ranges from 0 to 14 and quantifies the acidity or alkalinity of an aqueous solution. A pH of 7 is neutral (e.g., pure water). pH values less than 7 indicate acidity, with lower numbers representing stronger acids. pH values greater than 7 indicate alkalinity (or basicity), with higher numbers representing stronger alkalis. The pH scale is dependent on the concentration of hydrogen ions (H⁺) in the solution. The higher the concentration of H⁺ ions, the more acidic the solution and the lower its pH value.

The Logarithmic Relationship (Higher)

Crucially, the pH scale is logarithmic, not linear. This means that each whole number step on the pH scale represents a tenfold change in the hydrogen ion concentration. Therefore, a change of one pH unit means the hydrogen ion concentration is ten times greater or ten times smaller. This has significant implications for how we perceive and measure changes in acidity. For example, a solution with a pH of 2 is ten times more acidic than a solution with a pH of 3, and one hundred times more acidic than a solution with a pH of 4.

Impact of Hydrogen Ion Concentration (Higher)

As the hydrogen ion concentration in a solution increases, the pH value of that solution decreases. This is an inverse relationship. If the hydrogen ion concentration increases by a factor of 10, the pH value will decrease by 1. For example, if a solution's H⁺ concentration is increased tenfold, its pH will drop from 5 to 4. Conversely, if the hydrogen ion concentration decreases by a factor of 10, the pH value will increase by 1. This relationship allows us to quantitatively compare the relative strengths of acids and alkalis based on their pH values.

What is Electrolysis?

Electrolysis is a process that uses electrical energy to break down ionic compounds. For this to happen, the ions must be free to move. This can occur when the ionic compound is molten (melted) or dissolved in a solvent. In the case of molten binary ionic compounds, there are only two types of ions present: positive metal ions (cations) and negative non-metal ions (anions), which allows for straightforward prediction of products.

The Electrodes and Ion Movement

There are two electrodes in an electrolytic cell: the cathode (negative electrode) and the anode (positive electrode). When an electric current is passed through the molten ionic compound, the positively charged metal cations are attracted to the negatively charged cathode. At the same time, the negatively charged non-metal anions are attracted to the positively charged anode. This directed movement is essential for the decomposition to occur.

Reactions at the Electrodes

At the cathode (negative electrode), the metal cations gain electrons and are reduced to form neutral metal atoms. For example, Na+ + e- → Na. This typically results in the deposition of a solid metal. At the anode (positive electrode), the non-metal anions lose electrons and are oxidised to form neutral non-metal atoms or molecules. For example, 2Cl- → Cl2 + 2e-. This often results in the production of a gaseous non-metal.

Predicting Products of Molten Binary Ionic Compounds

To predict the products, identify the metal cation and the non-metal anion in the compound. The metal element will always be formed at the cathode, and the non-metal element will always be formed at the anode. For example, in molten sodium chloride (NaCl), sodium metal (Na) will be produced at the cathode and chlorine gas (Cl2) will be produced at the anode. This rule applies to all molten binary ionic compounds.