Topic 5: Energetics & Thermochemistry Flashcards
exothermic reaction
- heat energy is released
- energy released in bond formation on product side is greater than energy consumed in bond breaking on reactant side
enthalpy change of exothermic reactions
ΔH = - x
endothermic rxn
- heat energy is absorbed
- energy released in bond formation on product side is less than energy consumed in bond breaking on reactant side
enthalpy change of endothermic reactions
ΔH = + x
system
- a specified part of the universe
- under observation/where a chem rxn is taking place
surroundings
- the remaining portion of the universe
- NOT part of the system
enthalpy
heat content of the system
symbol: H
enthalpy change
- heat absorbed/evolved during a rxn
- measured at a constant temp/pressure
symbol: ΔH
unit: kJ mol^-1
enthalpy change equation
ΔH = Hp - Hr ΔH = enthalpy change Hp = enthalpy of products Hr = enthalpy of reactants
absolute zero
-273 degrees Celsius
the temp at which all particle movements cease completely
heat
- form of energy
- measure of total energy in a given amount of substance
- depends on the amount of substance present
temperature
- measures the ‘hotness’ of a substance
- avg kinetic energy of the substance
- independent of the amount of substance present
specific heat capacity
heat required to raise the temp of 1g of a substance by 1°C/K unit: J K^-1 g^-1 j = joules k = kelvin g = grams
heat capacity
heat needed to increase the temp of the object by 1°C/K
standard enthalpy change
enthalpy change when 1 mol of the gaseous bond is broken or formed
calorimetry
technique of measuring enthalpy change
what is a calorimeter?
- a well-insulated container (e.g. polystyrene cup)
- in which temp change of a liquid is measured
- before and after the change
assumptions made when making heat calculations
- no heat transfer between the soln, the thermometer, the surrounding air, and the calorimeter itself
- the solution’s density and specific heat capacity are equivalent to water’s
- rxn occured fast enough for max temp to be achieved before cooling begins
problems with calorimetry
- not having the desired rxn occur (e.g. incomplete combustion)
- loss of heat to surroundings in exothermic reactions
- absorption of heat from surroundings in endothermic reactions
- using incorrect specific heat capacity value in enthalpy calculations
what is a cooling graph?
- for slow rxns (e.g. metal ion displacement), the results will be less accurate
- due to heat loss over time
- this can be compensated by plotting a temp-time graph (cooling graph) to extrapolate backwards
how to draw a cooling graph
- draw a line of extrapolation backwards from cooling curve
- draw a vertical line where the reactants were mixed (i.e. when the curve begins to rise)
- the y-coordinate of the point of intersection is the temp that would’ve been reached if no heat was lost to surroundings
Hess’ Law
- the total enthalpy change for a chem rxn doesn’t depend on the pathway it takes
- only considers initial and final states
1st law of thermodynamics
AKA: law of conservation of energy
- energy can neither be created nor destroyed
- it can only be converted from one form to another
enthalpy of formation
- denoted by ΔHf
- enthalpy change when 1 mole of a substance is formed from its elements
- all substances being in standard states
- enthalpy of formation of every element in its standard state is assumed to be 0!
what happens to enthalpy change when you reverse a rxn
signs are reversed
e.g. negative value turns positive and vice versa
Energy
measure of the ability to do work
conditions of measuring standard enthalpy change
- pressure: 100kPa
- conc.: 1 mol/dm3 for all solutions
- temp.: 298K (usually)
- all substances in standard state
how to calculate reaction enthalpies from temp changes
ΔH (reaction) = - ΔH (water) = - m (H2O) x c(H2O) x ΔT(H2O)
- it’s negative bc water gains heat (positive enthalpy change) from the reaction
- so if water’s ΔH is positive then the reaction’s ΔH is negative
how to calculate enthalpy change using Hess’ Law
ΔH1 = ΔH2 + ΔH3
ΔHf (products) = ΣΔH (reaction) + ΣΔHf (reactants)
ΔH1 = enthalpy change of formation (of products) ΔH2 = enthalpy change of formation (of reactants) ΔH3 = enthalpy change of reaction
average bond enthalpy
- energy needed to break one mole of a bond in a gaseous molecule
- can be used to calculate enthalpy change
limitations of using bond enthalpies
- can only be used if all reactants and products are in gaseous state; in other states, more heat would be evolved
- bond enthalpies are averaged values that are obtained considering a number of compounds containing the bond, but in reality the bond energy varies between compounds
formula for specific heat capacity
q = mcΔT
q: heat change
m: mass
c: specific heat capacity
ΔT: change in temperature
formula for heat capacity
C = q/ΔT
C: heat capacity
q: heat change
ΔT: change in temperature
assumptions made in enthalpy change calculations for solutions
ΔH(system) = ΔH(water) + ΔH(reaction) = 0 ΔH(reaction) = - ΔH(water)
enthalpy change: does change in temperature differ depending on volume of reactants?
volume halved = enthalpy halved = heat required also halved = change in temperature will be constant!
