Physical Chemistry Flashcards
Enthalpy:
Enthalpy, H, is the thermal energy stored in a chemical system
The law of conservation of energy:
The law of conservation of energy states that energy cannot be created or destroyed, only moved from one place to another
Enthalpy change:
Exothermic reactions:
Endothermic reactions:
Activation energy:
Activation energy is the minimum energy required to start a reaction by breaking bonds in the reactants
Standard conditions:
Enthalpy change of reaction:
Enthalpy change of formation:
Enthalpy change of combustion:
Enthalpy change of neutralisation:
Calorimetry:
Calorimetry is the quantitative study of energy in a chemical reaction
Calorimetry equation:
Specific heat capacity:
The specific heat capacity is the energy required to raise the temperature of 1 gram of a substance by 1 Kelvin
Experiment: A coffee cup calorimeter
The reaction vessel can be the insulated coffee cup held in a beaker. You could use a polystyrene lid with a hole for the thermometer to reduce heat loss to the surroundings (minimise the biggest experimental error)
• Add a measured mass of the first liquid reactant. Take the temperature every minute until it is stable. This usually takes around 4 minutes
• At 5 minutes, add the second reactant. Do not take or record the temperature for the fifth minute
• Monitor the temperature of the reaction mixture every minute for a further 5 minutes
• Plot a graph to infer the maximum temperature change generated by the reaction (extrapolation)
Experiment: Copper calorimetry
When a fuel is combusted, the heat energy can be used to increase the temperature of a known mass of water
• Measure the starting mass of the fuel
• Add a known mass of water to a copper calorimeter
• Mount the copper calorimeter over the fuel and take the starting temperature of the water
• Combust the fuel for a few minutes and take the final temperature of the water
• Take the mass of the unused fuel and calculate the mass of the fuel burnt
A bomb calorimeter:
A bomb calorimeter is a sophisticated piece of equipment that minimises heat loss as much as possible. It uses pure oxygen, to ensure complete combustion is achieved
Bond enthalpy:
- Bond enthalpies tell you how much energy is needed to break each different bond
- When calculating bond enthalpies, you must include the overall enthalpy change of the chemical equation in your calculation
Average bond enthalpy:
The mean energy needed for 1 mole of a given type of gaseous bonds to undergo homolytic fission
Breaking and making bonds:
Hess’ law:
Hess’ law states that the enthalpy change on a chemical reaction is independent of the route it takes
Enthalpy cycle:
An enthalpy cycle is a pictorial representation showing alternative routes between reactants and products
Hess’ cycle - enthalpy change of reaction:
Hess’ cycle - enthalpy change of formation:
Rate of reaction:
The rate of reaction is defined as the change in concentration of a reactant or a product in a given time. Usually:
- Concentrations of reactants are highest at the start of a reaction - this is when the rate is at its fastest (the initial rate)
- The rate slows down as reactant concentrations decrease (reactants get used up)
- When any reactant has a concentration of zero (it is used up) the reaction stops, and the rate of reaction is zero
Factors affecting the rate of a chemical reaction:
- Temperature
- Pressure (for gaseous reactants only)
- Concentration
- Surface area
- Adding a catalyst
The effect of temperature on reaction rate:
An increased temperature results in an increased rate of reaction as the reacting particles gain kinetic energy. More of the particles now have more than or sufficient activation energy. After gaining kinetic energy, the particles start to move more, and move more quickly and vigorously. There are more collisions per second, and so more successful collisions per second, meaning that there is an increased rate of reaction
The effect of pressure on reaction rate:
When the pressure of a gas is increased, the molecules are pushed closer together. The same number of molecules occupies a smaller volume. As the molecules are closer together, they will collide more times per second, and so there will be more successful collisions per second, as more of the collisions are likely to occur with sufficient energy to overcome the activation energy. Therefore, the rate of reaction increases
What is a catalyst?
A catalyst is a substance that increases the rate of a reaction without being used up during the process. Catalyst provide an alternative pathway of lower activation energy
Homogeneous catalysts:
If a catalyst for a reaction is in the same phase (state) as the reactants, then it is a homogeneous catalyst. Enzymes within bodily fluids such as blood or saliva are examples of homogeneous catalysts
Heterogeneous catalysts:
If a catalyst for a reaction is in a different phase from the reactants, it is a heterogeneous catalyst. Phases do not just include states, for example if all of the chemicals are immiscible liquids, the catalyst could still be heterogeneous, as it could be in a different liquid layer from the reactants. An example of a heterogeneous catalyst is a catalytic converter in a car - solid metal (platinum, rhodium or palladium) acting as catalyst to turn harmful gaseous exhaust fumes into less harmful substances
The economic importance of catalysts:
Catalysts help to lower the energy demands of processes. By doing so, they reduce costs and also help the environment. Less fossil fuel needs to be burned to generate the required energy, and this also means lower carbon dioxide emissions
- An important industrial application of catalysts are within catalytic converters. Catalytic converters play an important role in improving our air quality by reducing toxic emissions from vehicles and preventing photochemical smog
Haber Process: Economic importance
- The Haber process for the production of ammonia is of great economic importance. The ammonia is used as the basis for fertiliser manufacture, improving crop yields to feed the ever-increasing world population
- A lot of energy is required to break the triple bond in nitrogen, contributing to a high activation energy. Iron is used to catalyse this reaction, weakening nitrogen’s triple bond, and lowering the activation energy and therefore the costs - both financial and environmental
What is the Boltzmann distribution?
