DAT Chem Acid-Base, Equilibrium, oxidation-reduction

Arrhenius definition of acids and bases

Acid: any species that dissolves in aqueous solutions to produce H+ ions (HCl in water forms H+ and Cl-)

Base: any species that dissolves in aqueous solution to produce OH- ions (NaOH in water forms Na+ and OH-)

Bronsted-Lowry acid and base definition

Acid: any species that acts as a proton donor

Base: any species that acts as a proton acceptor

-HCl produces and H+ ion in water that bonds to water to form H3O+, so HCl acts as a proton donor (acid) and H2O acts as a proton acceptor (base)

Lewis definition of acid and base

Acid: any species that accepts a pair of electrons (usually have a positive charge)

Base: any species that donates a pair of electrons (usually have a negative charge)

*usually metal ions

Conjugate acids and bases

Conjugate acid: base with an additional H+

-H2O (acid) and NH3 (base) react to form OH- (conjugate base) and NH4+ (conjugate acid)

Conjugate base: acid with one less hydrogen

-HCl (acid) becomes H+ and Cl- (conjugate base)

Inverse relationship between acid base strength

The stronger an acid or base is, the weaker its conjugate acid or base is and vice versa

Example: HCl is a strong acid and Cl- is a weak conjugate base

How does pKa relate to acid strength

A low pKa is a very strong acid

A high pKa is a very weak acid

Calculating pH

pH= -log[H+]

[H+]= 10^-pH

-the higher the concentration of H+ ions, the lower the pH

Calculating pOH

pOH= -log[OH-] where [OH-]= 10^-pOH

-the higher the concentration of OH- ions, the lower pOH

Relationship between pOH and pH

pH + pOH = 14

-inverse relationship between pH and pOH

-when acids are measured, pH is less than 7 and pOH is greater than 7

-when bases are measured, pH is greater than 7 and pOH is less than 7

Dissociation constants

Describes the extent at which a particular substance dissolves into ions

Ka (acids) = [H3O+][conjugate base]/[acid]

Kb (base) = [conjugate acid][OH-]/[base]

Kw (water constant) = [H+][OH-] = Ka*Kb where Kw = 1 × 10^-14

pKa and pKb

Measure how strong it hoe weak an acid or base is

pKa = -log(Ka) where the lower the pKa, the stronger the acid (likely to donate its protons)

pKb = -log(Kb) where the lower the pKb, the stronger the base (more likely to accept protons)

Strong acids vs weak acids

Strong acids completely dissociate in water and Ka is greater than 1

Weak acids partially dissociate in water and Ka is much less than 1

*pKa is inversely rekeyed to Ka, so a low pKa is a stronger acid

Strong base and weak base

Strong base: completely dissociated in water to release lots of OH- and Kb is greater than 1

Weak base: partially dissociates in water and Kb is much less than 1

*low pKb is a strong base

Binary acids

Usually include hydrogen halides (HBr, HCl, HF)

The larger the atomic radius of the halide, the more acidic (HI> HBr >HCl>HF)

Oxoacids

Acids that contain oxygen

-substances with more oxygen atoms are more acidic

*if there are equal numbers of oxygen atoms, more acidic substances have a more electronegative center atom

Neutralization reactions

Strong acids react with a strong base to produce water

HCl + LiOH = LiCl + H2O

Salts

The product of a neutralization reaction of acids and bases

-strong acid and weak base react to form an acidic salt

-weak acid and strong base react to form a basic salt

Determining whether a salt is acidic or basic

Break the salt down into its respective ions, then add H and OH. Whichever acid or base is stronger determines whether the salt is acidic or basic

Example: NH4Cl becomes NH4+ and Cl-, HCl is a strong acid and NH4OH is weak, so salt is acidic

*strong acid and strong base form a neutral salt

Buffer

Solution that resists a change in pH when small amounts of acid or base are added to it

*work best when there’s equal concentrations of weak acid, weak base, and their conjugate acids and bases

6 ways to make a buffer

1) weak acid + salt in a 1:1 ratio (HF and NaF)

2) Weak acid + strong base in 2:1 ratio (HF + NaOH)

