Textbook Notes (367,969)
Canada (161,538)
Chemistry (261)
CHEM 1AA3 (18)
Pippa Lock (13)
Chapter 17

Textbook Chapter 17 - Chem 1AA3
Premium

11 Pages
86 Views
Unlock Document

Department
Chemistry
Course
CHEM 1AA3
Professor
Pippa Lock
Semester
Winter

Description
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
More Less

Related notes for CHEM 1AA3

Log In


OR

Join OneClass

Access over 10 million pages of study
documents for 1.3 million courses.

Sign up

Join to view


OR

By registering, I agree to the Terms and Privacy Policies
Already have an account?
Just a few more details

So we can recommend you notes for your school.

Reset Password

Please enter below the email address you registered with and we will send you a link to reset your password.

Add your courses

Get notes from the top students in your class.


Submit