Energetics Flashcards
All chemical reactions involve…
… changes in energy, usually heat energy
Enthalpy change
- energy change during any change in a system under constant pressure
- ΔH
- dépendent on amount of material involved, so the values are standardised
When a bond is formed…
… one electron from each atom entered a new electron cloud of lower energy
Bond enthalpy
- heat energy obtained when a bond is formed
- heat energy taken in when a bond is broken
- 1dp
- approximation - lacks accuracy
Having a low bond enthalpy
- allows for low Ea; faster rate
* broken first in a series of steps
Why do bond enthalpies lack accuracy?
The energy required to break each of the bonds in a molecule is different
Mean bond enthalpy
- enthalpy needed to break the covalent bond into gaseous atoms, averaged over different molecules
- positive, because is energy is required to break a bond
- must start and end in gaseous state
- other states are much less accurate
E.g. 1/4CH4 -> C + H
Why is the real bond enthalpy différent to the average bond enthalpy?
The environment of the bonds differ
Exothermic
- -ΔH
- à réaction in which energy is given out
- more energy is given out from bonds being formed than energy is taken in from bonds being broken
- temperature increase in surroundings
- majority of reactions
Endothermic
- +ΔΗ
- a reaction in which heat energy is taken in
- breaking of bonds requires more energy than is given out by the formation of new bonds
- temperature decrease in surroundings
Standard enthalpy of ΔH
- molar quantities
- 298K
- 1 atm pressure
- normal state
- solutions at 1moldm^-3
Elements in their common state…
… are assigned an energy value of zero
What is important in enthalpy change values?
- state signs
- allotropic form
Because values are specific to quantity
Standard enthalpy of formation
- ΔHf
- enthalpy change that takes place when one mole of substance is formed from its elements in their standard states at a temperature of 298K and a pressure of 100kPa
- an element = 0kJmol^-1
Standard enthalpy of combustion
- ΔHc
- enthalpy change when one mole of substance is burned completely in oxygen at a temperature of 298K and 100kPa, in standard states
- incomplète combustion will lead to carbon (soot), CO and water; less exothermic
Standard enthalpy of neutralisation
- ΔHneut
- the enthalpy change when one mole of water is formed in a neutralisation reaction between an acid and an alkali at a temperature of 298K, and a pressure of 100kPa, in standard states
- always exothermic
Standard enthalpy of reaction
- ΔH
- then enthalpy change when the number of moles of substances in the equation react at a temperature of 298K and a pressure of 100kPa, in standard states
Calorimetry
Measures the energy given out in a reaction and absorbed by water, either as a solvent for the chemicals or in a metal calorimeter
Measuring calorimetry
1) mixing substances in an insulated container
2) using a calorimeter in a combustion experiment
Enthalpy in solution practical
1) place 30cm^2 of dilute HCl (measured with a measuring cylinder) into an expanded, insulated polystyrene cup in a beaker
2) record the temperature every 3 minutes using a thermometer
3) after 3.5mins, add 30cm^2 NaOH
4) stir continuously
5) take readings every minute for 10mins
6) plot s graph and extrapolate lines
Accuracy when measuring the enthalpy of a solution
- wash equipments with solutions before use
- dry equipment
- if solid reagent is used, weigh with balance
- if the reaction is too slow, cooling occurs, decreasing accuracy of ΔT; take temp at regular intervals and extrapolate back to 0
- measure initial temperatures of both for an average initial temperature
Errors in enthalpy in solution
- assuming volume (and mass) of water is equal to volume (and mass) of solution
- assuming Cp to be the same as water
- energy lost to surroundings over time (relatively small)
- neglecting Cp of calorimeter; ignoring energy absorbed by apparatus
- incomplète réaction
Combustion enthalpy practical
1) measure 100cm^3 water (w/ measuring cylinder)
2) place water in calorimeter
3) place liquid in spirit burner and weigh (w/ balance)
4) note initial temperature of water (w/ thermometer)
5) set up calorimeter above spirit burner on tripod
6) it tie the liquid
7) extinguish the flame after ΔT = 20°C
8) re-weigh the burner to find out how much liquid has been burnt
Error in the combustion calorimetry
- energy lost to the surroundings (relatively large); flame heat escaping as convection current; using a thermometer which reads to less than 1°C is not justified
- unlikely to give more than 2sf
- incomplète combustion or energy transfer; évaporation of fuel
- lack of special equipment to allow calorimeter to capture energy (e.g. draught excluder)
- H2O (g) is not standard
Hess’ law
- for a given chemical change, overall energy change will always be the same where the change takes place in one or multiple steps
- derived from first law of thermodynamics; energy is always conserved
Shorthand Hess’ law
Total enthalpy change for a reaction is independent of the rate by which the chemical reaction takes place
Hess law is used to…
… calculate enthalpy changes for reaction that cannot be measured directly by experiments; alternative reactions are carried out that can be measured experimentally
Examples of reactions that cannot be measured experimentally
- adding the right amount of water in hydration reactions
- measuring temperature changes of solids
- impossible to measure temperature to decompose a solid
- reactions that do not readily occur
Dashed arrow show…
… indirect routes; signs must be observed so that the reactions measured point towards the product
When drawing a Hess cycle, remember to show;
- intermediates
- data
- calculation steps
Homologous series enthalpy changes
- enthalpy of combustion increases in constant increments
* because there is a constant amount and type of extra bonds being broken and made