3) salt (conjugate base) + strong acid in a 2:1 ration (NaF and HCl)

4) weak base and salt in a 1:1 ratio (NH3 and NH4Cl)

5) weak base and strong acid in a 2:1 ratio (NH3 and HCl)

6) salt (conjugate acid) and strong base in a 2:1 ratio (NH4Cl and NaOH)

Calculating the pH or pOH of a buffer (The Henderson Hasselback Equation)

pH = pKa + log[conjugate base]/[acid]

pOH = pKb + log[conjugate acid]/[base]

-ph and pKa are equal when there are equal concentrations of acid and conjugate base

*log(1) = 0, log(10) = 1, log(100) = 2

Titration

An experiment used to determine the unknown concentration of an acid or base by adding a known concentration of an acid or base to an unknown concentration of an acid or base

Indicators in a titration

A substance used to estimate the equivalence point of an acid-base titration experimentally

-if the titration involves a strong base and weak acid, the indicator will have a higher pH

-us the titration involves a strong acid and strong base, indicator will have a neutral pH (7)

Equivalence point

Point in titration where the amount of titration is just enough to neutralize the analyte (unknown) solution

-the steepest part of the titration curve is usually the equivalence point (vertical line)

Half-equivalence point

The midpoint of the buffering region that is seen when a weak acid and strong base or vice versa are combined

-at this point, pH = pKa (supported by Henderson hasselback equation)

Polyvalent titrations

Polyvalent acids and bases can donate or accept more than one H+ (diprotic acid has two hydrogens, two humps on graph)

-these curves usually feature multiple buffering regions and multiple equivalence points

Dynamic equilibrium

The forward and reverse reactions occur at the same rate, resulting in no observable change in the system

-no net change in concentrations of reactants or product

-reaction is constantly moving (dynamic)

-concentration of reactants or products are not always equal

Equilibrium constant (Keq)

aA + bB ←> cC + dD

Keq= rate of forward reaction/rate of reverse reaction = ([C]^c*[D]^d)/[A]^a*[B]^b)

*eckudes ant reactants or products that are liquid (l) or solid (s), only looks at gases and aqueous

Rules about equilibrium constants

1) if Keq < 1, then there is greater concentration of reactants than products at equilibrium

2) if Keq=1, then the ratio of products to reactants at equilibrium is equal

3) if Keq > 1, then there is a greater concentration of products than reactants

Keq to Gibbs Free Energy equation

Delta G (Gibbs Free energy) = -RTln(Keq)

R= a constant, T = temperature in Kelvin, Keq = equilibrium constant

Relationship between Keq and Gibbs Free Energy

1) if Keq < 1, ln(Keq) < 0, delta G is positive so reactants are favored at equilibrium

2) if Keq = 1, ln(Keq) = 0, delta G is 0, so ratio of products and reactants are equal

3) if Keq > 1, ln(Keq) > 0, delta G is negative so products are favored at equilibrium

Reaction Quotient

Uses concentration of the reactants/products at any point in the reaction when concentrations are still changing

Q = ([C]^c [D]^d)/([A]^a[B]^b)

Reaction Quotient (Q) vs Equilibrium Constant (Keq)

Reaction quotient looks at the concentration of products/reactants at any point in the reaction but equilibrium constant is only once the reaction reaches equilibrium

Rules regarding Reaction Quotient

1) if Q < Keq, there us a higher concentration is reactants than there is at equilibrium, and reaction proceeds forward

2) if Q=Keq, the reaction is in dynamic equilibrium

3) if Q > Keq, there is a higher concentration of products than there is at equilibrium, and the reaction proceeds in the reverse direction

Le Chatelier’s Principle

Shift by a system that occurs to restore the equilibrium state when it is disturbed

