Chemical reactions Revision Notes

What are State Symbols?

State symbols are abbreviations used in chemical equations to indicate the physical state of each reactant and product. They are written in parentheses immediately after the chemical formula. These symbols are essential for providing a complete and accurate description of a chemical reaction, indicating whether a substance is solid, liquid, gas, or dissolved in water.

The Four Main State Symbols

There are four primary state symbols you need to know:

• (s): Represents a solid. Solids have a fixed shape and volume. Examples include metals like Fe(s) (iron) or ionic compounds like NaCl(s) (sodium chloride) before dissolving.
• (l): Represents a liquid. Liquids have a fixed volume but take the shape of their container. Examples include pure water, H₂O(l), or ethanol, C₂H₅OH(l).
• (g): Represents a gas. Gases have no fixed shape or volume and will expand to fill any container. Examples include oxygen, O₂(g), or carbon dioxide, CO₂(g).
• (aq): Represents an aqueous solution. This means the substance is dissolved in water. This is common for many ionic compounds and acids when mixed with water, such as NaCl(aq) (sodium chloride dissolved in water) or HCl(aq) (hydrochloric acid).

Importance of State Symbols

State symbols are crucial for several reasons. They help chemists understand the conditions under which a reaction occurs, such as whether reactants need to be dissolved or heated to a gaseous state. They also provide insight into the products' physical properties and how they might be observed or collected. For example, a reaction producing a solid (s) would likely result in a precipitate, while one producing a gas (g) would show effervescence (bubbling).

Exothermic Reactions: Releasing Heat

Exothermic reactions are chemical processes that release energy to their surroundings, typically in the form of heat. This energy release causes the temperature of the surroundings to increase. You would observe a thermometer placed in an exothermic reaction showing a rise in temperature. Examples include combustion (burning of fuels), neutralisation reactions (acid + alkali), and oxidation of metals. Many everyday phenomena, like a burning candle or a hot pack, involve exothermic reactions. The total energy of the products is lower than the total energy of the reactants, with the excess energy being dissipated into the surroundings.

Endothermic Reactions: Absorbing Heat

Endothermic reactions are chemical processes that absorb energy from their surroundings, typically in the form of heat. This energy absorption causes the temperature of the surroundings to decrease. If you place a thermometer in an endothermic reaction, you would observe a fall in temperature. Examples include dissolving certain salts like ammonium nitrate in water, photosynthesis, and the decomposition of limestone. Cold packs, often used for sports injuries, utilise endothermic reactions to create a cooling effect. The total energy of the products is higher than the total energy of the reactants, as energy has been taken in from the surroundings.

Identifying Reactions by Temperature Change (PAG C5)

A simple and reliable way to distinguish between endothermic and exothermic reactions is to measure the temperature change of the surroundings. For this practical activity (related to PAG C5), you would typically mix reactants in an insulated container (like a polystyrene cup) and use a thermometer to monitor the temperature over time. An initial temperature reading is taken, and then subsequent readings are noted as the reaction proceeds. A consistent rise in temperature indicates an exothermic reaction, while a sustained fall in temperature indicates an endothermic reaction. The extent of the temperature change can also give an indication of the magnitude of the energy transfer.

Key Differences Summarised

The fundamental difference between endothermic and exothermic reactions lies in the direction of heat energy transfer. Exothermic reactions release heat *out* into the surroundings, leading to a temperature *increase*. Endothermic reactions absorb heat *in* from the surroundings, leading to a temperature *decrease*. Think of 'exo' meaning 'out' and 'endo' meaning 'in'. This distinction is based entirely on the effect observed in the environment surrounding the chemical system, not on the internal energy changes within the reactants and products themselves.

Carbonates Reacting with Acids

When a carbonate reacts with an acid, the products are always a salt, water, and carbon dioxide. This is a reliable way to identify carbonates. The specific salt formed depends on the metal in the carbonate and the acid used. For example, sodium carbonate reacting with nitric acid will produce sodium nitrate, water, and carbon dioxide. The carbon dioxide gas is observed as bubbles (effervescence).

Metals Reacting with Acids

Some metals, specifically those more reactive than hydrogen, react with acids to produce a salt and hydrogen gas. Examples of such metals include magnesium, zinc, and iron. The vigour of the reaction depends on the metal's position in the reactivity series; more reactive metals react faster. For instance, zinc and sulfuric acid produce zinc sulfate and hydrogen gas. Metals less reactive than hydrogen (e.g., copper, silver) generally do not react with dilute acids.

Identifying Reaction Products

The presence of carbon dioxide can be confirmed by bubbling the gas produced through limewater (calcium hydroxide solution), which will turn cloudy. Hydrogen gas is identified by its characteristic 'squeaky pop' sound when a lit splint is held near it. Knowing these tests helps to verify the products of these reactions experimentally.

Writing Balanced Chemical Equations

To write a balanced chemical equation, you must ensure that there is an equal number of atoms of each element on both sides of the reaction arrow. For example, for magnesium carbonate reacting with hydrochloric acid: MgCO₃ + 2HCl → MgCl₂ + H₂O + CO₂. Here, there is one Mg, one C, three O, two H, and two Cl atoms on both sides, demonstrating conservation of mass.

Understanding the Law of Conservation of Mass

The Law of Conservation of Mass states that during any chemical reaction, atoms are rearranged but are never created or destroyed. This means the total mass of the reactants must always be equal to the total mass of the products when the system is closed. In simpler terms, no mass is lost or gained in a chemical reaction; it's simply transformed from one form to another. This fundamental principle is crucial for understanding mass changes in both enclosed and non-enclosed systems. The particle model supports this by showing that chemical reactions involve breaking and forming bonds between existing atoms, not the disappearance or creation of new ones.

Mass Changes in Non-Enclosed Systems

In a non-enclosed system (e.g., an open beaker or an object exposed to air), an observed change in mass does not mean the Law of Conservation of Mass has been violated. Instead, it indicates that substances have either entered or left the system being measured. If a gaseous product is formed during a reaction and allowed to escape into the atmosphere, the measured mass of the remaining substances in the container will decrease. An example is the reaction of an acid with a carbonate, which produces carbon dioxide gas that floats away. Conversely, if a gaseous reactant from the atmosphere is consumed in the reaction, the measured mass of the solid or liquid products can increase. For instance, when metals like iron react with oxygen from the air (rusting), the mass increases because oxygen atoms from the air become part of the solid product, effectively adding mass to the system.

Explaining Mass Changes Using the Particle Model

The particle model provides a clear explanation for observed mass changes in non-enclosed systems. When a gas escapes, it means reactant particles have been rearranged to form new gaseous product particles, which then move out of the system. Because these particles are no longer present within the boundaries of the measured system, the total mass measured appears to decrease. For example, during combustion of a candle, solid wax particles and gaseous oxygen particles react to form gaseous carbon dioxide and water vapour particles. These gaseous products escape, causing the candle's mass to decrease. Conversely, if a gas from the surroundings is incorporated into a solid or liquid product, its particles become part of the measured system, increasing its total mass. The overall number and type of atoms remain constant in the universe, demonstrating the conservation of mass, even if the observed mass within a specific non-enclosed system changes.