5.2 rates of reactions- enthalpy Flashcards
lattice enthalpy
a measure of the strength of ionic bonding in a giant ionic lattice
the enthalpy change that accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions
born-haber cycle
lattice enthalpy must be calculated indirectly using known energy changes in an energy cycle
2 routes
born-haber cycle route 1
formation of gaseous atoms:
-changing the elements in their standard states into gaseous atoms
-this change is endothermic as it involves bond breaking
formation of gaseous ions:
-changing the gaseous atoms into positive and negative gaseous ions
-overall this change is endothermic
lattice formation:
-changing the gaseous ions into the solid ionic lattice
-this is the lattice enthalpy and is exothermic
born-haber cycle route 2
converts the elements in their standard states directly to the ionic lattice
there is just one enthalpy change, the enthalpy change of formation and this is exothermic
standard enthalpy change of formation
the enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions, with all reactants and products in their standard states
Na(g) + 1/2Cl2(g) –> NaCl(s)
standard enthalpy change of atomisation
the enthalpy change that takes place for the formation of one mole of gaseous atoms from the element in its standard state under standard conditions
Na(s) –> Na(g)
1/2Cl2(g) –> Cl(g)
always an endothermic process as bonds are being broken to form gaseous atoms
first ionisation energy
the enthalpy change required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions
endothermic as energy is required to overcome the attraction between a negative electron and the positive nucleus
what is the opposite of ionisation energy
electron affinity
-measures the energy to gain elecs
first electron affinity
enthalpy change that takes place when one elec is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions
exothermic as the elec being added is attracted towards the nucleus
second electron affinities
endothermic
a second elec is being gained by a neg ion, which repels the elec away
so energy must be put in to force the neg charged elec onto the neg ion
standard enthalpy change of solution
the enthalpy change that takes place when one mole of a solute dissolves in a solvent
if the solvent is water, the ions from the ionic lattice finish up surrounded with water molecules as aqueous ions
Na+Cl-(s) + aq –> Na+(aq) + Cl-(aq)
dissolving process
when a solid ionic compound dissolves in water, two processes take place:
- the ionic lattice breaks up
- water molecules are attracted to, and surround, the ions
what types of energy change are involved in the dissolving process
1- the ionic lattice is broken up forming separate gaseous ions. this is the opposite energy change from lattice energy, which forms the ionic lattice from gaseous ions
2- the separate gaseous ions interact with polar water molecules to form hydrated aqueous ions. the energy change involved is called the enthalpy change of hydration
enthalpy change of hydration
enthalpy change that accompanies the dissolving of gaseous ions in water to form one mole of aqueous
ions.
factors effecting lattice enthalpy
ionic size
ionic charge
effect of ionic size on lattice enthalpy
as ionic size increases:
- ionic radius inc
- attraction between ions dec
- lattice energy less negative
- MP dec
effect of ionic charge on lattice enthalpy
as ionic charge inc:
- attraction between ions inc
- lattice energy becomes more neg
- MP inc
factors effecting hydration
ionic size
ionic charge
effect of ionic size on hydration enthalpy
as ionic size inc:
- ionic radius inc
- attraction between ion and water molecules dec
- hydration energy less neg
effect of ionic charge on hydration enthalpy
as ionic charge inc:
- attraction with water molecules dec
- hydration energy becomes more neg
entropy
dispersal of energy in a system which is greater,
the more disordered a system
units are JK-1 mol-1
entropies of diff states
solid- smallest entropies
gases- greatest
predicting entropy changes
- if a system changes to become more random, energy can be spread out more- there will be an entropy change which will be positive
- if a system changes to be less random, energy becomes more concentrated- the entropy change will be negative
changes of state - entropy changes
entropy incs during changes in state that give a more random arrangement of particles
solid -> liquid -> gas
when any substance changes state from solid to liquid to gas..
MP and BP incs the randomness of particles
energy is spread out more and change in entropy is positive
change in number of gaseous molecules - entropy changes
reactions that produce gases result in an inc in entropy
- production of a gas incs the disorder of particles
- energy is spread out more and change in entropy is positive
standard entropies
the entropy of one mole of a substance, under standard conditions (100kPa and 298K)
- units of JK-1 mol-1
- always positive
feasibility
used to describe whether a reaction is able to happen and is energetically feasibly
free energy
the overall change in energy during a chemical reaction is called the free energy change
is made up of two types of energy:
-enthalpy change ΔH. this is the heat transfer between the chemical system and the surroundings
-entropy change at the temp of the reaction TΔS. this is the dispersal of energy within the chemical system itself
gibb’s equation
relationship between the two types of energy, ΔH and TΔS.
