CHEM 162 - General Chemistry II
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CHEM 162 Exam 2 Review – Winter2015 page 1 of 3
CHEM 162: Exam 2 Study Guide
Chapter 17: Acids and Bases
• Recognize hydronium ion, H3O + = H+ + H2O
• Know the strong acids: HCl, HBr, HI, HNO3, HClO4, H2SO4.
• Know the common strong bases: LiOH, NaOH,
KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2. • Recognize strong acids and strong bases
dissociate or ionize (break up) completely. → Equilibrium lies far to the right.
• Recognize weak acids dissociate or ionize (break up) only to a small degree.
→ Equilibrium lies far to the left.
• Write balanced equations and equilibrium expressions for the dissociation of any acid. – Omit pure liquids and solids.
HA(aq) H+(aq) + A− (aq)
[HA]
]-[A ][H K a
+ =
• The strength of an acid is inversely related to
strength of its conjugate base. → Strong acids have conjugate bases that are
weaker than H2O. → Weak acids have conjugate bases that are
stronger than H2O. → Write net the ionic equation for the
weak acid’s conjugate base reacting with water to form the conjugate acid and OH−.
• Autoionization of water:
2 H2O(l) H3O +(aq) + OH−(aq)
Kw = [H +][OH–]=1.0×10–14 at 25°C
Kw = water’s ion product constant
• Know water is amphoteric, so it can act as an acid or a base.
pH scale: pH = 7: neutral & pH < 7: acidic & pH > 7: basic (or alkaline)
[H+] = 10−pH [OH−] = 10−pOH pH = −log [H+] pOH = −log [OH−]
pH + pOH = 14.00
Know # of sig figs in [H+] or [OH−] determines the # of decimal places in pH and pOH.
Calculate pH of Strong Acids and Strong Bases – Strong acids and bases ionize completely. – Accounting for all the H+ and OH– ions, → [H+] = original molarity of strong acid → [OH–] = molarity of base × (# of OH– in base)
pH Calculations for a Weak Acids – Weak acids remain mostly undissociated. – Write the equation for the dissociation of acid. – Set up ICE table with [HA] given or calculated. – Calculate pH of weak acid solution. – Use percent ionization of solution to get Ka.
– Use [HX] – x ≈ [HX] approx. for: [HA] x <5%
– Use Quadratic Method or method of successive
approximations for: [HA] x ≥0.05%
Percent ionization = [HA]
mequilibriu at ][H+ ×100%
Convert Ka and Kb: Kw = Ka ·∙ Kb = 1.0×10
–14
Acid-‐Base Properties of Salts – salt = ionic compound – Soluble salts dissociate into ions in water. – Classify a given salt as acidic, basic, neutral.
CHEM 162 Exam 2 Review – Winter2015 page 2 of 3
Chapter 17: Acids and Bases (Continued)
Acid-‐Base Properties of Salts (Continued)
Ions that produce acidic solutions – NH4+ and highly charged metal ions: Al3+, Zn2+, etc., except cations of strong bases
Ions that produce basic solutions – anions that are conj. bases of weak acids. – all anions except anions of strong acids – Note: SO4
2– + H2O → HSO4 – + OH–
Ions that produce neutral solutions – do not react with H2O to make H
+ or OH– Salts with an acidic cation and basic anion – Classify a given salt as acidic, basic, neutral. – If Ka>Kb, salt is acidic. – If Kb>Ka, salt is basic. – Account for different ion concentrations by
multiplying initial concentration with Ka or Kb. pH Calculations for Weak Bases – Write equation for reaction of base with H2O. – Write equilibrium expression for the weak base. – The weaker the base, smaller the Kb. – Set up ICE table with [A–] given or calculated. – Solve for x to calculate [OH–] and pOH of a weak base solution. Acid Rain – Explain why rain is naturally acidic (pH 5.6)
due to CO2 in the atmosphere and the formation of H2CO3(aq).
– Know the reactions for SO2 and NO2 with O2 and H2O in the atmosphere with rain to produce H2SO4 and HNO3 (or acid rain).
– Explain problems associated with acid rain.
Given == −+ ][OH ][H
1K 'w 1.0×10 14 at 25°C
– Know H+ (from strong acids) will completely react with any base, A–.
– Know OH– (from strong bases) will completely react with any acid, HA.
– Be able to write neutralization reactions for H+ or OH– added to any solution.
Polyprotic Acids – Write the stepwise dissociation of any polyprotic acid. – Recognize the [H+] is determined only by 1st dissociation step for most polyprotic acids. – Solve for pH and equilibrium ion
concentrations. The Common Ion Effect – The shift in equilibrium caused by the addition of a
salt with an ion in common with the dissolved substances
– Given a compound added to an acid or base at equilibrium, predict equilibrium shifts (to the left or right), if [H+] ↑ or ↓, if [OH–] ↑ or ↓, and if pH ↑ or ↓,
– Calculate the pH of an acid or base solution when a common ion is added.
Acid-‐Base Titrations – Distinguish between endpoint (when indicator
changes color) and equivalence point (when equal amounts of H+ or HA and OH– present)
– Know general pH range for equivalence point of following:
– strong acid-‐strong base titration (pH≈7) – weak acid-‐strong base titration (pH>7) – strong acid-‐weak base titration (pH<7) – Recognize that pH=pKa at halfway to
equivalence point for a weak acid-‐strong base titration.
– Given titration data, calculate the pH at any point during an acid-‐base titration.
− Write the net ionic equations for the neutralization reactions.
− Account for amount of H+ and OH– neutralized and the resulting initial [HA] or [A–]
− Set up ICE table then solve for x to get [H+] or [OH–]
− Calculate pH based on [H+] or [OH–] − Carry out titration calculations for polyprotic
acids or for salts containing the conjugate base of polyprotic acids.
CHEM 162 Exam 2 Review – Winter2015 page 3 of 3
Chapter 17: Acids and Bases (Continued)
Acid-‐Base Indicators – Know when color changes become apparent (at pKa±1). – Know effective range of indicators using pKa. – Given the pKa for different acid-‐base indicators, explain which indicators would be effective for
specific acid-‐base titrations. buffer: a solution of a weak acid or weak base and its conjugate that can resist large changes in pH upon
addition of a small amount of strong acid or strong base. Determining pH and [H+] for any buffered solution – Know pH=pKa for [HA]=[A
–] (at halfway to the equivalence point for a titration) – Calculate the number of moles of strong acid (H+) or strong base (OH–) added to the buffer. – Carry out the neutralization reaction, and determine the number of moles of weak acid/base and its
conjugate present after reaction with the strong acid (H+) or strong base (OH–). – Calculate the pH using the Henderson-‐Hasselbalch equation:
pH = pKa + log [HA]
]-[A
Know how to prepare a buffer system to maintain a given pH. – Choose a weak acid with pKa near buffer pH range.
– Use Henderson-‐Hasselbalch to determine the [HA]
]-[A ratio to get desired pH for a buffer system.
– Be able to determine the amount of weak acid (volume and concentration) and/or the mass of a salt containing the conjugate base to prepare a buffer to maintain the desired pH.
– Be able to determine how to prepare a buffer using a weak acid (volume and concentration) and a strong base or a salt containing the conjugate base of a weak acid and a strong acid.
Blood pH Buffer System – Know the chemical equation for the buffer system for blood pH. – Use Le Châtelier’s Principle to explain how the buffer system maintains the pH of blood. – Define acidosis and alkalosis and use Le Châtelier’s Principle to explain how the conditions occur.
The usual General Chemistry Periodic Table will be provided.
Be prepared to solve problems combining concepts from various chapters covered.