-affected by changes in concentration, pressure, and temperature

How does a change in concentration affect Le Chatelier’s Principle

1) Reactants are added: Q < Keq, so reaction shifts forward until Q=Keq

2) Reactants are removed: Q> Keq, so reaction shifts backwards until Q = Keq

3) Products are added: Q > Keq, so reaction shifts backward until Q = Keq

4) Products are removed: Q < Keq, so reaction shifts forward until Q = Keq

What happens tot the volume and pressure of a gas as it is compressed

Gas is compressed: decreases volume, increases collision, pressure increases

Gas is expanded: volume increases, collisions decrease, pressure decreases

How do changes in pressure and volume affect Le Chatelier’s Principle

Compression (increasing pressure and decreasing volume): reaction shifts towards the side with fewer moles of gas

Expansion (decreased pressure and increased volume): reaction shifts towards the side with more moles of gas

How does temperature affect Le Chatelier’s Principle

Endothermic reactions (delta H is positive): requires an input of energy so heat is considered a reactant

-decreasing temperature shifts reaction backwards, increasing temperature shift reaction forward

Exothermic reactions (delta H is negative): releases energy, so heat is considered a product

-decreasing temperature shifts reaction forward, higher temperature shifts reaction backwards

Precipitation reactions

Ions in a solution react to form an insoluble solid

-described using solubility product constant (Ksp)

Solubility Product Constant (Ksp)

Describes the extent to which an ionic compound dissolves (higher Ksp is more soluble)

*only looks at gases and aqueous components, NOT solids or liquids

Ksp = [product]^moles/[reactants]^moles

Solubility Product Quotient (Qsp)

Describes the current state of solution

Qsp = [products]^moles/[reactants]^moles

*only looks at gases and aqueous

-if Qsp < Ksp, reaction will proceed forward to make more product

Possible states for solutions

Unsaturated: Qsp < Ksp so solution can dissolve more solute

Saturated: Qsp = Ksp, so it cannot dissolve more solute

Supersaturated: Qsp > Ksp, so it cannot dissolve more solute and precipitate has formed

-solute concentration increases as you move from unsaturated to supersaturated

Molar solubility

Number of moles that can be dissolved per liter of solution until the solution becomes saturated

Common ion effect

Describes the decrease in solubility of an ionic precipitate when a soluble compound is added to the solution that shares an ion common to the precipitate

Example: adding AgCl to water solution with NaCl has lower solubility than just AgCl in water, so there’s more precipitate

Amphoteric species

Can act as an acid (proton donor) or base (proton acceptor)

-water is a very common example

Autoionization of water

Proton is transferred from one water molecule to another water molecule

-forms an H3O+ as the conjugate acid and OH- as the conjugate base

How does adding a strong acid or a strong base to an equilibrium reaction affect it

Adding a strong acid: strong acids completely dissociate, so it increases [H3O+], and equilibrium shifts left

Adding a strong base: strong bases completely dissociate, OH- reacts with H3O+ to reduce the [H3O+] so equilibrium shifts right

How does increasing the concentration of the acid in the reaction affect equilibrium? How does decreasing the concentration of the conjugate base of the acid affect equilibrium?

Increasing [HA]/acid concentration: equilibrium shifts forward/right

Decreasing [A-]/conjugate base: equilibrium shifts forward/right

How does increasing pH affect equilibrium

Increasing pH lowers the hydronic ion concentration [H3O]+ and makes it more basic

equilibrium shifts right to produce more [H3O]+

Oxidation vs Reduction

Oxidation: losing electrons (Fe2+ → Fe3+)

Reduction: gaining electrons (Cu2+ → Cu+)

*memorization tip: OIL RIG (oxidation is losing, reduction is gaining)

Oxidizing and Reducing Agents

Reducing Agent: reduces other species, is oxidized and loses electrons (common examples: H2, Fe, Zn, metals)

Oxidizing Agent: oxidizes other species, gains electrons and is reduced (common examples: O2, O3, H2SO4, halogens)

how to determine Oxidation States/Numbers

-atoms in their elemental state are assigned an oxidation number of 0 (H2, Li, etc.)