ΔG = ΔH - TΔS
free energy change= enthalpy change with surroundings - temperature x entropy change of system
condition for feasibility
feasibility of a reaction depends on the balance between ΔH and TΔS in the gibb’s equation
for a reaction to be feasible, there must be a decrease in free energy:
ΔG < 0
limitations of predictions made for feasibility
many reactions have a negative ΔG and do not seem to take place
although it indicates the thermodynamic feasibility, it takes no account of the kinetics or rate of a reaction
oxidising agent
takes elecs from the species being oxidised
contains the species that is reduced
reducing agent
adds elecs to the species being reduced
contains the species that is oxidised
construction of redox equations using half equations
write out two half equations reduction and oxidation
balance electrons in each equation
add and cancel electrons
cancel any species that are on both sides of the equation
construction of redox equations using oxidation numbers
assign oxidation numbers to identify the atoms that change the ox number
balance only the species that contain the elements that have changed ox number
e.g. to match the total increase of +6 for sulfur, you need a total decrease of -6 from nitrogen
balance any remaining atoms
predicting products of redox reactions
in equations you might not know all the species involved in the reaction and you might need to predict any missing reactants or products
in aqueous redox reactions, H2O is often formed
other likely products are H+ and OH- ions, depending on the conditions used
ensure both sides of the equations are balanced by charge
manganate (VII) titration procedure
- a standard solution of potassium manganate (VII), KMnO4, is added to the burette
- using a pipette, add a measured volume of the solution being analysed to the conical flask. an excess of dilute sulfuric acid is also added to provide the H+ ions required for the reduction of MnO4- ions
- during the titration the manganate solution reacts and is decolourised as it is being added. the end point of the titration is judged by the first permanent pink colour, indicating when there is an excess of MnO4- ions present.
- repeat the titration until you obtain concordant titres
reading the meniscus - manganate (VII) titration
KMnO4 is a deep purple colour and is very difficult to see the bottom of the meniscus through it
in manganate (VII) titrations, burette readings are read from the top rather than the bottom of the meniscus
titre is the same
examples of manganate (VII) titrations
used for the analysis of many diff reducing agents, e.g:
- iron (II) ions, Fe2+
- ethanedioic acid, (COOH)2
iodine/thiosulfate titration procedure
- add a standard solution of Na2S2O3 to the burette.
- prepare a solution of the oxidising agent to be analysed. using a pipette, add this solution to a conical flask. then add an excess of potassium iodide. the oxidising agent reacts with iodine ions to produce iodine, which turns the solution a yellow-brown colour
- titrate this solution with the Na2S2O3. during te titration, the iodine is reduced back to I- ions and the brown colour fades gradually, making it difficult to decide on a end point. this problem is solved by using starch indicator. when the end point is being approached, the iodine colour has faded enough to become a pale straw colour.
using starch for the end point
iodine/thiosulfate titration
when the end point is being approached and the iodine has faded enough to become a pale straw colour, a small amount of starch indicator is added. a deep blue-black colour forms to assist with the identification of the end point. as more sodium thiosulfate is added, the blue-black colour fades
at the end point, all the iodine will have just reacted and the blue-black colour disappears
examples of iodine-thiosulfate titrations
for the analysis of different oxidising agents e.g:
- chlorate (I) ions, ClO-
- copper (II) ions, Cu2+
Analysis of copper
Iodine/thiosulfate titrations can be used to determine the copper content of copper (II) salts or alloys.
For copper (II) salts, Cu2+ ions are produced simply by dissolving the compound in water
Insoluble copper (II) compounds can be reacted with acids to form Cu2+ ions
Cu2+ ions react with I- to form a solution of iodine I2 and a white precipitate of copper (I) iodide CuI. The mixture appears a brown colour
The iodine in the brown mixture is then titrated with a standard solution of sodium thiosulfate
Thiosulfate
S2O3^2-
Half cell
Contains the chemical species present in a redox half-equation
A voltaic cell can be made by connecting together two different half cells, which then allows electrons to flow
In the cell, the chemicals in the two half cells must be kept apart
What happens if chemicals in two half cells mixed?
