F322: Enthalphy Changes, Equilibria And Rates Flashcards
Bond enthalpy
The energy required to break 1 mole of a given type of bond in the gaseous state.
Always +ve numbers as bond breaking is a endothermic process.
Standard enthalpy change of reaction
The enthalpy change that accompanies a reaction. In the molar quantities expressed in a chemical equation under standard conditions (1atm, 25 degrees, conc of 1mol/dm3)
Equation to calculate enthalpy chane of reaction
Sum of bond energies of reactants - sum of bond energies of products
Exothermic reaction
A reaction that gives out energy to its surroundings. -ve enthalpy change of reaction. More energy released when bonds are formed in products, than is required to break the bonds in the reactants.
Endothermic
A reaction that takes energy to its surroundings. +ve enthalpy change of reaction. Less energy released when bonds are formed in products, than is required to break the bonds in the reactants.
Hess’s law
The total energy change of a reaction is the same, regardless of the route taken, provided the initial and final conditions are the same.
Standard conditions
25 degrees C (298K) and 1atm (100KPa)
Standard enthalpy change of formation
The enthalpy change when 1 mole of a compound is formed from it’s elements under standard conditions.
Standard enthalpy change of combustion
The enthalpy change when 1 mole of a substance completely combusts under standard conditions.
Collision theory
Particles must collide with a minimum about of energy for a reaction to occur
Activation energy
Minimum amount of energy required in a collision to break bonds in order for a reaction to occur.
Effect of increased concentration on rate of reaction
More particles per unit volume, increased frequency of collisions, increased rate
Effect of increased pressure (gas) on rate of reaction
Same number of particles in a smaller volume, increased frequency of collisions, increased rate
Effect of increasing temperature on rate
Particles now have more energy, move around faster (greater KE), collide more frequently, more molecules have >Ea so more collisions are successful
Catalyst
Speeds up the rate of reaction without being consumed in the overall reaction. It provides an alternative pathway/route with a lower Ea, so more molecules have >Ea.
Industrial importance of catalysts (examples)
Iron in ammonia production
Platinum/Palladium/Rhodium in catalytic converters
Industrial benefits of using catalysts
1) Lower temperatures can be used, reducing production costs and CO2 emissions from burning hydrocarbons to produce energy.
2) Enable different reactions to be used with better atom economy to reduce waste.
3) some catalysts such as enzymes function effectively at room temperature.
Boltzmann distribution
Shows the distribution of energies for a sample of molecules
Dynamic equilibrium
Achieved when the rate of the forward reaction is equal to the rate of the backward reaction. Requires a closed system, and no overall change in macroscopic properties.
Closed system
A system where no reactants or products can escape (energy can exchange with surroundings)
Equilibrium yield
The amount of product in the equilibrium mixture
Le Chatelier’s Principle
A change in the conditions to a system at equilibrium will cause a shift in the position of the equilibrium in the direction that MINIMISES the effect of the change.
Effect of increasing on the equilibrium position
Equilibrium shifts in the endothermic direction
Effect of increasing pressure on the equilibrium position
Equilibrium shifts to the side with FEWER moles of gas
Effect of changing concentration on the equilibrium position
Shifts right if: products removed, reactants added
Shifts left if: reactants removed, products added
Effect of adding a catalyst on the equilibrium position
No effect on equilibrium position
As it speeds up the rate of forward and backward reaction by the same amount.
Note: particularly useful in reactions where increasing the temp lowers the yield.