-monoatomic (one atom) ions are equal to their charge (Mg2+ = +2 oxidation state)

-Fluorine is -1, oxygen is often -2, hydrogen is usually +1

Half reactions

Part of a redox reaction that contains the reduction or oxidation components (an oxidation reaction and a reduction reaction combine for the complete redox reaction but these two components can be split into half reactions)

Balancing Redox reactions in acidic conditions

1) write out the half reactions and balance all atoms other than oxygen and hydrogen

2) balance the oxygen atoms by adding the same number of H2O molecules to the opposite side of the reaction

3) balance the hydrogen atoms by adding H+ ions to the opposite side of the reaction

4) if the charges on both sides are not equal, add electrons to the more positive side to balance charges

5) the number of electrons in both half reactions must be the same, if not multiple the coefficients to get them equal

6) add the half reactions together and cancel out the electrons and common terms

Balancing Redox Conditions in Basic Conditions

Same first few steps as balancing redox reactions in acidic conditions, but:

-after balancing the H+ atoms, for each H+ added, add the same number of OH- to both sides

-in the side where H+ ions were added, combine H+ and OH- to create H2On

-determine if the charges on both sides add equal. If not, ass electrons to the more positive side to balance both sides

Electrochemical cells

Convert chemical energy to electrical energy or vice versa using redox reactions and the movement of electrons

-batteries are electrochemical cells

What to remember about anode and cathode

An OX and a Red Cat

-anode: site of oxidation

-cathode: site of reduction

-anions (Cl-) flow to anode and cations (Na+) flow to cathode

*electrons always flow from anode to cathode

Chemical energy vs electrical energy

Chemical energy: produced from chemical reactions from energy stored in bonds of atoms/molecules

Electrical energy: produced from the flow of electrons

Galvanic (voltaic cells)

Spontaneous (negative delta G); example= battery

-electrons flow from anode (-) to cathode (+) spontaneously through the wire to produce electricity (does not require an external power source)

-this process converts chemical energy to electrical energy

-oxidation occurs at anode (loses electrons), reduction occurs at the cathode (gain electrons)

Eo (cell potential) is positive

Electrolytic cells

Nonspontaneous (positive delta G), example= rechargeable battery

-electrons flow from anode (+) to cathode (-) nonspontaneously requiring an external power source

-converts electrical energy to the chemical energy

-oxidation occurs at the anode (lose electrons), reduction occurs at the cathode (gain electrons)

Cell potential (Eo) is negative

Molten Electrolysis

Power source forces the nonspontaneous movement of electrons (-) to the cathode (+)

Electroplating

Process of playing one metal with another to cover it in a thin layer

-the metal being plated by the other metal is made into the cathode (-) which is the site of reduction

Eo (standard cell potential)

The sum of the oxidation potential and reduction potential measured in volts (V)

Eo cell = Eo reduction + Eo oxidation

-if Eo is positive, reaction is spontaneous (negative delta G)

-if Eo is negative, reaction is nonspontaneous (positive delta G)

What does standard mean/refer to

Concentration is 1M, partial pressure is 1 atm, temperature is 298K

How does standard potential affects strength if oxidizing and reducing agents

Stronger reducing agents (weak oxidizing agents): more negative reduction potential, want to lose electrons

Stronger oxidizing agents (weak reducing agents): more positive reduction potential, want to gain electrons

How to determine standard cell potential (Eo cell) in a galvanic cell

-write the oxidation and reduction half-reactions for the cell

-look up the reduction potential (Eo reduction) or the reduction half-reaction in a table

-balance the half-reactions and reverse the sign for the oxidation potential (Eo oxidation = -Eo reduction)

-add the potentials of the half-cells to get the overall standard cell potential

*half cell with higher reduction potential undergoes reduction, half cell with lower reduction potential is oxidized

-overall cell potential will be a positive value

4 Electrolysis variables

1) Coulomb (C) = SI unit of electric charge

2) Fardray’s constant = 96,500 C/mole e-

3) Ampere = coulomb/second

4) Volt = joule/coulomb

How to find moles of product using electrolysis calculations

moles of product = (I)(ts )/(n)(F)

-I: current (in amps)

-ts: time ( in seconds)

-n: number of moles of electrons transferred

-F: Fardray’s constant (96500 C/mol e-)

How to find change in free energy and cell potential using electrolysis calculations

Delta G (free energy) = -nFE

-n: number of moles of electrons transferred

-F: Fardray’s constant (96,500 C/mol e-)

-E: cell potential