AP Chemistry Unit 8:
Acids and Bases

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8.1 Introduction to Acids and Bases

• Arrhenius Acids/Bases

• In acid-base or neutralization reactions, an Arrhenius acid and base will usually react to form water and salt.
• Arrhenius Acids

• Any species that increases the concentration of H+

• H+ represents the same thing as Hydronium (H3O+) because of H+ 's immediate association with H2O

• Arrhenius Bases

• Any species that increases the concentration of OH–

• Example (in water):
  HCl → H+ (aq) + Cl– (aq)
  Type - Arrehnius Acid

• Brønsted-Lowry Acids/Bases

• Acids donate a proton
• Bases accept a proton

• Conjugate Pairs

• Conjugate Pairs are at equilibrium with competition for H+. The stronger base controls direction.
• Example: HA (aq) + H2O (l) ⇌ H3O+ (aq) + A– (aq)
                  Acid                Base              Conj Acid         Conj Base
• Conjugate Acid

• Species formed when the proton is transferred to the base
• Example: Hydronium (always a conjugate acid)

• Conjugate Base

• Everything that remains of the acid molecule after a proton is lost
• Species form after an acid donates its proton.

• Acids

• Monoprotic Acids

• One acidic Hydrogen
• Example: HCl

• Polyprotic Acids 

• More than one acidic hydrogen
• Diprotic Acids (H2PO3)

• Acid with two acidic hydrogens

• Triprotic Acids (H3PO4)

• Acid with three acidic hydrogens

• Oxyacids

• Acidic hydrogen is attached to the oxygen of an ion
• The more oxygens in an oxyacid, the stronger the acid.
• Example: HPO3
  

• Tastes Sour

• Bases

• Regular Naming
• Tastes Bitter

• Amphoteric (Amphiprotic)

• Water is Amphoteric, meaning it can behave as an acid or base.
• Representation:
  H2O + H2O ⇌ H3O+ + OH–

• pH Scale

• pH generally ranges from 0-14 but can be less and greater than this range
• Sig figs for pH only count the digits after the decimal place.
  Example: 123456.00 (2 sig figs)
• The pH scale is a logarithm scale based on 10.
  Thus, pH = –log[H+]
• Acidic vs. Basic
  Source: Abigail Giordano
• Water vs. Kw

• The chart above has exceptions. Use the equations below (if possible) as explanations of whether something is neutral, acidic, or basic.
• The equilibrium constant for the autoionization of water (Kw) is always 10–14 at 25°C
• At 25°C in pure water, [H+] = [OH–] = 1.0•10-7 M
  This means that Kw = [H+][OH–] = 1.0•10-14 M
• Neutral

• [H+] = [OH–]

• Acidic

• [H+] > [OH–]

• Basic

• [H+] < [OH–]

• pH Equations (Given on the equation sheet)

• Litmus Paper

8.2 pH and pOH of Strong Acids and Bases

• pH

• The concentration of H+ ions
• Lower concentration of H+ → More Basic
• Higher concentration of H+ → More Acidic

• pOH

• The concentration of OH– ions
• Lower concentration of OH– → More Basic
• Higher concentration of OH– → More Acidic

• Ka

• Larger Ka → Stronger Acid
• Lower Ka → Weaker Acid
• Generic monoprotic weak acid HA with conjugate base A– ‘s equilibrium constant
  

• Kb

• Larger Kb → Stronger Base
• Lower Kb → Weaker Base
• Generic weak base B with conjugate acid BH+ ‘s equilibrium constant
  

• Strong Acids/Bases

• Dissociates completely in water (essentially goes to completion)
• No equilibrium or dissociation constant
• Strong Acid + Strong Base = Water (Neutral pH)

• Strong Acids

• [H+] = [HA]
• [OH–] will be very small (equilibrium)
• If the molarity of the strong acid is very dilute (<10–7), water would be the dominant acid (pH of 7.00).

• Strong Bases

• Dissociates completely in water
• When comparing strong bases, the ones with 2 OH molecules will be more basic (and thus a higher pH value).

8.3 Weak Acid and Base Equilibria

• Weak Acids (Small Ka/High Kb)

• If the acid isn’t strong, then they are weak.
• Weak acids & bases do not dissociate completely. The strength depends on how much it dissociates.
• Equilibrium is far to the left (very few products)
  
  Source: Abigail Giordano
• Weak acids produce Hydronium, but only a very small percent of the acid dissociates.

• [H+]<<[HA]

• Weak Base (Small Kb/High Ka)

• Weak bases produce OH–, but only a very small percent of the base dissociates.

• [OH–]<<[HB+]

• Percent Dissociation (Ionization)

• % dissociation = amount dissociated (M)/initial concentration (M) • 100%
• For weak acids, percent dissociation increases as an acid becomes more dilute.

