Exam 3

Competitive reactivity

Exhibited in salts composed of weak acids and bases

Largest K

Components of reaction with largest K dictate concentration of each ion at equilibrium

Reactions that influence PH

Contain H3O+ and OH-, starting concentration comes from reaction with largest K

Stronger acid or base

Dominates reactivity with water

Neutral pH

Cation forms strong base, anion forms strong acid

Basic pH

Cation forms strong base, anion is conjugate base of weak acid

Acidic pH no metal

Cation is conjugate acid of weak base, anion forms strong acid

Acidic pH metal

Cation forms is transition metal, group 13/14 metal, , anion forms strong acid

PH depends on higher K

Cation is conjugate acid of weak base, anion is conjugate base of weak acid

Common ion effect

The addition of an ion to a system at equilibrium causes a shift, reduces solubility

Strongest acid reacts with

Strongest base

Buffer

Solution capable of resisting large changes in pH upon addition of strong acid or base

Best buffer

50/50 mix of conjugate acid/base pair

Moles HA in buffer solution

Max capacity for strong base

Moles A- in buffer solution

Max capacity for strong acid

Buffer capacity

Ability of a buffer to resist change, depends on concentrations of A- and HA

Ideal buffer standards

pKa of acid must be near desired pH of buffering, concentrations of A- and HA are 10x greater than expected strong acid/base concentrations, buffer range is within 1 pH unit of pKa

Why is controlling pH important

Enzymes function best under optimal pH conditions to form strong ionic bonds

Titration

Chemical reaction used to quantitatively determine pKa, molar masses, concentrations

Titrant

Known concentration, calibrated against a standard concentration

Analyte

Carefully measure mass/volume to determine pKa, molar mass, or concentration

Equivalence point

Base and acid concentrations are equal, analyte is completely reacted, occurs when HA is completely reacted with base (concentration dependent

Indicator

Signify end of titration, dyes are chosen so that color change occurs at pH=pKa of dye

Titration curve

Plotting pH against base volume using pH meter, pH as a function of titrant solution volume added

Titration curve equivalence point

Steep change in pH

Buffer zone (pKa)

Revealed by weak acids titrated with a strong base, controlled by acid identity, independent of concentration

Acid identity controls

Buffer zone, equivalence point pH

Acid quantity controls

The volume of titrant required to reach equivalence, more weak acid can consume more base and maintain desired pH

pH at equivalence

Increases as pKa increases, Ka decreases

Center of buffer region

pH=pKa, occurs at crossing point of of HA and A- speciation curves, independent of concentration

HA and A- speciation curves

Cross at pH=pKa, molecular interpretation of pH change

Operate similarly to weak acids

Buffer solutions and titrations of weak bases

Titrating a strong base with a weak acid

Identical principle to weak acid/strong base titration, start with base end with conjugate acid

Weak base speciation curves

Usually plotted as function of pH not pOH,

Multiprotic acids titration step 1

As H2CO3 is consumed, HCO3- is produced, CO3 concentration remains around 0

Multiprotic acids titration step 2

Once H2CO3 is fully consumed, HCO3- is dominant species, CO3 concentration rises as HCO3- is consumed, at high pH only CO3 remains, both protons are eventually removed

Equilibrium of dissociation

Subtly dictated by delicate balance of enthalpy and entropy, when rate of dissolution equals rate of recrystallization dynamic equilibrium is reached

Thermodynamics of dissolution

Lattice dissociation, ion solvation by H2O-water and ion interactions promote dissolution, must overcome lattice energy to dissolve, ion-ion electrostatic interactions break

Solvation enthalpy

Ion-water interactions are electrostatic attractions that stabilize dissociated ionic solids, ion-water dipole attractions always negative, but lower than lattice attraction, r-big ions yield low attraction ot water and low solubility

Lattice enthalpy

Ion-ion interactions are always net negative and dominate trends in solubility, r-big ions yield low electrostatic attraction

Solvation enthapy ion properties

Ion-water attraction for +1/-1 ions are purely electrostatic, +2/-2 charge ion-water interactions comparable to covalent bonding energies

Dissolution chemistry

Driven by breaking solute-solute interactions and forming solvent-solute interactions

Trends in r

Two big ions in lattice are insoluble because solvation enthalpy is too low and possibly covalent, one small ion and one big ion is soluble because enthalpy solvation is stabilizing for small ion

Salts with ion of +2/-2 charge or higher

Can undergo exothermic dissolution if lattice is weaker than ion-water bonds

Heat of hydration

Enthalpy of solvent + enthalpy of mix, always exothermic

Enthalpy of solution

Enthalpy of solute + enthalpy of mix + enthalpy of solvent

Enthalpy of solute

Always endothermic = -lattice enthalpy

Salt entropy dissolution

salt goes from fixed ions to mobile ions, translational motion, ion-water rotation and vibrations, entropy is gained

