Acids & Bases, Buffers, Titration

Binary Acids

Hydrogen and a nonmetal

• Example: HCl, H2S, HF, etc.

Binary Acids: Gas Molecules

Hydrogen ide

• Ex: hydrogen chloride, hydrogen sulfide, hydrogen fluoride

Binary Acids: Dissolved in Water

Hydro ic acid

• Ex: hydrochloric acid, hydrosulfuric acid, hydrofluoric acid

Oxoacids

Contains hydrogen, oxygen, and other elements

• Oxygen normally combines with another nonmetal for a polyatomic anion

• Draw out lewis structures of each anion and the # of charges correlates to how many H’s you’ll add in the name

Bases

Ionic or molecular compounds

• May contain hydroxide ions (NaOH, Ba(OH)2, KOH, etc

• Molecular compound; bases due to their molecular structure (like NH3)

Arrhenius Definition

Based on H+ and OH-

• Acids are substances that produce H+ in aqueous solutions

  • Do not exist alone, combine with water to form H3O+

• Bases produce OH-

Bronsted Lowry Definition

• The acid is capable of donating a proton (H+)

• The base is capable of accepting a proton

Bronsted Lowry Example

HCl in water

• HCl(aq) + H2O → Cl- + H3O+

• HCL is the acid because it transfers a proton to water, forming H3O+

Conjugate Pairs

• A base accepts a single proton and becomes a conjugate acid

• An acid donates a single proton and becomes a conjugate base

• Going from neutral molecules to charged molecules

Amphoteric Substances

Can act as an acid and a base because they have both a transferable H and an atom with a lone pair of electrons

• Example: water

Autoionization of Water

H2O + H2O ⇌ H3O+ + OH-

• Water is amphoteric; can act as a base and an acid

• Concentration of H3O+ and OH- are equal in DI water

Ion Product of Water

The product of the H3O+ and OH- [ ] is always the same regardless of the composition of the solution at 25C

• [H3O+] x [OH=] = Kw = 1.00 × 10^-14

• H3O+ and OH= are inversely proportional (if one increases, the other must decrease so so product stays constant

Neutral Solutions

Equal [H3O+] and [OH-]

• 1.00 × 10^-7

Acidic Solutions

Higher [H3O+] than [OH-]

• [H3O+] > 1.00 × 10^-7

• [OH-] < 1.00 × 10^-7

Basic Solutions

Higher [OH-] than [H3O+]

• [H3O+] < 1.00 × 10^-7

• [OH-] > 1.00 × 10^-7

pH

The acidity or basicity of a solution

• pH = -log[H3O+]

• pH_water = -log[10^-7] = 7

• B/c of the - in front of the log, when [H3O+] increases, pH decreases and vice versa

pH Scale

• Lower pH = more acidic

• High pH = more basic

• It is possible to have a pH less than 0

pOH

Another way of expressing pH

• pOH = -log[OH-]

• pH + pOH = 14.00

Strong Acid

Donate practically all of their hydrogens to water

• 100% ionized in water

• Strong electrolyte

Strong Acid List

• Hydrochloric acid (HCl)

• Hydrobromic acid (HBr)

• Hydriodic acid (HI)

• Nitric acid (HNO3)

• Perchloric acid (HClO4)

• Sulfuric acid (H2SO4)

Strong Bases

The stronger the base, the more willing it is to accept H+

• Practically all units are dissociated into OH- or accept protons for ionic bases

• Dissolved ionic bases: 100% ionized in water

• Strong electrolyte

Examples of Strong Bases

• NaOH

• KOH

• LiOH

• Ca(OH)2

• Ba(OH)2

Hydronium and Hydroxide in SA or SB Solutions

• Two sources of hydronium in a SA solution: the acid and the water

• Same for SB: the base and the water

• The contribution of the water to the total concentration of H3O+ or OH- is negligible

Weak Acids

Donate a small fraction of the hydrogen’s

• Most weak acid molecules do not donate H+ to water

• % ionization is usually low, <10%

• Acid dissociation constant (Ka)

Weak Bases

Small fraction of the substance dissolved will accept protons from H2O

• Base ionization constant (Kb)

• Most weak base molecules do not react with water

Strengths of Acids and Bases

Measured by determining the equilibrium constant of a substances reaction with water

• The farther the equilibrium lies to the products, the stronger the acid or base

pK

A way of expressing the strength of an acid or base (using Ka or Kb)

• pK = -log(K)

• K = 10^-pK

• pK has an inversely proportional relationship with the strength of the acid or base

Conjugate Pairs and Kw

When you add two equations together (an acid dissociating and a base reacting with water), you multiply the K’s

• Ka x Kb = [H3O+][OH-] = Kw

Strength Table Trends

• Top: Strong acids have very weak conjugate bases

  • Complete acid ionization; no base ionization

  • Water would rather react with itself than these bases

• Middle: Weak acids with weak conjugate bases

• Bottom: Strong bases with very weak conjugate acids

  • Complete base ionization; no acid ionization

Predicting A/B Reactions

A product favored reaction must have weaker products than reactants

• Ex: SA with a SB produce a very weak base and very weak acid; so K »1

• SA with a WB or WA and SB produces a similar result with K » 1

• WA and WB reactions can go either way, so use the Ka and Kb values of the products and reactants to see which is weaker

