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Chapter 17

Textbook Chapter 17 - Chem 1AA3

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Chem 1AA3 Chapter 17: Additional Aspects of Acid-Base Equilibrium Acid-Base Chemistry Review (Ch 16) Bronsted Lowry Theory +  Acid = H donor  Base = H acceptor  Acid/Base reactants and conjugate acid/base products:  H2O is an amphoteric species (behaves as both an acid or a base) Strong Acids/Bases  Strong Acids: o HCl o HBr o HI o HClO 4 o HNO 3 o H2SO 4  Strong Bases: o Group I and Group II Hydroxides (many Group II partially soluble), hydrides (H ) and oxides (O ) Water is Amphoteric Typical pKaValues Acid-Base Equilibrium  Weak Acid and Ka o Equilibrium favors reverse process:  Acidity of 3 O > CH3COOH Chem 1AA3  Basicity of CH 3OO > H O 2  Weak Base and K b o Equilibrium favours reverse process: -  Basicity of OH > NH 3  Acidity of NH 4 H O 2 Review of Relationships: pH and pK w  Water Auto ionization: H 2 (l) O 2 (l) H O3 +(aq) OH (aq) + - -14 K w [H O3][OH ] = 1.0x10 (25 degree °s C) o pK w pH + pOH = 14.0 Aqueous Solution  For a conjugate acid-base pair: o pK a pK = bK = 1w.0 Weak Acids/Bases: Assumptions pH of Salt Solutions  The pH of 0.1 NaCl (aq) acidic, basic or neutral? o NaCl (aq) Na +(aq) Cl(aq) + o Na (aq) H 2 (l) o Cl-(aq) H 2 (l)R o pH = 7 (neutral)  The pH of 01 M NH Cl 4 (aq) acidic, basic or  The pH of 01 M CH CO3Na (aq) acidic, basic or neutral? neutral? o NH C4 (aq) NH 4 (aq)+ Cl(aq) o CH 3OONa (aq) CH C3O -(aq) Na +(aq) - + o Cl (aq) H 2 (l)R o Na (aq) H 2 (l) o NH 4 (aq) H2O (l)H 3(aq)+ H3O +(aq) o CH 3OO -(aq) H2O (l)H COO3 (aq)+ OH -(aq) o pH < 7 (acidic) o pH > 7 (alkaline) Acid-Base Review  Bronsted-Lowry acid: Acid = H donor, Base = H acceptor  Conjugate acid-base pairs  Strong acids/bases: ~ 100% ionized in water  Weak acids/bases: equilibrium reaction with water o K and K vblues  pH of weak acid or weak base = equilibrium calculation  pH + pOH = 14 = pK + aK b  pH of salt: acidic, neutral or basic Chem 1AA3 Ch 17 Acid Base Reactions  Key to solving acid-base problems is: o Writing the correct chemistry o Determining which species acts as a base and which acts as an acid o Knowing what remains after the reaction has occurred  Looking for the limiting reagent o Also look at what remains and what has been formed Acid-Base Reactions  Consider 3 Options 1. Strong Base + Strong Acid 2. Strong Base + Weak Acid 3. Strong Acid + Weak Base 1. Strong Base + Strong Acid  Limiting regent: consider moles before and after reaction  pH = -log10H 3 }  [H3O ] = mol STRONG ACID remaining total solution volume 2. Strong Base + Weak Acid  pH is calculated as that of a weak base solution  pH is determined only from the STRONG BASE  NO 2 (aq)H 2 (l)H (aq) HNO 2(aq)  pH = 14 – pOH Chem 1AA3 3. Strong Acid + Weak Base  pH is calculated as that of a weak acid solution  pH is determined only from the STRONG ACID  NH 4 (aq) H2O(l)NH 3(aq) H3O+(aq) 17.4 Acid-Base Titrations  Limiting reactant questions: either acid or base is always in excess, except at the equivalence point* o *mol acid = mol base for 1:1 acid-base titration  In 3 titrations different species are formed as pH changes: o Strong Acid + Strong Base  Strong Acid  Neutral Salt  Strong Base o Strong Base + Weak Acid  Weak Acid  Buffer  Basic Salt  Strong Base o Strong Acid + Weak Base  Weak Base  Buffer  Acidic Salt  Strong Acid  Strong Acid–Strong Base Titration (Lab 1, 6)  Sample Calculations Chem 1AA3  Weak Acid– Strong Base Titration Chem 1AA3 o Features different from Strong Acid– Strong Base:  Start at higher pH  Buffer region  ½ equivalence point, pH = pKa  Equivalence point pH > 7  Titrations – Key Concepts o Equivalence point: mol acid = mol base (if 1:1 reaction( o Strong acid – Strong Base titration  pH = 7 at equivalence point  Sharp pH change at equivalence o Weak Acid- Strong Base / Strong Acid– Weak Base titrations  Buffer region  pH = pKa at ½ equivalence  pH ≠ 7 at equivalence (pH depends on the salt present – acidic or basic) 17.3 Acid-Base Indicators  Equivalence point: theoretical point where mol acid = mol base (1:1 titration)  With an indicator, we observe the endpoint (color change) Acid-Base Indicators  Color depends on pH -  Indicators are weak acids/bases (HIn, In ): o HIn (aq) H 2 (l) O 3 +(aq)+ In(aq) acid color base color Chem 1AA3 o K HIn= [In-][H3O] [HIn] o pH = pK HIn+ log[In-] [HIn]  We see: o Acid color when pH < (pK HIn – 1) (10-fold excess of HIn) o Base color when pH > (pK HIN+ 1) (10-fold excess of In )  Color change of ~ 2 pH
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