UWorld Solutions and Electrochemistry

Neutral Salts

Formed from a strong acid reacting with a strong base (pH = 7)

Acidic Salts

Formed from a strong acid reacting with a weak base (pH < 7)

Basic Salts

Formed from a weak acid reacting with a strong base (pH > 7)

Galvanic Cell

A spontaneous, exergonic chemical reaction (Delta G < 0) produces an electric curren.

Electrolytic Cell

Electric force from an external source (like a battery) creates a nonspontaneous, endergonic chemical reaction (Delta G > 0) within the electrochemical cell

Reduction takes place at the:

Cathode (negatively charged), attracts positively charged particles

Oxidation takes place at the:

Anode (positively charged), attracts negatively charged particles

Electrolysis

Caused by an electrolytic cell and is a nonspontaneous (endergonic) electrochemical decomposition of a compound

External Source

In an electrolytic cell, it removes electrons from the anode and toward the cathode. The electron flow drives the nonspontaneous process.

Ionic Salts

When dissolved, they dissociate into ions that can freely move and assist the passage of electrons through the solution.

Molecular Liquids

Nonelectrolyes because they do not associate into ions and do not convey electrons through the gel.

Gibbs Free Energy Equation (Electrical Current)

Delta G = -nFE(cell)

n = number of electrons

F = Faraday’s constant

Ecell = Cell potential

Delta G < 0 and E cell > 0

Spontaneous

Delta G > 0 and E cell < 0

Non Spontaneous

Kw self ionization equation

Kw = [H3O+][OH-]

Buffer System

Resists changes in pH when OH- or H+ is added to the solution. Accomplish this by containing large amounts of a mixture of either a weak acid (HA) and its conjugate base (A-) or a mixture of a weak base (B) with its conjugate base (HB), which can neutralize lesser amounts of any additional H+ or (BH).

Henderson Hasselbach Equation

HA → ← H+ + A-

Acid Dissociation Constant

Ka = [H+] [A-] / [HA]

pH

-log [H+]

pKa

-logKa

Indicator

A weak acid or base is used to visually detect the endpoint of an acid base titration. Does not react with the substance being titrated and changes color close to the equivalence point.

Equivalence Point

Stoichiometric amount of titrate needed to react with all the solution being titrated. A titration is halfway to the equivalence point when a pH is achieved that is near the pKa value of the buffer’s acid.

Polyprotic Acids

Have more than one proton that can dissociate in water. The titration curve of a polyprotic acid will show plateaus and steep climbs as the number of acidic protons in the acid, representing the buffer regions and equivalence points.

pH = pKa

[HA] (acid) = [A-] (conjugate base)

pH > pKa

[HA] (acid) < [A-] (conjugate base)

pH < pKa

[HA] (acid) > [A-] (conjugate base)

Dominant Species

pKa dominant species is the most acidic one and the pH dominant species is the most basic onet

Bronsted Lowry

Acid = donates a proton/H+ Base = accepts a proton/H+

Conjugate Acid = The species the acid becomes after losing a proton, Conjugate Base = The species the base becomes after gaining a proton

Redox Titration

A solution with a known concentration of an oxidizing agent is added to a solution with an unknown concentration of species to be measured (analyte). As the oxidizing agent is added, the analyte is oxidized and the electrode potential changes. When all of the analyte solution has been oxidized, the voltage sharply rises when passing through the equivalence point. Impurities undergoing this same behavior as the analyte may decrease the accuracy.

Solubility

The maximum amount of solute that can dissolve in a given solvent at a given temperature. Can be g/L, mg/mL, g solute/g solvent and mol/L

Increasing Temperature (Solute)

Increases the amount of solute that can dissolve

Parts Per Million (ppm)

ppm = mass of solute / mass of solution (solute + solvent) multiplied by 1o^6

Common Ion Effect

If a solution contains two dissolved compounds that have an ion in common, the common ion supplied by one compound shifts the equilibrium of dissolution of the other compound to the left (reactants) and decreases the solubility of the other compound

Decreasing the pOH/Increasing the OH in solvent solutions will:

cause an acid base reaction to disrupt the solubility equilibrium, causing a shift toward the products

Ksp

Ksp = [X]^b[Y]^c , provides an indication of the solubility of a compound at a given temperature

X and Y = molar concentrations

b and c = The respective coefficients of the balanced dissolution reaction.

Balancing Chemical Equations (Charges)

When balancing cheical equations charges must be balanced. Soichiometric coefficients effect charges.