Electrons would flow in an uncontrolled way and heat energy would be released rather than electrical energy
Voltaic cell
Electrochemical cell
Converts chemical energy into electrical energy
Metal/metal ion half-cells
A half cell that consists of a metal rod dipped into a solution of its aqueous metal ion
At the phase boundary where the metal is in contact with its ions, an equilibrium will be set up
In an isolated half cell, there is no net transfer of electrons either into or out of the metal
When two half cells are connected, the direction of electron flow depends on the relative tendency of each electrode to release electrons
Ion/ion half cells
Contains ions of the same element in different oxidation states
There is no metal to transport electrons either into or out of the half cell, so an inert metal electrode made out of platinum is used
Equilibrium in a half cell
Written so that the forward reaction shows reduction and the reverse shows oxidation
Negative electrode
The electrode with the more negative metal loses electrons and is oxidised
Positive electrode
The electrode with the less reactive metal gains electrons and is reduced
Standard electrode potential
Measures the tendency to be reduced and gain electrons
The standard chosen is a half cell containing hydrogen gas and a solution containing H+ ions
An inert platinum electrode is used to allow electrons into and out of the half cell
The Standard electrode potential is the emf of a half cell connected to a standard hydrogen half cell under standard conditions
The Standard electrode potential of a standard hydrogen electrode is exactly 0V
Standard electrode potential standard conditions
Solutions have a conc of 1 mol dm-3
298K
100kPa
Measuring a Standard electrode potential
The half cell is connected to a standard hydrogen electrode
- the two electrodes are connected by a wire to allow a controller flow of electrons
- the two solutions are connected with a salt bridge which allow ions to flow. The salt bridge typically contains a concentrated solution of an electrolyte that does not react with either solution. E.g. strip of filter paper soaked in aqueous potassium nitrate KNO3
The more negative the E^o value the greater the reactivity of a metal in losing electrons
The more positive the E^o value the greater the reactivity of a non metal in gaining electrons
Cell potentials
Cells can easily be assembled using any half cells
The emf measured is then a cell potential, Ecell
Measuring standard cell potentials
- Prepare two standard half cells
- Connect the metal electrodes of the half cells to a voltmeter using wires
- Prepare a salt bridge by soaking a strip of filter paper in a saturated aqueous solution of potassium nitrate
- Connect the two solutions of the half cells with a salt bridge
- Record the standard cell potential from the voltmeter
Calculation of a standard cell potential from standard electrode potentials
E^o cell= E^o (positive electrode) - E^o (negative electrode)
Predicting redox reactions
Predictions can be made about feasibility of any potential redox reactions using standard electrode potentials
- the most negative system has the greatest tendency to be oxidised and lose electrons
- the most positive system has the greatest tendency to be reduced and gain electrons
Oxidising agents take elecs away from the species being oxidised, so are reduced and are on the left
Reducing agents adds elecs to the species being reduced, so are oxidised and are on the right
You can predict the feasibility of redox reactions from E^o values
Strong reducing agent
More negative E^o value
As is oxidised so loses elecs more easily
Strong oxidising agent
More positive E^o value
As is reduced so gains elecs more easily
Limitations of using E^o values
Reaction rate- electrode potentials May indicate the thermodynamic feasibility of a reaction but they give no indication of the rate of a reaction
Conc- standard electrode potentials are measured using concs of 1moldm-3. Many reactions take place using concentrated or dilute solutions. If conc different then value of electrode potential will be diff from the standard value
Standard electrode potentials and ΔG
ΔG^o = -nFE^ocell
n is moles of elecs transferred in the balanced equation
F is faradays constant
Primary cells
Non rechargeable
When in use, electrical energy is produced by oxidation and reduction at the electrodes
However, the reactions cannot be reversed
Eventually the chemicals will be used up, voltage will fall, the battery will go flat, and the cell will be discarded or recycled
Are alkaline based on zinc and manganese dioxide, Zn/MnO2, and a potassium hydroxide alkaline electrolyte
Secondary cells
Rechargeable
Cell reactions producing electrical charge can be reversed during recharging
The chemicals in the cell are then regenerated and the cell can be used again
Examples:
- lead-acid batteries used in car batteries
- lithium ion cells used in laptops, phones etc
Fuel cells
Uses energy from the reaction of a fuel with oxygen to create a voltage
- the fuel and oxygen flow into the fuel cell and the products flow out. The electrolyte remains in the cell
- fuel cells can operate continuously provided that the fuel and oxygen are supplied into the cell
- don’t have to be recharged
Hydrogen fuel cells
Can have either an alkali or acid electrolyte.
Produce no carbon dioxide during combustion, with water being the only combustion product.
