Acids and Bases Properties, Reactions, and Calculating pH

Acids and Bases Properties, Reactions, and Calculating pH

Acids and Bases Properties, Reactions, and Calculating pH Standards **STANDARD SET 5: Acids and Bases 5. Acids, bases, and salts are three classes of compounds that form ions in water solutions. As a basis for understanding this concept: a. Students know the observable properties of acids, bases, and salt solutions. 5. b. Students know acids are hydrogen-ion-donating and bases are hydrogen-ionaccepting substances.

5. c. Students know strong acids and bases fully dissociate and weak acids and bases partially dissociate. 5. d. Students know how to use the pH scale to characterize acid and base solutions. 5. e.* Students know the Arrhenius, Brnsted-Lowry, and Lewis acidbase definitions. 5. f.* Students know how to calculate pH from the hydrogen-ion concentration. 5. g.* Students know buffers stabilize pH in acidbase reactions. Properties Acids Electrolytes React with metals to

produce H2 gas Form H3O+ ions in water Tastes sour Bases slippery Form OH ions in water Tastes bitter Examples: soap, cocoa Examples: vinegar, lemon powder, ammonia,

juice, battery acid Drano Calculating pH [H3O ] = 10 + pH H+ concentration= 1 1012

pH = 12 pH Scale pH 1 Strong acids Weak acids 7 Neutral : water Weak bases

14 Strong bases pH < 7 = acidic pH = 7 = neutral pH > 7 = basic Various Definitions of Acids/Bases Arrhenius Acid a substance which forms

hydronium (H3O+) ions in water. Arrhenius Base a substance which forms hydroxide (OH) ions in water. Brnsted-Lowry Acid a substance which can donate a proton (an H+). Brnsted-Lowry Base a substance which can accept a proton (an H+). Lewis Acid a substance which can accept a lone pair of electrons. Lewis Base a substance which can donate a lone pair of electrons.

Acid Reactions Strong acid: HCl (aq) + H2O (l) Cl (aq)+ H3O+ (aq) hydronium ion Weak acid: CH3COOH(aq) +H2O(l)

CH3COO(aq)+ H3O+(aq) Base Reactions Strong base: NaOH (aq) + H2O (l) Na+ (aq)+ H2O + OH(aq) hydroxide ion Weak base: NH3 (aq) + H2O(l)

NH4+(aq)+ OH(aq) Conjugate Acid/Base Pairs CH3COOH(aq) +H2O(l) acid CH3COO(aq)+ H3O+(aq) conjugate base

NH3 (aq) + H2O(l) NH4+(aq)+ OH(aq) base conjugate acid Conjugate Acid/Base Pairs CH3COOH(aq) +H2O(l)

acid CH3COO(aq)+ H3O+(aq) conjugate base NH3 (aq) + H2O(l) NH4+(aq)+ OH(aq) base

conjugate acid Conjugate Acid/Base Pairs CH3COOH(aq) +H2O(l) H+ acid CH3COO(aq)+ H3O+(aq) conjugate

H+ base Conjugate Acid/Base Pairs NH3 (aq) + H2O(l) + H base NH4+(aq)+ OH(aq) +

H conjugate acid Water is an Acid/Base H2O (l) + H2O (l) H3O+ (aq) + OH(aq) hydronium ion

hydroxide ion But water is a weak acid and a weak base, so only a small amount of water will form these ions, the rest will remain as H2O. Measuring the Strength of Acids/Bases We often speak of acid concentrations in molarity (ex. 2.0 M HCl solution, aka

) 2.0 mol HCl 1 L solution But which is stronger 2.0 M HCl solution or 2.0 M CH3COOH solution? Its more important to know the concentration of hydronium ions. [H+] really means [H3O+]

Calculating pH pH = log10[H3O+] Example: If [H3O+] = 1 x 105 M, what is the pH? pH = log10[1 x 105] pH = 5 Back-calculating concentration (1) x log10[H3O+] = pH x (1) log10[H3O+] = pH 10

10 [H3O+] = 10pH H2O (l) + H2O (l) H3O+ (aq) + OH(aq) hydronium ion

hydroxide ion H3O+ (aq) + OH(aq) H2O (l) + H2O (l) H2O (l) + H2O (l) H3O+ (aq) + OH(aq)

H3O+ (aq) + OH(aq) H2O (l) + H2O (l) neutral solution H3O+ (aq) + OH(aq) H2O (l) + H2O (l) acidic solution

H3O+ (aq) + OH(aq) H2O (l) + H2O (l) basic solution Calculating [OH] from [H3O+] H2O (l) + H2O (l) H3O+ (aq) + OH(aq)

Kw = [H3O+][OH] Kw = 1 x 1014 always this number Example: If [H3O+]= 1 x 105M, what is [OH]? 1 x 1014 = [1 x 105M][OH] [1 x 105M] [1 x 105M] 1 x 109 M = [OH]

pH Scale pH 1 Strong acids Weak acids 7 Neutral : water Weak bases 14 Strong bases

pH < 7 = acidic pH = 7 = neutral pH > 7 = basic Neutralization H3O+ (aq) + OH(aq) H2O (l) + H2O (l) Large amounts of H3O+ and OH cannot exist at the same time in a solution. Whichever

ion has the larger amount will reduce the lesser amount. What you would see: HCl (aq) + NaOH (aq) NaCl (aq)+ H2O (l) Titration H3O+ (aq) + OH(aq) H2O (l) + H2O (l) Titration is the experimental process of

figuring out the pH of a mystery solution by neutralizing it by incrementally adding small amounts of a known solution. The titration is complete when we reach the equivalence point (where [H3O+]=[OH] ). n=CxV moles = concentration x volume Titration H3O+ (aq) + OH(aq) H2O (l) + H2O (l)

Example: How many moles of H3O+ would it take to neutralize 5 liters of a 0.1 M NaOH solution? n=CxV mol n = (0.1 L )(5 L) n = 0.5 mol H3O+ Buffers

Buffers stabilize the pH of acids. The salt of the conjugate base to the acidic solution can act as a buffer. CH3COOH(aq)+H2O(l) acid CH3COO(aq)+ H3O+(aq) conjugate base How would the equilibrium shift if we added

NaCH3COO to the acidic solution? adding more CH3COO H He Li Be B C

N O F Na Mg Al Si

P S Cl Ar K Ca Ne

Br Kr I Xe H He Li Be B

C N O F Na Mg Al

Si P S Cl Ar K Ca

Ne Br Kr I Xe 4 e in valence shell Measuring the Strength of Acids/Bases

We often speak of acid concentrations in molarity (ex. 2.0 M HCl solution, aka ) 2.0 mol HCl 1 L solution But which is stronger 2.0 M HCl solution or 2.0 M CH3COOH solution? Since we use both strong and weak acids, a

more consistent measurement would tell us just the concentration of hydronium ions. [H+] really means [H3O+] H3O+ (aq) + OH(aq) H2O (l) + H2O (l) H3O+ (aq) + OH(aq) H2O (l) + H2O (l)

Conjugate Acid/Base Pairs CH3COOH(aq) +H2O(l) acid CH3COO(aq)+ H3O+(aq) conjugate base NH3 (aq) + H2O(l)

NH4+(aq)+ OH(aq) base conjugate acid Conjugate Acid/Base Pairs CH3COOH(aq) +H2O(l) H+ acid

CH3COO(aq)+ H3O+(aq) conjugate H+ base Conjugate Acid/Base Pairs NH3 (aq) + H2O(l) + H base

NH4+(aq)+ OH(aq) + H conjugate acid

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