Solute entropy dissolution

Goes from free water molecules to water-ion molecules with restricted motion, entropy is lost, entropy loss is greater than salt entropy gain for ions with high charge or small size

Solute dissociation

Spread solute into dense gas, positive enthalpy, positive entropy

Solvent dissociation

Increase intermolecular separation of solvent, positive enthalpy, positive entropy

Mixing dissociation

Combines solute and solvent dissociation, negative enthalpy, positive entropy

Ksp

Thermodynamic quantity related to limit of solubility for a compound, dissolution usually enthalpy driven, entropy can help with dissolution

Ksp greater than 1

Soluble salt

Unsaturated solution

Q is less than Ksp, no solid will reside in the solution, all ions in aqueous phase

Saturated solution

Q=Ksp, maximum limit in solubility achieved

Supersaturated solution

Q is greater than Ksp, thermodynamically unstable solution where salt will spontaneously crystallize out of aqueous phase

Strong acid

Can be used to promote dissolution of salts containing weakly basic anions, protonation of base gives charge neutral product, eliminates ion-ion attraction

Salt whose anion is a conjugate base of weak acid

Solubility increases as pH decreases

Strong acids can be used to

Promote dissolution of salts containing weakly basic anions

Bronsted-Lowry chemistry

Can be used to increase solubility through the anion

Lewis acid/base chemistry

Can be used for the cation

Coordination complex

Coordinate covalent bonds, ligand, and metal center, metals typically form cations, attract lone pairs of molecular species

Coordinate covalent bonds

Electron pair comes from ligand, lewis base

Ligand

Molecular species that donates an electron pair to bond with metal, Lewis base

Metal center

Accepts electron pair from ligand, Lewis acid

Complexation

Lewis acid-Lewis base reaction, can be used to solubilize insoluble salts by forming a coordination complex in solution

Nucleophile

Nucleus lover

Electrophile

Electron lover

Kb of ligands

Can be used to predict equilibrium position of complexation reactions

Stronger bases

Larger Kb, displace weaker bases in complex ion equilibrium, ligand exchange

Chelation effect

Preferred binding of multidentate ligands to metals, driven by entropy, can be more effective at increasing solubility of insoluble salts than multidentate ligands

Le Chatelier

Drive reaction into forward direction by removing metal cation, ligand-metal bonding and entropic effects drive complex formation

Binding affinity

Increases as number of anchoring groups on a chelating ligand decreases

One binding group detaches

Probable the other will remain anchored to prevent ligand from dissociating, entropically driven dissolution due to fewer number of molecules needed to bind to metal center

Binding strength of ligand

Increases as number of coordination sites on ligand increases

Galvanic cells

Where reduction and oxidation spontaneously occur, electron kinetic energy converted to heat, thermal energy lost as light to surroundings, current is generated by a spontaneous redox reaction

Electron transfer in galvanic cell

Spontaneously transfer from anode to cathode, from negative voltage to positive voltage, from high potential energy to low potential energy because it decreases their electrostatice potential energy

Anode

Oxidation

Cathode

Reduction

Standard cell potential

Measured cell potential under standard conditions

Potential difference

Amount of electrical work that can be done by 1 Coulomb of charge, due to relative in Fermi energy of metal electrodes

Voltage

Potential difference between two points in a circuit, free energy per charge, measure of electrical work

Electromotive force, emf

Total energy per unit charge provided by a source, measure of the thermodynamic driving force (free energy) for approaching equilibrium

Current

Rate of electron flow per second

Ampere

1 C per second, allows us to count electrons, independent of time

Current flows because

Charge spontaneously transfers from a half-cell of high potential energy to low potential energy

Negative voltage

Repels negatively charged electrons

Positive voltage

More stable for electrons,

Electrochemical potentials

Measure of free energy

q

Amount of charge, quantified by moles of electrons

Electrochemical work

Can be converted to heat, light, motion, or chemical bonds

Standard cell potentials

Standard hydrogen electrode, defines 0 potential and 0 voltage, arbitrary, and required since it is impossible to measure Ecell for a single reaction, potential difference is independent of SHE, report reduction so change sign for oxidation

Standard reduction potential

Standard hydrogen electrode is the anodic cell

Cell potentials

Enable us to measure Keq for many reactions

Electrochemistry

Can be used to measure very large Kf and very small Keqs

Nernst equation

Ecell=Ecell-(RT/nF)lnQ, describes how far we deviate from Keq, represents thermodynamic driving force for electrochemical systems

Q less than Keq

Products are favored

Q=Keq

Equilibrium

Q is greater than Keq

Reactants are favored