Unit 3 - Electrochemistry and Potentiometry Flashcards
Counterions
salts that shield ions from interacting electrostatically. This is because of the opposite charges of the counterions being attracted to the ions trying to interact electrostatically
Ex. Na+ being attracted to SO42- and Cl- being attracted to Ca2+
This improves solubility because it increases the distance of the electrostatic compound, thus, is better at separating. Ksp increases with added salt
What would a graph look like if its eq’m constant was truly constant? Why wouldn’t a line be straight?
It would be a horizontal line. A line would not be straight because high salt concentrations change the “constants” to shift more towards the ion products
4 properties of Ionic strength
1) ionic strength increases as concentration and ion charge increases
2) strength of effect: monovalent < divalent < trivalent < etc.
3) For 1:1 salts (monovalent), ionic strength equals molarity
4) For other salts, ionic strength is greater than molarity
Activity
The effective concentration of each species in an eq’m. It is related through concentration through the activity coefficient, gamma
Activity in neutral molecules vs. ionic molecules
For neutral molecules, the activity coefficient is generally unity (gamma = 1)
For ions, the activity coefficient can be calculated from ionic strength
Redox reaction
change in the oxidation states of the reactants
Reduction
gain of electrons (oxidation state decreases)
Oxidation
loss of electrons (oxidation state increases)
Salt bridge
feeds ions that make the solution electrically neutral. It completes the circuit by allow ions to conduct current across the salt bridge
usually a tube with a porous frit at each end and filled with electrolyte (ex. KCl). It prevents direct mixing of the two electrolyte solutions. It avoids direct reaction of Cu2+ at Zn(s) electrode. Forces electrons through external circuit
Anode
electrode where oxidation occurs
Cathode
electrode where reduction occurs
Cell potentials
As the cell operates, the potential will gradually decrease to zero (at eq’m). The potential difference (voltage) measured at any point reflects the driving force of the associated redox reaction towards eq’m
Positive cell potential
galvanic cell (spontaneous reaction; can perform electrical work)
Negative cell potential
electrolytic cell (non-spontaneous; electrical work must be done to drive the chemical reaction)
Standard half-cell potentials are measured at standard conditions. Name standard the 4 conditions
1) activity of all reactants and products is unity, a = 1
2) for solutes, a concentration of approx. 1.0 M
3) For gases, a pressure of 1.0 atm
4) Temperature = 25 degrees Celsius
Standard potential
a measure of the driving force for a reaction from a state of unit activity for reactant(s) and product(s) to their eq’m concentrations when coupled with the SHE half-cell. That is, how far the rxn is from eq’m
Standard half-cell potentials are _____ for reductions reactions
positive when the reduction rxn is spontaneous relative to the SHE
negative when the reduction rxn is non-spontaneous relative to the SHE
independent of the # of moles of reactants or products in the balanced half-rxn
Standard Hydrogen Electrode (SHE)
the reference point to which other half-cells are compared. It is assigned a zero potential, E0 = 0.000V (at all temps)
Bubble 1 atm of H2(g) to maintain saturated solution
Serves as either cathode or anode
Reversible rxn
Saturated Calomel Electrode (SCE)
SHE is not convenient in lab experiments, so SCE is used. It consists of a paste of Hg metal and calomel (Hg2Cl2), has KCl filling solution. Salt bridges are built into these electrodes
Ag/AgCl reference electrode
more convenient to use than SCE
consists of a silver wire with layer of silver chloride
KCl filling solution
Salt bridge is built in
Potentiometry
use of electrode potentials to determine analyte concentrations
Inert electrodes
Respond to redox couples without participating directly in the reaction. It serves as an interface for transferring electrons from one half-cell to the other. Very useful for redox titrations
Ex.
Platinum, gold, palladium, carbon
Electrodes of the first kind - Metal electrodes
not widely used, but can be used for precipitation and complex formation titrations
Not all metal/metal cation systems establish equilibrium quickly
Poor selectivity (prone to reaction with other metals)
Some metals dissolve at low pH or are easily oxidized
Ex. Ag/Ag+, Hg/Hg2+ (neutral solutions) and Cu/Cu2+, Zn/Zn2+ (deaerated solutions)
Electrodes of the second kind
Respond to ions that form a sparingly soluble salt with the metal
ex. Silver electrode with chloride
AgCl(s) Ag(s) + Cl-
can be used for precipitation and complex formation titrations, but not usually used
Electrodes of the second kind is just an electrode of the first kind coupled to a Ksp eq’m
Redox Titrations
An electrochemical analog of an acid-base titration. A redox active analyte can be titrated with a strong oxidizing or reducing agent to determine the quantity of analyte
Standard oxidizing and reducing agents definition
Strong, standardized oxidizing or reducing agents
Redox reactions go nearly to completion
Equivalence point
Amount of oxidant/reductant added is equal to the amount of the analyte
Endpoint
Observable change that (approximately) signals the equivalence point
2 techniques that can determine the endpoint of a titration
1) tracking changes in potential
2) redox indicator dyes
2 purposes of auxiliary oxidizing and reducing agents
1) used for pre-oxidation or pre-reduction of samples
2) excess auxiliary agent is relatively easy to quench or remove
Auxiliary reducing agents examples
solid metal (ex. Zn, Al, Ag, etc.) A very large negative E0 means it is a very strong reducing agent
Auxiliary oxidizing agents examples
ex. Bismuthate, peroxydisulfate, hydrogen peroxide
A very large positive E0 means it is a very strong oxidizing agent
Standard reducing agents examples
ex. Ferrous ion (Fe2+), Iodine/Sodium thiosulfate (S2O3 2-)
The lower the E0, the better. It makes the cell potential high, making the reaction more spontaneous
Standard oxidizing agents examples
ex. Permanganate ion (MnO-4), Ceric ion (Ce4+), Dichromate ion (Cr2O7 2-)
Have positive E0
When will the potential of a cell be 0V? How can there be potential?
