Topic 1: Chemical Equilibrium Systems

• Irreversible Reactions: A reaction where the reactants convert into products but cannot convert back into reactants
• Collision Theory: Explains why reactions occur faster when there are more particles present, as it increases the chance that two molecules will collide with enough energy to cause a reaction.
• Reversible Reactions: Reactions in which the reactants convert into products and where the products can convert back into reactants
• Open System: allows matter and energy to become exchanged with it's surroundings
• Closed system: does not allow any exchange of matter or energy
• Activation Energy: The minimum amount of energy required for a chemical reaction to take place
• Dynamic Equilibrium: can only occur with a closed system
• in equilibrium reactions, the concentration of products increase
• the rate of forward reaction is equal to the rate of reverse reaction at dynamic equilibrium
• The catalyst lowers activation energy but doesn’t change the position of equilibrium
• at the start of the equilibrium reaction, the product concentration is always high
• Le Chatelier's Principle: when a chemical equilibrium is disturbed by changing conditions, the system will react in such a way to counteract the change. A new equilibrium is then formed.
• Changes which affect systems include, Concentration of Products and Reactants, Temperature and Pressure.
• Factors that increase concentration: More solute is added to the solution, Increasing the pressure of the gas by decreasing it's volume
• factors that decrease concentration: Removing some of the solute from the solution, Decreasing pressure of the gas by increasing it's volume.
• Enthalpy: given off in the reaction as heat energy into the surroundings
• Endothermic reactions absorb heat energy from their surroundings.
• Delta (triangle) H = H (products) - H (reactants)
• When a reaction has reached equilibrium: there will be no change in the net amount of products or reactants of the reaction.
• A catalyst increases the rate of a chemical reaction without being used up in the process. It lowers the activation energy required for the reaction to occur.
• Acid: a substance which in solution produces hydrogen ions or otherwise known as hydronium ions. Also known as the "proton donors" to bases
• Neutralisation: is the reaction between an acid and a base, forming salt and water only.
• pH scale: measures how acidic/basic a solution is on a logarithmic scale ranging from 0-14. The pH value of pure water is 7.
• Base: A substance that can accept protons/hydrogen ions and contains OH- ions (hydroxide ions)
• pH scale measures how acidic or alkaline a solution is on a logarithmic scale ranging from 0-14. The pH value of pure water at room temperature is 7.
• Strong Acid: Have more H+ ions which have dissociated in water
• Weak Acid: Have less H+ ions that have dissociated in water
• Strength of Acid or Base: Refers to the level of dissociation in solution
• Acid + Base = Salt + Water (Neutralisation)
• pH of solution = negative of the logarithm (base ten) of the hydrogen ion concentration
• The lower the pH, the higher the [H+]
• The higher the pH, the lower the [H+]
• pH = -log 10 [H30+] (brackets refer to concentration of)
• Kw: The ionic product constant of water
• Concentrated Acids: Total concentration of H+ ions in the solution is greater than the concentration of OH- ions
• Dilute acids: Total concentration of H+ ions in the solution is less than the concentration of OH- ions
• Bronsted Lowry's Law: Acid is a "proton donor" whilst a base is a "proton acceptor"
• Acid gives up proton to form a conjugate base
• Monoprotic Acid = Gives up one proton
• Diprotic Acid = Gives up two protons