w6 Flashcards
Oxidation-Reduction Reactions Characteristics
- Occurs simultaneously when electrons are transferred from one species to another
The reactant species that gains electrons is said to be reduced.
* That which loses electrons is oxidized.
* The species from which electrons are taken is oxidized (and is the reducing agent).
* The species that takes the electrons is reduced
(and is the oxidizing agent).
Oxidation States
a measure of the degree of oxidation of an element in a compound compared with when it is uncombined
Oxidation
accompanied by an increase in the oxidation state of an element.
ANODE
Reduction
accompanied by a decrease in the oxidation state of an element.
CATHODE
GUIDELINES to Redox reactions
- The oxidation state of each atom in a pure element is zero.
- For a simple, monatomic ion, the oxidation state is equal to the charge on the ion.
- The algebraic sum of oxidation states in an ion is equal to the ionic charge. In a molecule, the sum is zero.
- In all of its compounds and ions, fluorine is in oxidation state -1.
- In most compounds and ions, hydrogen is in oxidation state +1.
- In most compounds, O is in oxidation state -2.
Like oxygen, the halogens (F2, Cl2, Br2, I2) are always oxidizing agents in their reactions with metals and non-metals.
Electrochemistry
a field of study about the interaction between electricity and chemistry.
Voltaic cells (or galvanic cells)
direct the transfer of electrons in a spontaneous redox reaction from one compartment of the cell, through a conductor, to another compartment.
Electrolysis cells (or electrolytic cells)
in which the application of an electrical potential forces a non-spontaneous redox reaction to occur. This forced non-spontaneous process is called electrolysis
Voltaic Cells
Contains anode, cathode, salt bridge, voltmeter and stuff idk
- The electrons transferred during a spontaneous reaction are harnessed as an electrical current—electrons move from the anode, the site of oxidation, through the external circuit to the cathode, the site of reduction.
- Charge balance in each half-cell is achieved by migration of ions through the salt bridge (internal circuit). Anions move toward the anode, and cations move toward the cathode.
Inert electrodes
made of conducting materials that are too unreactive to be oxidized or reduced.
Cell EMF
a measure of the difference between the abilities of species in the half-cells to compete for electrons.
Standard Conditions
– Solutes in aqueous solution have a conc of 1.0 mol L-1
– Gaseous substances have a pressure of 1.0 bar.
– Solids and liquids (not solutions) are pure.
Standard Reduction Potentials
Different combinations of half cells will give rise to a different cell EMF.
The EMF is the difference between the two electrode potentials.
E(cell) = E(cathode) – E(anode)
E(cathode) and E(anode) are half-cell reduction potentials orelectrode potentials (Ehalf-cell)
Potential ladder for reduction half-reactions
The higher ranked the half equation is, the more powerful an oxidizing agent the species is.
At bottom: stronger reducing agents
The tendency for the rxn to occur as a reduction (as written) DECREASES down the table.
For any cell:
* The ½ reaction that occurs highest up the table is written in the FORWARD direction (as is).
* The one that is lower goes in the REVERSE direction. It is the oxidation reaction.
Primary batteries: Alkaline cell
Cannot be recharged
Anode (-):
Zn(s) + 2OH- (aq) → ZnO(s) + H2O + 2e
Cathode (+): 2MnO2(s) + H2O + 2e- → Mn2O3(s) + 2OH- (aq)
Advantages:
No gas buildup, more suitable for high drain applications. last ~50% longer than older lead acid batteries
secondary batteries: Lead Storage Battery
Reactions can be reversed (rechargable)
PbO2(s) + Pb(s) + 2H2SO4(aq) -> 2PbSO4(s) + 2H2O(l)
6 cells each generating 2.04 V = 12 V output
The reactants Pb and PbO2 are electrodes:
- as both are solids no need to separate into
different half cells - because both are solids have no effect on rxn quotient (Q)
- Thus the EMF is constant during discharge.
Voltaic Cells under Nonstandard Conditions
Dependence of Cell emf on Concentrations
* When the species in a voltaic cell are not in their standard states, cell emf can be calculated using the Nernst equation:
E (cell) = E (nought cell) - (RT/nF) lnQ
= E (nought cell) - (0.0257V/n) lnQ at 25 degrees
- R is the gas constant (8.3144 J K mol-1);
- T is the temperature (K);
- n is the moles of electrons transferred
- F is the Faraday constant (9.6485338 x 10¹ C mol-1)
- Q is the reaction quotient
pH Meters and Ion-Selective Electrodes: Nerst equation
used to determine an unknown concentration by using a measured cell potential, using a pH meter.