Potential of a cell is 0 when any reaction is at equilibrium.
ex. Ce4+, Ce3+, Fe3+, Fe2+ are mixed in the same beaker with nothing to prevent them from reacting, therefore, any reaction between these species will come to an eq’m
For a cell to have potential, there must be at least one redox reaction that is never at equilibrium
ex. the reference electrode is immersed in the beaker with Ce4+, Ce3+, Fe3+, Fe2+, but its filling solution is separated from the beaker solution by a salt bridge. The redox rxns between the beaker and reference electrode do not come to an eq’m
Redox indicators
Electrochemical cell is not necessary, so colour changes can be used as indicators in redox titrations. A colour change is observed when the system changes from a large excess of the oxidized form to a large excess of the reduced form (or vice versa)
Colour changes are observed within +/- 59.2/n mV of E0 for the indicator (n=2 for many indicators)
Common redox indicators
1,10-Phenanthrolines, Starch/Iodine, Diphenylamine derivatives
When is there electric potential? How are they like in potentiometry?
Anytime there is a separation of charge over a distance. In potentiometry, these potentials usually appear as liquid junction potentials or membrane potentials
Liquid junction potential (E_j)
a potential difference that develops across an ion permeable boundary between different electrolyte solutions
Caused by the diffusion of cation and anion at different rates (make sure ions have similar mobility)
Resultant separation of charge generates a voltage (that can be measured)
Can reach voltages on the order of approx. 10^-2 V
Limitation of liquid junction potential
it limits the accuracy of potential measurements (their contribution to the net potential is unknown). Minimize liquid jxn potentials (few mV) by using concentrated electrolytes where the cation and anion have similar mobility (ex. KCl instead of NaCl)
Membrane potentials
ex. glass interface, bulk SiO2 terminates in SiOH groups
Depending on the pH, the membrane will acquire an amount of negative charge (for visual look on review sheet near the end, or part 3, slide 7). The negative charges on the glass interface cannot diffuse. The anionic interface is screened by cations in solution forming an electric double layer
The separation of charge at the double-layer generates a membrane potential, E_m. The magnitude of E_m will depend on the charge on the membrane (ie. interface)
Boundary potentials
If a membrane is thin and conductive, the difference between two membrane potentials (each side of the interface) can be measured as a boundary potential, E_b. The boundary of a thin glass membrane is pH-sensitive
Glass as a membrane and its physical properties
it is amorphous SiO2. No long range structure, and is cooled to rigidity without crystallization
ex. Soda-lime silicate glass. Soda (Na2O) is used a “flux” to lower melting point in manufacture. Lime (CaCO3) is used to prevent glass from dissolving
Consists of irregular arrangement of SiO4 tetrahedra
Incomplete bonding, Si–O- groups associated with cations
Can hydrate surface layer; exchange monovalent ions for protons
***binding of hydrogen ions is more strongly favoured than alkali metal cations
Glass as a pH sensitive membrane
Na+ and H+ ions conduct electricity in the hydrated outer layers of the glass membrane. Na+ conducts electricity in the dry interior
Ion selective electrodes
Pressure difference is coming from ion-selective membrane, such as the different concentration of protons in the internal solution and external solution. This is why we need 2 reference electrodes, because they give constant potential
Combination glass electrode
internal Ag/AgCl reference electrode. Junction to maintain electrical contact
Alkaline error
At high pH, the concentration of H3O+ is very low
Sodium (and some other monovalent metals) can associate with the glass membrane
Generate a change in potential
Measure a pH between 0 and -1 units lower than the true value (positive bias)
Modern glass membranes are relatively immune to alkaline error up to approx. pH=12
Two methods used to determine the selectivity coefficient (of the Nikolsky equation)
1) separate solution method - measure the analyte and interferent ion(s) separately
2) fixed interference method - prepare a calibration curve with different activities of analyte and a fixed activity of interference
Total Ionic Strength Adjustment Buffer (TISAB)
A buffer solution that increases the total ionic strength to a high level (to make them similar)
The ionic strength of the buffer overwhelms the contribution from the sample, such that the sample and calibration standards can be measured at approximately the same ionic strength, even though the ionic strength of the sample is unknown
This is important because we are using activity (which is directly related to ionic strength) to be measured, which can be correlated to the analyte concentration
Sources of error in pH measurements
Chemical errors:
Junction potential different between standard and sample
Junction potential drifts over time (ex. material in salt bridge could leak out)
Alkaline error (aka sodium error)
Acid error (electrode response decreases at pH<1)
User errors: (there are more listed)
Standard solutions used for calibration are inaccurate
Glass membrane is not clean
Calibration and sample temp not matched
Types of ion-selective electrodes
Glass
Solid-state
Liquid membranes
Ionophore
a hydrophobic organic molecule that serves as an ion exchanger (ie. reversibly binds an analyte ion)
Liquid membrane ion-selective electrodes
function is analogous to glass membrane. Glass membrane is replaced with an ionophore in a hydrophobic matrix. A membrane potential develops when the activity of the analyte ion is different between the two sides of the membrane
Solid-state ion-selective electrodes
Inorganic crystals
Mobile monovalent ions conduct electricity through crystal vacancies
Vacancies created by doping
