Chapter 16: Acids and Bases

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Chapter 16: Acids and Bases

Students will learn… 3 definitions of Acids and Bases
Acid Strength – pH scale Water as acid and base Calculating pH of strong acids and bases Buffered solutions *Review naming acids, Pg 132 & 133

Three Definitions of Acids-Bases
Arrhenius BrØnsted-Lowry Lewis

3 Definitions of Acids and Bases
Arrhenius Produce H+ ion* (= proton) Produce OH– ion BrØnsted-Lowry Donate H+ ions Accept H+ ions Lewis Accept an electron-pair Donate an electron-pair *H+ ion always combines with water molecule, forming hydronium ion, H3O+

Arrhenius Acids: HCl, HC2H3O2, HCN Arrhenius Bases: NaOH, Ca(OH)2
BrØnsted-Lowry Acids and Bases: can only be determined from chemical equations H2O + HCl → H3O+ + Cl- base acid conjugate conjugate acid base conjugate pair conjugate pair

Conjugate Acid-Base Pair
An acid and a base that only differ by one H+ ion Always made of one reactant and one product Often in reversable reactions (Ex) HSO4- + H2O ↔ SO42- + H3O+ acid base base acid HSO4- + H2O ↔ H2SO4 + OH- base acid acid base

Examples 1. Which of the following represent conjugate acid–base pairs? a) HCl, HNO3 b) H3O+, OH– c) H2SO4, SO42– d) HCN, CN–

2. Write the conjugate base for each of the following.
HClO4 H3PO4 CH3NH3+

Strong vs. Weak Strong: 100% dissociation of H+ or OH–
Strong electrolytes (conduct electricity very well) Weak: Less than 100% dissociation of H+ or OH– Weak electrolytes (conduct electricity somewhat)

Common Strong and Weak Acids/Bases
Strong acids: H+7A, H2SO4, HClO4, HNO3 Weak acids: HCN, HC2H3O2 Strong bases: 1A+OH, Sr(OH)2, Ba(OH)2 Weak bases: NH3 (ammonia) in water = NH4OH (ammonium hydroxide)

Relative Strength of Acids and Bases

Example Acetic acid (HC2H3O2) and HCN are both weak acids. Acetic acid is a stronger acid than HCN. Arrange these bases from weakest to strongest : Cl– CN– C2H3O2–

Terms related Acids Amphoteric: Ions or molecules that can become an acid and a base (Ex) HSO4- + H2O → SO42- + H3O+ acid base HSO4- + H2O → H2SO4 + OH- base acid (Ex) H2O + H2O → H3O+ + OH- Oxyacids: acids containing oxygen (Ex) HNO3, H2SO4, HClO

Monoprotic Acids Diprotic Acids Triprotic Acids
Have one proton (H+) to donate (Ex) HCl, HNO3, HI, HClO3 Diprotic Acids Have more than two protons to donate (Ex) H2SO4 Triprotic Acids Have three protons to donate (Ex) H3PO4

Self-ionization of Water
In neutral water, H2O + H2O ↔ H3O+ + OH– [H3O+] = [OH–] = 1×10-7 mol/L [H3O+] · [OH–] = (1×10-7)·(1×10-7)=1× @ 25 °= Kw Kw = ion-product constant for water *The constant changes to the temperature

Acidic or Basic Solutions
Kw =1×10-14 always maintained [H3O+] [OH-] pure water [H3O+]=1×10-7 mol/L [OH-]= 1×10-7 mol/L acidic solution [H3O+]>1×10-7 mol/L [OH-]< 1×10-7 mol/L basic solution [H3O+]<1×10-7 mol/L [OH-]> 1×10-7 mol/L

Common Misunderstanding
An acidic solution has only H+ ions A basic solution has only OH‒ ions

Example (1) In an acidic aqueous solution, which statement below is correct? a) [H+] < 1.0 × 10–7 M b) [H+] > 1.0 × 10–7 M c) [OH–] > 1.0 × 10–7 M d) [H+] < [OH–]

(2) In an aqueous solution in which [OH–] = 2
(2) In an aqueous solution in which [OH–] = 2.0 x 10–10 M, the [H+] = ________ M, and the solution is ________. a) 2.0 × 10–10 M ; basic b) 1.0 × 10–14 M ; acidic c) 5.0 × 10–5 M ; acidic 5.0 × 10–5 M ; basic

pH and pOH Scale pH = –log[H+] [H+] = 10-pH pOH = ‒log [OH‒]
A compact way to represent solution acidity pH decreases as [H+] increases [H+] = 10-pH pOH = ‒log [OH‒] pOH decreases as [OH‒] increases [OH‒] = 10-pOH pH + pOH = 14

pH = 7; neutral pH > 7; basic Higher the pH, more basic. pH < 7; acidic Lower the pH, more acidic.

Acidity and pH Scale high pH = weak acid high pOH = weak base

Examples (1) Calculate the pH for each of the following solutions.
1.0 × 10–4 M H+ b) M OH–

(2) The pH of a solution is 5.85. What is the [H+] for this solution?

(3) Consider an aqueous solution of 2. 0 × 10–3 M HCl. What is the pH
(3) Consider an aqueous solution of 2.0 × 10–3 M HCl. What is the pH? *If the solution was 2.0 × 10–3 M HNO2, a weak acid, you can’t take the same approach. Why?

Buffered Solutions resists a change in its pH when either an acid or a base has been added. Presence of a weak acid and its conjugate base buffers the solution.

Characteristics of Buffered Solutions
The solution contains a weak acid HA and its conjugate base A–. The buffer resists changes in pH by reacting with any added H+ or OH– so that these ions do not accumulate. Any added H+ reacts with the base A–. H+(aq) + A–(aq) → HA(aq) 4. Any added OH– reacts with the weak acid HA. OH–(aq) + HA(aq) → H2O(l) + A–(aq)

Example If a solution is buffered with NH3 to which has been added NH4Cl, what reaction will occur if a strong base such as NaOH is added? a) NaOH + NH3 → NaNH4O b) NaOH + NH4+ → Na+ + NH3 + H2O c) NaOH + Cl– → NaCl + OH– d) NaOH + H2O → NaH3O2