Chemical reactions Study Guide

Chemical reactions involve changes in energy, specifically the breaking and forming of chemical bonds. These energy changes manifest as either a release or absorption of heat from the surroundings. We classify reactions into two main types based on this heat exchange: exothermic and endothermic. Exothermic reactions are processes that release energy, usually as heat, into their surroundings. This causes the temperature of the surroundings to increase. Common examples include combustion (burning) and neutralisation reactions. When you feel a burning sensation or observe a rise in temperature, you are likely experiencing an exothermic reaction. Conversely, endothermic reactions are processes that absorb energy, usually as heat, from their surroundings. This absorption of heat causes the temperature of the surroundings to decrease, making them feel colder. Examples include dissolving certain salts, such as ammonium nitrate in water, or processes like photosynthesis. The key distinction lies in observing the temperature change of the environment around the reaction. A thermometer placed in the reacting mixture will clearly show whether the temperature is rising (exothermic) or falling (endothermic), providing a direct indicator of the energy transfer. This fundamental concept is crucial for understanding chemical processes and their applications in everyday life and industry.

Electrolysis is a process that uses electrical energy to drive non-spontaneous chemical reactions. For ionic compounds, this means using electricity to break them down into their constituent elements. When a binary ionic compound, such as sodium chloride (NaCl) or lead bromide (PbBr₂), is in a molten state, its ions are free to move. This mobility is crucial for electrolysis to occur, as solid ionic compounds do not conduct electricity. The molten compound acts as the electrolyte.

During electrolysis, an electrical current is passed through the molten ionic compound. This involves two electrodes: the anode (positive electrode) and the cathode (negative electrode). The positively charged ions, called cations, are attracted to the negative cathode. At the cathode, they gain electrons (reduction) and are converted into neutral atoms of the metal. For example, in molten sodium chloride, Na⁺ ions move to the cathode, gain an electron, and become sodium metal (Na). The half-equation is Na⁺ + e⁻ → Na.

Conversely, the negatively charged ions, called anions, are attracted to the positive anode. At the anode, they lose electrons (oxidation) and are converted into neutral atoms or molecules of the non-metal. For example, in molten sodium chloride, Cl⁻ ions move to the anode, lose an electron, and become chlorine gas (Cl₂). The half-equation is 2Cl⁻ → Cl₂ + 2e⁻. It's important to remember that non-metal elements often form diatomic molecules. By understanding these fundamental principles, one can predict the products formed at each electrode for any given molten binary ionic compound.