8.4 Acid-Base Reactions and Buffers

• Buffer

• It must contain a common ion:

• A weak acid and its conjugate base present in some form
  Example: HF & F–
• A weak base and its conjugate acid present in some form
  Example: NH3 & NH4+

• Buffers resist changes in pH
• H+ is the strongest acid and OH– is the strongest base.

8.5 Acid-Base Titrations

• Titration

• A lab procedure using a burette to drip small amounts of the titrant (solution of known concentration) into the analyte until the equivalence point is met.
  

• Titration (pH) Curve

• pH changes gradually until the titration is near its equivalence point, where it changes more dramatically. This is because there is a large amount of H+ or OH–, and adding the acid/base would only produce a small change.
• Equivalence Point

• When enough titrant has been added to react exactly with the solution being analyzed, it is when the graph’s slope is highest.
  Mol acid = mol base
  
  Source: chemistrystudent

• Half-Way Point (Half Equivalence Point)

• When half of the analyte has been neutralized

• Weak vs Strong Acid/Base Titrations

• Weak acids/bases will be affected more at the start of titrating.
• Strong acids (pink) will barely be affected and the equivalence point will be ~7.
  
  Source: okstate.edu

8.6 Molecular Structures of Acids and Bases

• Bond Strength (bond dissociation energy)

• Weaker → Stronger Acid
• Stronger → Weaker Acid

• Bond Polarity (electronegativity)

• Lower → Stronger Acid
• Higher → Weaker Acid

• Oxyacids

• The hydrogen for that acid is always bonded to oxygen.
  Example: X–O–H
•  When electrons are pulled to the side of the electronegative elements (X), the opposite side (O–H bond) will become weaker.
  
  Source: Abigail Giordano
• The more oxygens (highly electronegative), the more the electron density will be pulled towards its side, away from the hydrogen.
• When comparing Oxyacids with the same number of oxygens, the oxyacid with the most electronegative element will be the stronger acid.

8.7 pH and pK

• Acid Dissociation Constant

• pKa = –log(Ka)

• Base Dissociation Constant

• pKb = –log(Kb)

• Titration Curve

Source: Quizlet

• pH < pKa (Before buffer region)

• Acid > Conjugate base

• pH > pKa (After buffer region)

• Acid < Conjugate base

• pH = pKa (Half-equivalence point)

• Acid = Conjugate base

• Indicators

• Indicators are substances that change color at a certain pH level.
• Example: colorless in acidic solution, but pink in basic solution.
  
  Source: wisc.pb.unizin.org
• Effective Range

• When an indicator’s pKa is ±1 to a solution’s equivalence pH.

• Phenolphthalein is a commonly used indicator because its effective range is 8 to 10.
• For the AP Exam, students will not need to memorize speicifc indocators or their effective ranges, but will be asked to choose the most effective indicator for a certain experiment (given a list of different options).

8.8 Properties of Buffers

• Buffer

• It must contain significant amounts:

• A weak acid and its conjugate base present in some form
  Example: HF & F–
• A weak base and its conjugate acid present in some form
  Example: NH3 & NH4+

• Buffers resist changes in pH
• H+ is the strongest acid and OH– is the strongest base.
• Most effective when the pH is similar to pKa

8.9 Henderson-Hasselbalch Equation

• Henderson-Hasselbalch Equation

• pH = pKa + log [A–]/[HA] = pKa + log [base]/[acid]
• Only to be used for buffer systems.
• When the acid and conjugate base are equal in molarity, then pH=pKa

8.10 Buffer Capacity

• Buffer Capacity

• The maximum amount of acid or base that can be added to a buffer without causing a significant change in pH.
• To increase a buffer's capacity, the concentrations of the conjugate acid base pair must be increased in the same ratio.
• Example: If there are 3 moles of Weak Acid, adding 3 moles of strong base would use up all the weak acid, ending the buffer.

8.11 pH and Solubility

• Solution’s pH Effect on Solubility

• Basic Salt (Example: Fe(OH)3 ⇌ Fe3+ + 3OH–)

• Dissolving in high pH (basic solution)

• Basic solutions have excess OH– and adding more OH– to the equation shifts it in the reverse direction, decreasing solubility by producing more salt.
• Decreases Solubility (Q>Ksp)

• Dissolving in low pH (acidic solution)

• Acidic solutions have excess H+ which will react with the OH– in the basic salt, causing the equation to shift in the forward direction, increasing solubility by using more salt.
• Increases Solubility (Q<Ksp)

• Acidic Salt (Example: NH4+ ⇌ NH3 + H+)

• Dissolving in high pH (basic solution)

• Basic solutions have excess OH– which will react with the H+, causing the equation to shift in the forward direction, increasing solubility by using more salt.
• Increases Solubility (Q<Ksp)

• Dissolving in low pH (acidic solution)

• Acidic solutions have excess H+ and adding more H+ to the equation shifts it in the reverse direction, decreasing solubility by producing more salt.
• Decreases Solubility (Q>Ksp)