- The electrodes used to measure ion conc are known as ion-selective electrodes.
pH-Dependence of Oxidizing
Power of Oxoanions
ions with the generic formula AxOy ^z−
The more acidic the solution is, the more powerful oxoanions like nitrate and permanganate etc are as oxidixing agents.
If the half-equation for the reduction process involves H+ (aq) ions, then the half-cell reduction potential depends on the pH.
When pH is low, the concentration of H+ is high, which pushes the equilibrium to the right. This makes nitrate a more powerful oxidizing agent.
In general:
The lower the pH of solution, the more powerful oxoanions are as oxidizing agents.
Electrolysis: Chemical Change Using Electrical Energy
reactions are forced the non-spontaneous direction by electrical energy.
In the schematic, electrons flow in the spontaneous direction in Galvanic cells (left).
A potential drives electrons in the non-spontaneous direction in the electrolytic cell (right).
Widely used in the refining of metals such as
aluminum and in the production chlorine gas.
Electrolysis of water makes hydrogen and oxygen gases.
Electrolysis of Molten Salts
The melting point of table salt, NaCl, is 803 C.
Above 803 C sodium ions NaCl melts, and forms a molten salt of Na+ and Cl ions.
Electrodes are places in the molten NaCl and a
(sufficiently high) external voltage applied.
Sodium ions accept electrodes and form sodium metal at the cathode.
At the anode, electrons are taken from chloride ions, and chlorine gas is formed.
The production of sodium metal and chlorine gas is not spontaneous. It is driven by the electric current.
Electrolysis of Aqueous Solutions
Electrolysis of aqueous solutions is more complicated than in molten salts, due to the presence of water.
The water molecules is dissolved ions could gain electrons at the cathode, or lose electrons at the anode. The water and ions compete.
Reduction decreases the H3O+ concentration,
so the concentration of OH- ions increases.
A drop of phenolphthalein has been added to
the solution so that the formation of OH-(aq) can be detected (by the red colour).
Overvoltage
In industrial situations, large currents are used to produce large quantities of products quickly.
Gaseous products can form electrically insulating layer on the electrode.
A potential above that predicted by standard
reduction potentials must be applied to overcome the insultaing effect of the gas, called an overvoltage.
The overvoltage cannot be theoretically calculated.
Spontaneous Direction of Change and Equilibrium
the direction of reaction that would take a reaction mixture closer to a state of chemical equilibrium.
- may be extremely fast or so slow that change is undetectable.
- A decrease of enthalpy of the reaction mixture is not a sufficient criterion to determine whether a process is spontaneous.
Entropy: Dispersal of Energy and
Matter
Entropy: using a thermodynamic property to predict whether a process is spontaneous
- built on the idea that spontaneous change results in dispersal of energy
- The Second Law of Thermodynamics: Any
spontaneous process is accompanied by an
increase in the entropy of the universe.
- also observed in solitions.
- formation of a mixture does not always lead to greater disorder
The Boltzmann Equation for Entropy
S = k log W
where k is the Boltzmann constant, and W represents the number of different ways that the energy can be distributed over the available energy levels.
Entropy and the Laws of Thermodynamics
The Third Law of Thermodynamics: There is no disorder in a perfect crystal at 0 K, or S = 0.
All substances have positive entropy values at temperatures above 0 K.
Standard molar entropy (S°)
the entropy of 1 mol of a substance in its
standard state at the specific temperature
Large molecules»_space; smaller molecules.
Substances whose molecules have complex structures have more entropy than those with simpler molecules.
increases as thetemperature is raised. There are large increases in entropy when the state changes.
At a temperature T above the boiling point, the standard molar entropy includes contributions from:
1.The heat capacity of the solid between 0 K and M.P.
2.The heat of fusion of the solid
3.The heat capacity of the liquid from M.P. to B.P.
4.The heat of vaporization of the liquid
5.The heat capacity of the gas from the boiling point to T K.
