3.2 Physical chemistry Flashcards
enthalpy definition
measure of heat energy stored in a system (H)
enthalpy change definition
(ΔH) difference between the enthalpy of reactants and enthalpy of products
heat definition
process whereby thermal energy (J) is transferred from hotter object to a cooler object
temperature
direction of energy transfer is determined by temperature of objects (measured in K, °C, °F)
system definition
substances directly involved in a chemical reaction
surroundings definition
apparatus, laboratory, anything not part of the system
universe definition
system+surroundings
law of conservation of energy
energy cannot be created or destroyed
exothermic meaning
energy released to and gained by surroundings, temperature increased
ΔH is negative
more energy released from making bonds than energy absorbed from breaking bonds
endothermic meaning
system takes in energy from surrounding, temperature decrease
ΔH is positive
more energy absorbed from breaking bonds than energy released from making bonds
enthalpy of reaction definition
overall energy change in a reaction
activation energy definition
minimum energy required for a reaction to take place
standard enthalpy of reaction definition
overall energy change in a reaction under standard conditions
100 kPa, 298 K / 25°C, conc. of 1.0 mol dm^-3 (for reactions with aqueous solutions)
standard conditions
100kPa,
298K/25°C
1 mol dm^-3 (for solution)
standard state definition
physical state of substance under standard conditions
standard enthalpy change of combustion definition
energy change when one mole of substance is combusted in excess oxygen
standard enthalpy change of formation definition
overall energy change of 1 mol of a compound from its elements
standard enthalpy change of neutralisation definition
overall energy change when 1 mole of water is produced in a reaction with an acid and base
enthalpy change of reaction formula
ΔrH = sum of bond enthalpies of broken bonds - sum of bond enthalpies of bonds made
average bond enthalpy definition
breaking of 1 mol of bonds in gaseous molecules
why average bond enthalpy differs from actual bond enthalpy (and less accurate than formation enthalpies)
different bond enthalpies are measured in different conditions (no universal, unchanging standard saying which molecules used to determine each bond)
Hess law definition
if reaction takes place by more than one route from initial to final conditions, total enthalpies change is the same for each route
simple collision theory
reaction takes place when molecules collide with sufficient energy and correct orientation
factors affecting rate of reaction
concentration/pressure increases rate of collision (particles closer together, more likely to collide)
temperature increases heat energy (more kinetic energy so faster molecules so higher rate of collision, also more molecules with heat energy equal or more than activation energy)
surface area increases likelihood of collision between molecules
Boltzmann distribution curve features
area under graph = number of molecules
shaded area = number of molecules able to react
Ea line = activation energy
peak of graph = most likely energy of a molecule
how catalyst affects Boltzmann distribution curve
moves Ea line to the left (more shaded area)
how temperature affects Boltzmann distribution curve
lowers peak of the graph then right
how dynamic equilibrium exists
rate of forward and reverse reactions are equal
concentration of reactants and products do not change
Le Chatelier’s principle
if a system at equilibrium exposed to a stress (e.g. concentration, pressure, temperature change), system will shift its equilibrium to relieve that stress
how concentration affects equilibrium
more concentration of reactants, faster rate of forward reaction
equilibrium shifted right
more concentration of products, faster rate of reverse reaction
how temperature affects equilibrium
when temperature increases
if reaction is exothermic, rate of reverse reaction increases (equilibrium shifts left)
if reaction endothermic, rate of forward reaction increases (equilibrium shifts left)
and vice versa
how changing volume/pressure affects equilibrium
lower pressure (more volume) = rate of reaction making more gas molecules increasing higher pressure (less volume) = rate of reaction making less gas molecules increasing
how catalyst affects equilibrium
increases rate of forward and reverse reaction
no effect on position of equilibrium
reduces time taken to reach equilibrium
how to determine position of equilibrium
find out equilibrium constant
if Kc > 1, position of equilibrium closer to the right (products)
if Kc < 1, position of equilibrium closer to the left (reactants)
rate of chemical reaction definition
how fast a reaction is taking place
how fast a reactant is used up
how fast a product is made
what concentration time graphs look like
steepest at start (rate of reaction fastest in the beginning) gradient lowers the longer the reaction is taking place (reactants concentration lower, less likely to collide and react) eventually plateaus (no more reaction taking place, one of reactants completely used up)
how to measure rate of reaction if gas is produced
monitor volume of gas at regular intervals (with gas collection)
monitor loss of mass of reactants with mass balance
monitor production of gas with gas collection method
place measuring cylinder (held by clamp) submerged under trough of water
measure initial volume in measuring cylinder
place one reactant in conical flask
place bung on top of conical flask
quickly add other reagents to conical flask and replace bung and instantly start time
record volume of gas at regular intervals until reaction is complete (no more gas produced)
plot graph
gradient (found using tangent) = rate of reaction at point of tangent
monitor loss of mass of reactants using mass balance method
record mass of conical flask and reactants initially and regular intervals
reaction complete when no more gas produced
plot graph
catalyst definition
substance that changes rate of chemical reaction without undergoing any permanent change itself
how catalyst increases rate of reaction
provides alternate reaction pathway of lower activation energy
catalyst facts
not used up in chemical reaction
may react with reactant to form intermediate or provide surface where reaction can take place
catalyst is regenerated at end of reaction
homogenous catalyst definition
same physical state as reactants
reacts with reactant to form intermediate that breaks down to give products and catalyst
heterogenous catalyst definition
different physical state from reactants
usually solids in contact with gaseous reactants or reactants in solution
reactant molecules adsorbed onto surface of catalyst, reaction takes place
product molecules leave surface of catalyst via desorption
adsorption definition
weakly bonding to surface
desorption definition
molecules leaving surface of another substance
autocatalysis definition
reaction products acts as catalyst for that reaction
importance of catalysts
increases rate of industrial chemical reactions by lowering activation energy
also reduces temperature needed for processes and energy requirements (less fossil fuels and electricity used)
cuts cost
more sustainable (higher atom economies, fewer pollutants, cuts CO emissions)