Galvanic Voltaic Cell

Involves a spontaneous (delta -G) and (E+) and converts chemical energy into electrical energy

Standard Cell Potential for Galvanic Cell Equation

E cell = E cathode - E anode

To get reduction potential from oxidation galf reaction you must”

Switch the sign to negative

Finding Mass Consumed from Electrical Charge

1. Electric Charge is a faraday

2. 1 mol electron / 1 mol faraday

3. moles of atom / moles of electron

4. Molar mass of atom / 1 mol atom

Electron Affinity and Reduction Potential (E cell)

Species that have a stronger affinity for electrons has a more positive reduction potential (E cell) and is a weaker reducing agent. A species that has a weaker affinty for electrons has a more negative potential and is a stronger reducing cing agent.

Determining the strength of a reducing agent based on an reduction potential (E)

Weaker reducing agent (positive E) = Metal less likely to be oxidized (can be used as cathodes)

Stronger reducing agent (Negative E) = Metal more likely to be oxidized (can be used as anodes)

Litmus Paper

A pH indicator. Blue litmus paper turns red in acidic conditions and red litmus paper turns blue in basic conditions.

Titration of an Acid

A measured amount of base solution (titrant) with a known concentration is added to an unknown concentration of acid. As the titration proceeds, acidic protons react with the base in acid base neutralization reaction which changes the pH of the acid titration.

Titration curve of an Acid shows:

One buffer region and an equivalence point for each acidic proton in its structure. When enough OH- is added to the titration curve, adding more OH- will have no effect on the curve and the curve will flatten as pH of solution matches pH of titrant.

Buffer Region

Where the curve is flat

Equivalence Point

Midpoint between two buffer regions

Net Charge

The sum of all nonzero formal charges from each atom in the structure. Atoms acquire a nonzero formal charge by acid base interactions in a solution. Functional groups that are acidic act as H+ donors (decrease formal charge by 1) and basic ones act as H+ acceptors (increase formal charge by 1).

Volume and Concentration (Relationship_

As volume increases, concentration decreases

Concentration Sell (Galvanic Cell)

Use the same electrode material and ionic solution, but the concentration of the ionic solution is greater at the cathodic half cell and less at the anodic half cell. Electrons will flow from the anode (fewer cations) to the cathode (more cations) until the concentrations are equal in each half cell.

Ksp Equation

Ksp = [B]^b [C]^c (B and C are the products) (Indication of maximum solubility)

Saturated Solution

Contains the maximum amount of solute that can be dissolved per unit volume of solvent under a set of given conditions.

Common Ion Effect

When two compounds have an ion in common, the common ion supplied by one compound will cause the equilibrium to shift toward the reactants and decrease the solubility of the compounds.

Temperature Dependence of Heat

Depends on whether the dissolution reactions is endothermic or exothermic. An exothermic reaction (-Delta H) releases heat and increases the temperature (adding heat) causes the reaction to shift toward reactants. An endothermic reaction (+Delta H) absorbs heat (acts as a reactant) and the reaction will shift towards the products with decreasing temperature.

Ksp decreases for:

exothermic reactions

Ksp increases for:

endothermic reactions

Acid-Base Neutralization

Acidic acids have highest solubility in basic conditions. Acid base neturalizations can be used to increase the amount of acidic and basic solutes that will dissolve in a solvent by forming more soluble ionic salts.

Buffered Solution

Consists of a weak acid and a salt of its conjugate base or a weak base and its conjugate acid. Buffer resists change in pH by neutralizing H+ and OH- ions. The pH of a buffered solution depends on the ratio of weak acid and its conjugate base.

Buffers

The best buffer for maintaining a given pH should have a pKa that is within one unit of the desired pH. So the concentrations of pH = pKa.

Determine the number of moles recovered

1. Get the initial volume

2. Get the grams from g/mL

3. Use molar mass to get moles

4. Use ratio of moles recovered (percent of moles recovered) over total moles (100) to get moles recovered

pH of the First Equivalent Point

Between the pKa1 and pKa2

Relationship between pKa and pH

If the pH of the solution is less than the pKa (the solution is more acidic), the equilibrium is shifted toward a greater concentration of the protonated acid (HA > A-). If the pH is greater than the pKa

Conjugate Acid

The species the base becomes after gaining an H+

Conjugate Base

The species the acid becomes after losing an H+

Osmosis

When two solutions with different solute concentrations are separated by semi permeable membrane, the solvent from the solution of lower solute concentration will diffuse across the membrane into the solution of higher solute concentration until the solution concentrations equalize.

