Chemical reactions Revision Notes

What is the Law of Conservation of Mass?

The Law of Conservation of Mass states that mass is neither created nor destroyed during a chemical reaction. This means the total mass of the substances you start with (reactants) will always be equal to the total mass of the substances you end up with (products). Atoms are simply rearranged, not lost or gained. For example, if solid sodium reacts with chlorine gas to form sodium chloride, the total mass of sodium and chlorine before the reaction will be the same as the total mass of sodium chloride formed.

Applying the Law in Chemical Reactions

When a chemical reaction occurs, the atoms are rearranged into new combinations, but the same atoms are always present. This is why chemical equations must be balanced; to show that the number of atoms of each element on the reactant side is exactly equal to the number of atoms of each element on the product side. If you know the mass of the reactants, you can predict the total mass of the products. For example, if 20g of hydrogen reacts with oxygen, and 180g of water is formed, then the mass of oxygen that reacted must be 160g (180g - 20g = 160g).

Importance for Closed vs. Open Systems

The law of conservation of mass holds true whether the reaction container is open or closed. However, measuring the mass change might appear different. In a closed system (where nothing can enter or leave), the total mass of the system will physically remain constant. If a gas is produced or consumed in a closed system, its mass is counted within the sealed container. In an open system, if a gas is produced, it might escape, making the system appear to lose mass. Conversely, if a gas from the environment (like oxygen) is consumed, the system might appear to gain mass. Even in these cases, the overall mass of all substances involved (including gases that escaped or entered) is conserved.

What is Electrolysis?

Electrolysis is the process of using electrical energy to break down ionic compounds. This occurs by passing an electric current through molten or dissolved ionic substances, known as electrolytes. The electrical energy forces non-spontaneous chemical reactions to happen, resulting in the decomposition of the compound into simpler substances or elements. It's a fundamental industrial process used for extracting reactive metals, refining metals, and electroplating.

Components of an Electrolytic Cell

An electrolytic cell consists of three main components:

1. Electrolyte: The molten or dissolved ionic compound that conducts electricity. It contains free-moving ions.
2. Electrodes: Two electrical conductors immersed in the electrolyte. The positive electrode is the anode, and the negative electrode is the cathode.
3. Power Supply: Provides the electrical energy to drive the non-spontaneous reactions. It pushes electrons into the cathode and pulls them from the anode.

Role of Electrodes: Inert vs. Non-Inert

Electrodes are crucial in electrolysis and their reactivity determines the products.

• Inert Electrodes: These electrodes, like graphite (carbon) or platinum, are unreactive. They do not take part in the chemical reaction themselves; they simply provide a surface for ions to gain or lose electrons. For example, in the electrolysis of aqueous sodium chloride with inert electrodes, chlorine gas forms at the anode and hydrogen gas at the cathode.
• Non-Inert (Active) Electrodes: These electrodes, like copper, are chemically reactive and can participate in the reaction. For instance, in the electrolytic refining of copper, the impure copper acts as the non-inert anode, dissolving into copper ions, while pure copper is deposited at the cathode.

Defining Oxidation and Reduction (Higher)

Oxidation and reduction describe processes involving electron transfer. Oxidation is the loss of electrons by a chemical species. This often results in an increase in the oxidation state or a more positive charge. For example, a metal atom losing electrons to become a positive ion is an oxidation. Reduction is the gain of electrons by a chemical species. This typically leads to a decrease in the oxidation state or a more negative charge. For instance, a non-metal atom gaining electrons to become a negative ion is a reduction. Understanding 'OIL RIG' (Oxidation Is Loss, Reduction Is Gain) is crucial.

Identifying Oxidised and Reduced Species (Higher)

To identify which species is oxidised and which is reduced in a redox reaction, focus on the electron changes. The species that releases electrons is oxidised, demonstrating electron loss. The species that accepts these electrons is reduced, showing electron gain. For example, in the reaction Cu(s) + 2Ag⁺(aq) → Cu²⁺(aq) + 2Ag(s), copper goes from a neutral atom to a 2+ ion, meaning it has lost two electrons, so copper is oxidised. Silver ions go from a 1+ charge to neutral atoms, meaning each Ag⁺ ion has gained one electron, so silver ions are reduced.

Redox Reactions Always Occur Together (Higher)

Oxidation and reduction are inseparable processes. One cannot happen without the other, as electrons must be transferred from one substance to another. The total number of electrons lost during oxidation must equal the total number of electrons gained during reduction. This ensures that the overall charge is conserved in the chemical reaction. These reactions are often represented using half-equations, showing either the electron loss (oxidation half-equation) or electron gain (reduction half-equation) separately, which can then be combined to form the overall ionic equation.

The pH Scale

The pH scale ranges from 0 to 14. A pH of 7 indicates a neutral solution (e.g., pure water). Solutions with a pH less than 7 are acidic, with lower numbers indicating stronger acidity. Solutions with a pH greater than 7 are alkaline (also called basic), with higher numbers indicating stronger alkalinity. Extremes of the scale, such as pH 0-1 and pH 13-14, represent very strong acids and alkalis, respectively.

Using Universal Indicator

Universal indicator is a mixture of dyes that produces a range of colours depending on the pH of the solution. To use it, a few drops are added to the solution to be tested. The colour produced is then compared to a colour chart to estimate the pH. While convenient and inexpensive, universal indicator provides only an approximate pH value, typically to the nearest whole number. For instance, red/orange suggests acid, green suggests neutral, and blue/purple suggests alkali.

Using a pH Meter

A pH meter is an electronic device used for precise pH measurements. It consists of a glass electrode probe connected to a digital display. Before use, the pH meter must be calibrated using buffer solutions of known pH values (e.g., pH 4, 7, 10). The probe is immersed in the solution, and a stable numerical pH reading is obtained. pH meters provide a much more accurate and precise measurement of pH compared to universal indicator, often to two decimal places, which is crucial for experiments requiring exact pH control.

Comparison of Techniques

Universal indicator is simple, quick, and cost-effective, offering an approximate pH reading. It is suitable for general classification of substances as acidic, neutral, or alkaline. A pH meter offers significantly higher accuracy and precision, providing exact numerical values. However, it is more complex, requires calibration, and is more expensive. The choice of technique depends on the required accuracy for the particular task or experiment.