Acids And Bases Chemistry Ms. Piela

Key Characteristics of Acids & Bases Taste sour Reacts with alkali metals (forms H2 gas) Forms electrolyte solutions (conducts electricity) pH paper color: Red Neutralizes Bases Bases Tastes bitter Slippery feel pH paper color: Blue Neutralizes Acids

The 3 Main Theories of Acids/Bases Lewis Acids/Bases This course will mainly deal with BL theory Bronsted-Lowry Acids/Bases Arrhenius Acids/Bases

Theories of Acids & Bases Arrhenius Theory of Acids & Bases: Properties of acids are due to the presence of H+ ions Example: HCl  H+ + Cl- Properties of bases are due to the presence of OH- ions NaOH  Na+ + OH-

HCl(g) + H2O(l) H3O+(aq) + Cl-(aq) What is an H+? H+ ions are bare protons These are so reactive that they do not exist naturally, but will bond with water to form a hydronium ion, or H3O+ ion Oftentimes H+ and H3O+ are used interchangeably HCl  H+ + Cl- HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)

Problems with the Arrhenius theory Only deals with aqueous solutions (solutions in water) Not all acids and bases contain H+ and OH- ions Example: NH3 is a base Considered the most incomplete theory of acids and bases

Theories of Acids & Bases Brønsted-Lowry Theory of Acids & Bases Acids are substances that donate H+ ions Acids are proton (H+) donors Bases are substances that accept H+ ions Bases are proton (H+) acceptors Example: HBr + H2O  H3O+ + Br- A B

Brønsted-Lowry Theory The behavior of NH3 can be understood now: NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq) NH3 becomes NH4+, so NH3 is a proton acceptor (or a Brønsted-Lowry base) H2O becomes OH-, so H2O is a proton donor (or a Brønsted-Lowry acid)

Brønsted-Lowry Theory

Brønsted-Lowry Theory Conjugate Acid-Base Pairs Definition: An acid and a base that differ only in the presence or absence of H+ Every acid has a conjugate base. Every base has a conjugate acid. These pairs only ever differ by exactly one hydrogen ion

Brønsted-Lowry Theory Example Problems Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base NH3 + H2O  NH4+ + OH- B A CA CB

Brønsted-Lowry Theory Example HCl (g) + H2O (l) ↔ H3O+(aq) + Cl- (aq) HSO4- + HCO3- ↔ SO4-2 + H2CO3 A B CA CB A B CB CA

Theories of Acids & Bases Lewis Acids & Bases Acids are electron acceptors Bases are electron donors Amphoteric – substances that can act as both an acid and a base Examples: H2O, HCO3-

Summary Of Theories Acids release H+ Bases release OH- Arrhenius Acids release H+ Bases release OH- Brønsted-Lowry Acids – proton donor Bases – proton acceptor Lewis Acids – electron acceptor Bases – electron donor

The pH scale Developed by Søren Sørensen in order to determine the acidity of ales Used in order to simplify the concept of acids and bases for his workers The pH scale goes from 0 to 14 The acidity/basicity of the solutions depends on the concentration of H+ (or H3O+)

The pH scale Acidic pH < 7 Neutral pH = 7 Basic pH > 7

pH scale Low pH values means a high concentration of H+ (acidic) High pH values means a low concentration of H+ (basic)

H2O (l) ↔ H3O+ (aq) + OH- (aq) Calculations of pH The Self Ionization of Water In pure water (pH = 7), the concentrations of the ions (H3O+ and OH-) are equal. [H3O+]=[OH-]= 1x10-7 This is because water will spontaneously dissociate naturally: H2O (l) ↔ H3O+ (aq) + OH- (aq) Writing the equilibrium expression for the self-ionization of water gives:

The Self-ionization of Water Plugging in the concentrations in pure water, this gives an equilibrium constant of 1x10-14 This is referred to as the ion product constant of water The ion product constant of water has its own symbol: Kw Unlike other equilibrium constants, the Kw will always be the same value

Calculations of H3O+/OH- Example #1 What is the H3O+ concentration in a solution with [OH-] = 3.0 x 10-4 M? Kw = [H3O+][OH-] 1 x 10-14 = [H3O+][3.0 x 10-4] ______ ___________________ 3.0 x 10-4 3.0 x 10-4

Calculations of H3O+/OH- If the hydronium-ion concentration of an aqueous solution is 1.0 x 10-3 M, what is the hydroxide ion concentration in the solution?   Kw = [H3O+][OH-] 1 x 10-14 = [1 x 10-3][OH-] ______ ___________________ 1.0 x 10-3 1.0 x 10-3

pH = -log [H3O+] or [H3O+] = 10-pH Calculations of pH pH can be expressed using the following equation: pH = -log [H3O+] or [H3O+] = 10-pH Example #1 What is the pH of a solution with 0.00010 M H3O+? Is this solution an acid or a base? Acid!

Calculating pH of a solution Example #2 What is the pH of a solution where the concentration of hydroxide ions is 0.0136 M? Is this an acid or a base? Kw = [H3O+][OH-] pH = -log [H3O+] Base!

Calculating pH of a solution Practice #1 Practice #2

Calculating H3O+/OH- from pH Example #1 What is the hydronium ion concentration in fruit juice that has a pH of 3.3? [H3O+] = 10-pH

Calculating H3O+/OH- from pH What are the concentrations of the hydronium and hydroxide ions in a sample of rain that has a pH of 5.05? [H3O+] = 10-pH Kw = [H3O+][OH-]

Calculating H3O+/OH- from pH Practice #1 Practice #2

Strength of Acids & Bases When a solution is considered strong, it will completely ionize in a solution Nitric acid is an example of strong acid: HNO3 (l) + H2O (l) ⇋ NO3- (aq) + H3O+ (aq) In a solution of nitric acid, no HNO3 molecules are present! Strength is NOT equivalent to concentration!

Strength of Acids & Bases Knowing the strength of an acid is important for calculating pH If given concentration of strong acid (such as HNO3) assume it is the same as the concentration of hydronium, H3O+, ions Given concentration of a strong base, assume it has the same concentration as the hydroxide, OH-, ions

Strong Acids & Bases Ionize 100% Example NaOH  Na+ + OH- 1 M 1 M 1 M OH- Na+ Na+ Na+ OH- OH-

Weak Acids & Bases Ionize X% Example HF  H+ + F- 1 M ? M ? M F- HF H+ H+ HF H+ F- F-

Naming Bases Bases are soluble metal hydroxides Examples Follow identical naming rules for ionic compounds Examples NaOH Ba(OH)2 NH3 NH4+ Sodium hydroxide Barium hydroxide Ammonia Ammonium

Naming Acids Binary Acids (HX) If the acid has an anion that ends in “-ide” use the following basic format to name the acid: “Hydro – root – ic acid” Example HCl Hydrochloric acid

Naming Acids Example Practice HBr HI H2S Hydrobromic acid Hydroiodic acid Hydrosulfuric acid

Naming Acids Polyatomic acids (aka oxoacids, HxAyOz) Name depends on the polyatomic used: If polyatomic ends in “-ite”, replace with “ous acid” If polyatomic ends in “-ate”, replace with “ic acid” Trick: “I ate something icky”

Naming Acids Examples HClO4 HClO2 Sulfuric acid Perchloric acid Chlorous acid H2SO4