Osmotic Pressure

The pressure made by the diffusion of the solvent during osmosis. (Osmosis = iMRT)

i = van’t hoff factor (the number of species a molecule will form when dissolved in a solution, each ion counts as 1)

M = Molarity

R = 0.0821

Ksp (Solubility Product Constant)

Ksp = [X]^b [Y]^c, Represents the limit of solubility for the compound

X and Y = Molar Concentrations

b and c = Stoichiometric Coefficients

Solving for Ksp (Solubility Product Constant)

PbCl2

Ksp = (n)(2n)² = (n)(4n²) = 4n³

Lewis Acid

lewis Acid Accepts electrons

Lewis Base

lewis base donates electrons

Bronsted Acid

Bronsted acid donates protons (H+)

Bronsted Base

Bronsted base accepts protons

Relating pH to pKa

Ka = 10^-pKa = Products (pH)(pH) / Reactants (M)

1. Make ICE table, initial is given M and the other initials are 0

2. Change = -x and +x +x

3. Equilibrium = M -x and x, x

4. pKa = x² / M -x

5. solve for x (x² = half o exponent of M - x)

[OH-]

10^-pOH (mol/L)

To Solve for Ksp

1. find pOH from OH concentration

2. Use the stoichiometric equation to find moles of atom asked for

3. Ksp = [Molar concentration solved] [OH concentration] (exponent added if there is a stoichiometric coefficient)

ppm

Basis for ppm = 1,000,000 / 100

1 ppm = 1 mg/L

Delta G

-nFE cell

n = electrons transferred

F = Faraday’s constant

E = Energy potential (Ered - Eox)

The more positive the reduction potential (E):

The more likely reduction reaction will occur. Has a strong oxidizing agent and a weak reducing agent.

The more negative the reduction potential (E):

The less likely reduction reaction will occur. Has a strong reducing agent and a weak oxidizing agent.

Redox Titration

The equivalence point = when all the analyte (unknown) in the solution has been oxidized by the added titrant (an oxidizing agent). At the half equivalence point, the molar concentrations of the non oxidized analyte and oxidized analyte are equal/

A Buffer Requires:

A weak acid or a weak base to be present with the salt of its corresponding conjugate base or acid

The best buffer for a solution with pH (6.5) is

(blank) × 10^-7

Osmolarity Equation

iM = Osmotic Pressure / RT (OSmol/L)

Concentration Cell

Electrons flow from the anode to the cathode and cathode always has higher concentratiton than anode

Osmotic pressure is proportional to:

The concentration of the osmotically active solute.

Osmotic Pressure 1/Osmotic Pressure 2 = M1/M2

Calculating mmol from ph

1. H+ = 10^-1 pH = .1 M

2. .1 M = .1 mol/L

3. 1000 mmol/1 mol

4. 1 L / 1000 mL

5. = 0.1 mmol/mL

Standard Gibbs Free Energy Equation

Delta G = -RT lnKeq

Self Ionization Constant Kw

Kw = [H+][OH-] = 1 × 10^-14 at 25 c

pOH decreases as:

OH- increases

The magnitude of [H3O+]

Changes by a power of 10 for each pH unit and is given by the expression: Factor Delta [H30+] = 10^-pH

The terms [H+] and [H3O+] as hydrogen ions are not free in aqueous solutions.

Boundary Lines

Between boundary lines indicate the transition from one species and can also indicate equilibrium aqueous and precipitating species.

When is the E value flipped:

Because half reactions are usually written as reductions, one of the half reactions must be reversed to make it oxidation

Standard Reduction Potentials E

Indicates whether a molecule spontaneously gains or loses electrons and are related to the gibbs free energy Delta G. In a spontaneous process, E is positive, Delta G is negative. Electrons flow from molecules with low reduction potential to molecules with high reduction potential.

Electron Loss occurs at the:

Anode

Electron Gain occurs at the:

Cathode

MCAT Strong Acids

H2SO4 (or sulfuric acid), HI (hydrologic acid), HBr (hydrobromic acid), HNO3 (nitric acid), HCl (hydrochloric acid) and HClO4 (perchloric acid).

MCAT Strong Bases

Sr(OH)2, Ca(OH)2, LiOH, NaOH, KOH, Ba(OH)2

Concentrations > Ksp

Ions combine and precipitate

Solubility and Corrosion

Solubility is proportional to corrosion. As corrosion increases so does the solubility.

Ksp in Exothermic Reactions and Endothermic Reactions

Ksp decreases for exothermic reactions and increases for endothermic reactions when temperature increases