Rates of reaction

• The rate of the reaction is how fast the reactants change into products. A slow reaction is the rusting of iron. A fast reaction is burning.
• Interpreting graphs - Steeper the line on the graph, the faster the rate of reaction. The quickest reactions are the steepest and become flat in the least amount of time. When the line becomes less steep the reactants are used, the reaction has finished.
• Collision theory - For a chemical reaction to happen particles must collide with each other. Reactant molecules must collide with enough energy for the collision to be successful. The more collisions the faster the reaction. Activation energy is the minimum amount of energy needed for a collision to be successful. This energy is needed so particles can break bonds and start the reaction.
• 4 factors which increase the rate of reaction as they increase the number of successful collisions: temperature, concentration of a solution or the pressure of gas, surface area, presence of a catalyst.
• Increasing the temperature - When the temperature increases, particles move faster and collide more frequently. Also the faster they move, the more energy they have, so more collisions will have enough energy to make the reaction happen.
• Increasing concentration or pressure - If a solution is made more concentrated, it means more particles are in the same volume of water (or solvent). When the pressure of a gas is increased, it means the same number of particles occupies a smaller space. This makes collisions between the reactant particles more frequent.
• Increasing surface area - If a reactant is solid, then breaking it up into smaller pieces increases its surface area to volume ratio. This means that for the same volume of the solid, the particles around it will have more area to work on, so there will be collisions more frequently.
• Adding a catalyst - A substance that speeds up a reaction, without being used up in the reaction itself, this means it's not part of the overall reaction equation. Different catalysts are needed for different reactions. They work by decreasing the activation energy by providing an alternate reaction pathway with a lower activation energy.
• Adding a catalyst:

  A) Activation energy without catalyst

  B) Activation energy with catalyst

  
• Rates of reaction can be observed either by how quickly the reactants are used up or how quickly products are formed.
• Rate of reaction = amount of reactant used or amount of product formed / time
• Measuring rates of reaction- Precipitation and colour change: 1) You can record the visual change in a reaction if the initial solution is transparent and the product is a precipitate which clouds the solution (2) You can observe a mark through the solution and measure how long it takes for it to disappear, the faster it disappears, the quicker the reaction (3) If the reactants are coloured and the products are colourless (or vice versa), you can time how long it takes for the solution to lose or gain its colour (4) Results are subjective
• Measuring rates of reaction - Change in mass (gas given off usually): 1) Measuring the speed of a reaction that produces a gas can be carried out using a mass balance. (2) As the gas is released, the mass disappearing is measured on the balance. (3) The quicker the reading on the balance drops, the faster the reaction (4) If you take measurements at regular intervals, you can plot a rate of reaction graph and find the rate easily (5) Most accurate as the mass balance is very accurate.
• Measuring rates of reaction - The volume of gas given off: 1) Uses a gas syringe to measure the volume of gas given off (2) The more gas given off during a given time interval, the faster the reaction (3) Gas syringes measure to the nearest cm^3, so are quite accurate (4) Take measurements at regular intervals and plot a rate of reaction graph.
• RQ, Magnesium + HCL react to produce Hydrogen gas: 1) Start by adding a set volume of dilute hydrochloric acid to a conical flask (2) Add some magnesium ribbon to the acid and quickly attach an empty gas syringe to the flash (3) Start the stopwatch. Take readings of the volume of gas in the syringe at regular intervals. (4) Plot the results in a table and then on a graph - As the concentration of acid increases, the rate of reaction will increase
• RQ, Sodium thiosulfate + HCL produce a cloudy precipitate: 1) Both chemicals are clear. They react to form a yellow precipitate of sulfur. (2) Add a set volume of dilute sodium thiosulfate to conical flask (3) Place the flask on a piece of paper with a black cross on it (4) Add some dilute HCL to the flask and start the stopwatch (5) Time how long it takes for the black cross to disappear through the cloudy sulfur - As the concentration of acid increases, the solution goes cloudy more quickly
• Finding the mean rate of reaction from a graph - mean rate of reaction = change in y /change in x

  
• Find the rate of reaction at a particular point - 1)Position the ruler on the graph at the point where you want to know the rate of (2) Adjust the ruler until the space between the ruler and the curve is equal on both sides of the point (3) Draw a ling along the ruler to make the tangent. Extend the line right across the graph (4) Pick two points on the line that are easy to read. Use them to calculate the gradient of the tangent in order to find the rate.

  
• Reversible reactions - A + B --> <-- C + D, the products can react to form the reactants. The forward reaction goes right and the backward reaction goes left.
• Dynamic equilibrium - As the reactants react, their concentrations fall so the forward reaction will slow down. But as more and more products are made their concentrations rise and the backward reaction will speed up. If the reaction is in a closed system, after a while the reaction will reach equilibrium. This means that the forward and backward reactions are still happening and they have the same rate of reaction so the concentration of all reacting substances remain constant.
• Equilibrium - Doesn't mean the amount of reactions and products are equal. If the equilibrium lies to the right, the concentration of products is greater than that of the reactants. If the equilibrium lies to the left, the concentration of reactants is greater than that of the products. The position of equilibrium depends on temperature, pressure and concentration of the reactants and products.
• Endothermic - A reaction that takes in energy from the surroundings.
• Exothermic - Energy is released to the surroundings.
• In reversible reactions if the reaction is endothermic in one direction, it will be exothermic in the other.
• Le Chatelier's Principle - If you change the conditions of a reversible reaction at equilibrium, the system will try to counteract that change.
• Le Chatelier's Principle - Temperature: All reversible reactions have an exothermic and endothermic direction. If the temperature is increased, the equilibrium position moves in the endothermic direction.
• Le Chatelier's Principle - Pressure: Changing the pressure only affects an equilibrium involving gases. If you increase the pressure in a reaction, the equilibrium position moves in the direction of the fewest molecules of gas to reduce the pressure.
• Le Chatelier's Principle - If you increase the concentration of the reactants the system tries to decrease it by making more products. If you decrease the concentration of products the system tries to increase it again by reducing the amount of reactants.