The Boltzmann distribution is the distribution of energies of molecules at a particular temperature, often shown as a graph
Important features of the Boltzmann distribution:
- The area under the curve is equal to the total number of molecules in the sample. The area does not change with conditions
- There are no molecules in the system with 0 energy - the curve starts at the origin
- There is not maximum energy for a molecule - the curve gets close to (the bottom square) but does not touch or cross the Energy (x) axis
- Only the molecules with an energy greater than the activation energy, Ea, are able to react
The effect of temperature on the Boltzmann distribution:
The Boltzmann distribution flattens and shifts to the right. The number of molecules in the system does not change, so the area under the curve remains the same, The end of the curve must finish in the same square (but above) the original temperature curve
The effect of catalysts on reaction rate and Boltzmann distributions:
- Catalysts lower the activation energy of reactions. They do not change the distribution of energy within molecules. However, by lowering the activation energy, more particles will automatically be above the activation energy barrier
- On collision, more molecules in the system will overcome the new lower activation energy of the reaction. There will be more successful collisions per unit time, and the rate of reaction will increase
When is a system in dynamic equilibrium?
A chemical system is in dynamic equilibrium when:
- The concentrations of the reactants and products remain constant
- The rate of the forward reaction is the same as the rate of the reverse (backward) reaction
The system must remain isolated
Factors affecting the position of equilibrium:
- Concentrations of reactants or products
- Pressure in reactions involving gases
- Temperature
The effect of pressure on equilibrium (Le Chatelier’s principle):
Increasing the total pressure of the system, causes the equilibrium position to shift to the side with fewer moles of gas, as this will decrease the pressure
The effect of concentration on equilibrium (Le Chatelier’s principle):
Increasing the concentration of a reactant favours the forwards reaction, and causes the equilibrium position to shift to the right , in order to form more products
The effect of temperature on equilibrium (Le Chatelier’s principle):
Increasing the temperature favours the endothermic reaction, and so the equilibrium positions shifts in the direction of the endothermic reaction
The effect of a catalyst on equilibrium:
A catalyst does not alter the position of equilibrium or the composition of an equilibrium system:
- A catalyst speeds up the rate of the forward reaction and reverse reactions equally
- A catalyst increases the rate at which equilibrium is established but does not affect the position of the equilibrium
Equilibrium vs yield:
Many important chemical processes exist as equilibrium systems. Examples include:
- the preparation of ammonia from nitrogen and hydrogen in the Haber process
- the conversion of sulphur dioxide into sulphur trioxide in the Contact process
In industry, chemists strive to achieve the highest possible yield of a desired product. However this has to be balanced against the optimum position of equilibrium, wichita allows industrial processes to be as cheap and energy-efficient as possible - to create the highest profit
The Haber process:
The raw materials, nitrogen and hydrogen, must be readily available:
- nitrogen is obtained from the air by fractional distillation
- hydrogen is prepared by reacting together methane (from natural gas) and water
Ammonia is produced by the forward reaction in this equilibrium. The optimum conditions are high pressure and low temperature. This is because:
- the products have less moles than the reactants, so is favoured by high pressures
- the forward reaction is exothermic (ΔH is negative), so is favoured by using a low temperature
However, there are drawbacks to using these theoretical conditions:
- although a low temperature should produce a high equilibrium yield, the reaction would take place at a very low rate. At low temperatures, comparatively few N2 and H2 molecules have enough energy to overcome the required activation energy
- a high pressure increases the concentration of the gases, increasing the reaction rate. So, a high pressure should produce both a high equilibrium yield and a high rate. However, large quantities of energy are required to compress gases, adding significantly to the running costs. There are also safety implications - any failure in the systems could potentially allow chemicals to leak into the environment, endangering those working on site
The modern ammonia plant:
A modern ammonia plant needs to produce a sufficient yield of ammonia at a reasonable cost and in as short a time as possible. In practice, a compromise is made between yield and rate:
- Temperature - this must be high enough to allow the reaction to proceed at a realistic rate, whilst still producing an acceptable equilibrium yield. A temperature of 400-500°C is typically used
- Pressure - a high pressure must be used, but it must not be so high that the workforce is put in danger or the environment threatened. A pressure of 200 atmospheres is typically used
- Catalyst - an iron catalyst is added to speed up the rate of reaction, allowing the equilibrium to be established faster, and lower temperatures to be used. Less energy is needed to generate heat, reducing costs
The compromise conditions used covert only 15% of nitrogen and hydrogen into ammonia. The ammonia produced is liquefied and removed. Unreacted nitrogen and hydrogen gases are then passed through the reactor again. Eventually, virtually all the nitrogen and hydrogen will have converted into ammonia
The equilibrium constant, Kc, and the equilibrium law
- Chemists consider the position of an equilibrium using the equilibrium constant, Kc
- Kc = products/reactants
The significance of the Kc value
- If Kc = 1, then the equilibrium position is halfway between the reactants and products. This would mean that the ratio of the product concentration to the reactant concentration is 1:1
- If Kc>1, the reaction favours the products (equilibrium lies to the right)
- If Kc<1, the reaction favours the reactants (equilibrium lies to the left)