Finding Keq of a Reaction

Keq = Ka x Kb x 1/Kw

• Remember, to find Kb you divide Ka out of Kw

Acidic/Basic Nature of Salts

If the Ka > Kb, it is an acidic solution and vice versa

• If its just a salt that dissociates and doesn’t react with water, the solution is neutral (no H3O+ or OH- formed)

Acidic Hydrogens

Hydrogen that can be donated as a proton in an acid base reaction

Monoprotic Acids

Only one acidic hydrogen can be donated

• Ex: HCl, HBrO4, HNO2, etc.

Polyprotic Acids

More than one acidic hydrogen can be donated

• Diprotic: H2SO4, H2S, H2C2O4, etc.

• Triprotic: H3PO4, etc.

Finding pH of Strong Polyprotic Acids

Two step process (H2SO4 example)

• First step: complete ionization of H2SO4 to produce H3O+ and HSO4- ions

• Second step: partial dissociation of the weak HSO4- acid (make an ICE table)

Finding pH of Weak Polyprotic Acids

Three-step process

• 1st: partial dissociation into hydronium and H2PO4-

• 2nd: partial dissociation into HPO42-

• 3rd: partial dissociation into PO43-

• Find the Ka each time and make an ICE table each time (same process for a weak polyprotic base)

Leveling Effect of Water

All strong acids and bases are equally strong in water because H2O exerts a leveling effect

• Must let two acids react with each other without including water to determine which is stronger

Strength of Binary Acids

The stronger the H-X bond, the weaker the acid

• Binary acid strength increases down a group

• Acid strength ex: HF < HCl < HBr

Conjugate Base Strength of Binary Acids

Opposite trend of the binary acids

• The ion with a greater charge density has a stronger attraction to H, meaning it’s more willing to accept another H

• Base strength ex: HF > HCl > HBr

Strengths of Oxoacids

The more electronegative the Y atom, the stronger the oxoacid (H-O-Y)

• Acidity of the oxoacid decreases down a group

• Acid strength ex: HOCl > HOBr > HOI

Conjugate Base Strength of Oxoacids

The more electronegative the Y atom is, the more it attracts electrons away from the oxygen

• Has a weaker attraction to H

• The charge is delocalized towards the ion, not the oxygen, so it’s less likely to accept a proton

• Base strength example: HOI > HOBr > HOCl

Strength of Oxoacids: Structure

The more oxygens attached to Y, the stronger the oxoacid

• Example: HClO4 > HClO3 > HClO2 > HClO

Conjugate Base Strength of Oxoacids: Structure

More oxygens → more resonance structures → more delocalization → weaker attraction to H+

• Example: ClO4- is a weaker base than ClO3-

Buffers

Resist changes in pH by neutralizing an acid or base that is added

• Contain either significant amounts of a WA and its conjugate base or a WB and its conjugate acid

Buffers: Process of Adding an A or B

• Before addition: the weak acid is in equilibrium with its conjugate base

• Adding a base: base will react with the weak acid; HA is consumed and A- is produced

• Adding a acid: acid reacts with the weak base; HA is produced and A- is consumed

• Use ICE table for each step!

Finding pH Changes: Adding a Base

Adding a base; two reactions involved

• Strong base reacting with the weak acid

• The weak acid then dissociating

ICE Tables for pH of Buffers

• Before adding OH-: can just use Henderson-Hasselbalch eqtn if significant amount of the weak acid and its conjugate base are present

• If the concentrations of the WA and SB are the same, you can just use the equation and skip the ICE table calculations

• Addition of OH-: rxn between WA and SB

  • Must use moles

  • Find the LR and subtract or add those moles to the other molecules' initial [ ]

Buffer Range

The range of pH values over which a buffer is most effective

• Most effective: the [ ] of the WA and its conjugate base are very similar

  • pKa = pH

• Non-effective: if either the WA or its conjugate base is less than 10% of the other

  • pH = pKa + 1 or pH = pKa - 1

Buffer Capacity

The amount of a strong acid or base that can be added to a given volume of the buffer before the pH changes significantly

• The greater amounts (# of moles) of the WA and its conjugate base (per given volume of buffer), the higher the buffer capacity

• The more it has, the more it can resist change

Titration

Type of quantitative analysis that can determine the amount and/or concentration of a substance

Titrant

Solution of a substance with known concentration

Analyte

Substance whose amount/concentration is unknown

• If the acid is the analyte, then the base is the titrant (and vice versa)

Titration Chart

“Zones”

• Before titration: HA + H2O ⇌ A- + H3O+

• Buffer region: [HA] > [A-], then after halfway point [A-] > [HA]

• Half-way point: [HA] = [A-]; pH = pKa

• Equivalence point: [OH-] = [HA]; A- + H2O ⇌ H2O + OH-

• After equivalence point: [OH-] > [HA] when there’s excess base