lattice enthalpy
AKA ΔH(lat)
- endothermic: enthalpy change when 1 mol of gaseous ions form from 1 mol of a solid crystal
- exothermic: enthalpy change when 1 mol of solid crystal forms from 1 mol of gaseous ions
- sign (+/-) of ΔH(lat) denotes whether it’s endothermic/exothermic
factors affecting lattice enthalpy/enthalpy of hydration
↓ size + ↑ charge = ↑ ΔH(lat)
↓ size + ↑ positivity of charge = ↓ ΔH(hyd) – because the enthalpy of hydration is always negative
enthalpy change of hydration
- enthalpy change occurring when 1 mol of gaseous ions is dissolved to form an infinitely dilute solution of 1 mol of aqueous ions
- always exothermic so ΔH(hyd) is always negative
entropy
- i.e. randomness/disorder
- denoted by S
- defined as the distribution of available energy among particles
- higher energy distribution (i.e. diffused state) = higher entropy
- all entropy values are positive except for perfectly ordered solids at absolute zero (S = 0)
how to predict ΔS with relation to state
- conversions from s –> l/g have positive ΔS
- vice versa for g –> l/s
formula to calculate ΔS
ΣΔS(system) = ΣΔS(products) - ΣΔS(reactants)
why are exothermic reactions more common than endothermic reactions?
- because ΔS(surroundings) ∝ -ΔH(system)
- exothermic reactions cause an increase in the entropy of the surroundings
relationship between ΔS and temperature
ΔS(surroundings) ∝ 1/T
- i.e. entropy of surroundings is inversely proportional to absolute temperature
formulae for ΔS
ΔS(total)* = ΔS(system) - ΔS(surroundings)
*total denotes total entropy of the universe.
ΔS(surroundings) = - ΔH(system) / T
how to deduce whether a reaction is feasible
as long as ΔS(total) > 0, a reaction will occur spontaneously
Gibbs free energy
- criterion of feasibility of a reaction occurring
- feasible reactions = spontaneous
- ΔG is always negative for spontaneous processes and positive in the reverse case
- ΔG is NOT the same as ΔS
formula for Gibbs free energy
ΔG(system) = -TΔS(total)
ΔG(system) = ΔH(system) - TΔS(system)
T must be in Kelvin (in standard conditions, T is 298K)
how can ΔG(system) predict feasibility?
- as ΔS is inversely proportional to T, the T value adjusts the importance of ΔS(system) in the gibbs free energy equation
- at low temps, ΔS is high but T is low so overall TΔS(system) ≈ 0. Thus ΔG(system) = ΔH(system)
- at high temps ΔH is negligible so ΔG(system) = TΔS(system)
formula for ΔG(reaction)
ΣΔG(reaction) = ΣΔG(products) - ΣΔG(reactants)
Relationship between ΔG(reaction) and equilibrium
as ΔG(reaction) becomes more negative (i.e. as reaction feasibility increases), the equilibrium position shifts more to the right
linking ΔS to ΔG
ΔS(total) = ΔS(system) - ΔS(surroundings)
Knowing ΔS(surroundings) = - ΔH(system) / T, we can substitute that:
ΔS(total) = ΔS(system) - ΔH(system) / T
-TΔS(total) = ΔH - TΔS(system)
Knowing ΔG(system) = -TΔS(total), we can substitute that:
ΔG(system) = ΔH(system) - TΔS(system)
enthalpy of atomization
ΔH(at)*
- standard enthalpy change when 1 mol of gaseous atoms are formed from the element in its standard state under standard conditions
- for diatomic molecules the ΔH(at) value is half that of the bond dissociation enthalpy
*standard enthalpies like ΔH(atom) are always written with a theta in superscript
difference between bond dissociation enthalpy and bond enthalpy
- bond dissociation enthalpy refers to the enthalpy of a specific bond in a specific compound
- bond enthalpy refers to the average enthalpy of bond, taking into account different compounds
- a specific bond’s energy differs somewhat between compounds, so bond dissociation enthalpy is more specific/accurate
enthalpy of solution
ΔH(sol)
- enthalpy change when 1 mol of ionic substance dissolves in water to form a solution of infinite dilution
Hess’ Law for the solubility of salts
ΔH(sol) = ΔH(lat) + sum of ΔH(hyd) of ALL constituents
NOTE: ΔH(sol) and ΔH(lat) are values of compounds while ΔH(hyd) is individual values of ions
why can’t H2O be added to SO3?
- taking SO3’s non-ionic nature into account, its ΔH(lat) value will be small
- considering ΔH(sol) = ΔH(lat) + ΔH(hyd of ALL constituents), remember the fact that ΔH(hyd) is always negative (exothermic)
- as ΔH(lat) will be small, the overall ΔH(sol) value for H2O + SO3 will be highly exothermic
- the evolved heat could be high enough to cause the product H2SO4 to boil
factors affecting ΔH(hyd)
- ionic radius
- charge of ion
factors affecting ΔH(hyd): ionic radius
smaller ionic radius = higher ΔH(hyd)
factors affecting ΔH(hyd): charge of ion
more positive charge = higher ΔH(hyd)
- due to increased attraction between the positive ion and the partially negative O atoms in water