Entropy Change of Reaction (∆(r)S°)
∆S° (univ) = ∆S°(sys) + ∆S° (surr)
(∆(r)S°) = (sum of the entropies of the products) - (sum of the entropies of the reactants)
= ∑n(i) S°(products) - ∑n (i) S°(reactants)
S is the entropy per mole of each product and reactant
N is the number of moles of each reactant or product in the equation
Contributions of ∆(r)S° and ∆rH° to
Spontaneity of Reaction
The entropy change of the surroundings is
∆S° (surr) = q (surroundings) / T = -∆H (sys) / T
There are 4 categories of reaction, depending on the signs of ∆S° (sys) and ∆S° (surr)
Gibbs Free Energy
e used determine whether reaction is
spontaneous from just the system.
At a specified temperature T, the free energy is defined mathematically as:
G = H – TS
The free energy of a substance is usually not important.
∆rG = ∆rH ̶T∆(r) S
∆(r)G and Spontaneity of Reaction
The second law states that a reaction is spontaneous if ΔSuniv > 0.
Reaction spontaneity can be worked out more conveniently from:
∆S (univ) = -∆(r) G / T
If ∆(r)G < 0, the reaction is spontaneous as written.
If ∆(r)G = 0, the reaction is at equilibrium.
If ∆(r)G > 0, the reaction is not spontaneous (but the reverse reaction is).
Standard Free Energy Change of
Reaction (∆rG°)
∆rG° <0, the spontaneous direction of reaction is the direction of the chemical equation.
∆rG° is directly related to the value of the equilibrium constant, thus whether reactions are reactant-favoured or product-favoured.
The free energy change is the maximum energy available for work:
∆rG = w(max)
Calculating ∆G0 (rxn)
- Determine ∆Ho (rxn) and ∆So (rxn) and use Gibbs equation:
∆G = ∆H - T∆S
ΔHrxn = ΔHprod – ΔHreact
ΔSrxn = ΔSprod – ΔSreact
Values for H and S can be found in SI data book.
- Use tabulated values of free energies of formation, ∆Gfo
∆Gfo is the standard free energy per mole for the formation of a compound from its
elements in their most stable form under standard conditions (298 K).
∆G0 (rxn) = Σ∆G0 f(products) - Σ∆G0 f(reactants)
Dependence of ∆rG° and Spontaneity
on Temperature: if ∆H is negative, ∆S is positive
reaction is always spontaneous, reverse is
always non-spontaneous
∆S positive favours reaction
∆H negative favours reaction
Dependence of ∆rG° and Spontaneity
on Temperature: if ∆H is negative, ∆S is negative
reaction is spontaneous at low T, becomes
non-spontaneous as T increases
∆H negative favours reaction
Have to do the calculation to determine spontaneity.
Dependence of ∆rG° and Spontaneity
on Temperature: if ∆H is positive, ∆S is positive
non-spontaneous at low T, spontaneous at
high T
∆S positive favours reaction
Have to do the calculation to determine spontaneity.
Dependence of ∆rG° and Spontaneity
on Temperature: if ∆H is positive, ∆S is negative
reaction is non-spontaneous at all
temperatures, reverse is spontaneous
∆rG for Non-standard Reaction Mixtures
∆rG = ∆rG° + RT ln Q
where R is the gas constant, 8.314 J K-1 mol-1, and T is the temperature (in K).
The Relationship Between ∆rG° and K
When equilibrium is reached Q = K and ∆rG = 0. If we substitute into:
∆rG = ∆rG° + RT lnQ
0 = ∆rG° + RT ln K
∆rG° = ̶RT ln K
This equation allows calculation of:
* an equilibrium constant from thermochemical data in tables.
* the standard free energy change of a reaction an experimentally equilibrium constants.
∆rG° and E°cell for Voltaic Cell Reactions
Recall the cell emf (E° cell) is related to K by:
ln K = + (nFE°cell) / RT
We can rearrange ∆rG° = ̶RT ln K to get:
ln K = -∆rG° / RT
Equate these two requations equations (at constant T):
∆rG° = ̶nFE° cell
Which relates the standard free energy of reaction to emf.
For a spontaneous reaction in a standard mixture
(Q = 1): K > 1,
E°cell > 0, and ∆rG° < 0
Dependence of Equilibrium Constants on Temperature
The dependence of the equilibrium constant on temperature is given by the following relationship:
ln K = (-∆rH° / RT) + (∆rS°/R)
R is independent of temperature changes and enthalpy and entropy are approximately constant over small temperature ranges.
The equation is linear
The van’t Hoff equation again i think
used to estimate K at one temperature from an experimentally measured value at another temperature:
ln K {1} - lnK {2} = -∆rH°/R (1/T{1} - 1/T {2})
