Ch 14 acids and bases
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Ch. 14: Acids and Bases. Dr. Namphol Sinkaset Chem 152: Introduction to General Chemistry. I. Chapter Outline. Introduction Acid/Base Definitions Titrations Strong vs. Weak The pH Scale Buffers. I. Introduction. I. Acids. Acids have characteristic properties. Taste sour.

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Ch. 14: Acids and Bases

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Ch. 14: Acids and Bases

Dr. Namphol Sinkaset

Chem 152: Introduction to General Chemistry

I. Chapter Outline

  • Introduction

  • Acid/Base Definitions

  • Titrations

  • Strong vs. Weak

  • The pH Scale

  • Buffers

I. Introduction

I. Acids

  • Acids have characteristic properties.

    • Taste sour.

    • Dissolve many metals.

    • Turn blue litmus paper red.

  • Common acids include: hydrochloric, sulfuric, nitric, acetic, carbonic, and hydrofluoric.

I. Two Acids

I. Bases

  • Bases have characteristic properties.

    • Taste bitter.

    • Feel slippery.

    • Turn red litmus paper blue.

  • Common bases include: sodium hydroxide, potassium hydroxide, sodium bicarbonate, and ammonia.

II. Acid/Base Definitions

  • There are several different definitions for acids and bases.

  • What definition you use depends on what kinds of compounds you are studying and what’s convenient.

  • We will cover the two most commonly used definitions.

II. The Arrhenius Definitions

  • An acid is a substance that produces H+ ions in aqueous solution.

  • A base is a substance that produces OH- ions in aqueous solution.

  • Note that these definitions are restricted to water-based solutions.

II. An Arrhenius Acid

  • HCl is an example of an Arrhenius acid.

  • Note that H+ always attaches to a water molecule to form H3O+, the hydronium ion.

  • H+(aq) = H3O+(aq)

II. An Arrhenius Base

  • Sodium hydroxide is an example of an Arrhenius base.

II. Brønsted-Lowry Definitions

  • An acid is a proton (H+ ion) donor.

  • A base is a proton (H+ ion) acceptor.

  • Notice that the focus in these definitions is on transfer of H+.

  • Notice that there is no dependence on aqueous solutions, so this definition is more widely applicable.

II. Acid/Base Pairs

  • To use the Brønsted-Lowry definitions, you have to analyze an entire reaction and see what’s giving up H+ and what’s accepting the H+.

  • Under this definition, acids and bases always occur together!

II. A Brønsted-Lowry Acid

HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)

  • We see that the HCl gives up an H+; HCl is the acid.

  • We see that H2O accepts an H+; H2O is the base.

II. A Brønsted-Lowry Base

NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

  • We see that NH3 accepts an H+; NH3 is the base.

  • We see that H2O gives up an H+; H2O is the acid.

II. Water is Amphoteric

  • Notice in the last two slides that H2O was acting as a base in one and as an acid in the other.

  • H2O is amphoteric, a substance that can act as either an acid or a base.

  • Another example would be bisulfate, HSO4-.

II. Conjugate Acid-Base Pairs

  • Under Brønsted-Lowry:

    • The acid loses H+ to become a conjugate base.

    • The base gains H+ to become a conjugate acid.

II. Conjugate Acid-Base Pairs

  • The formulas of conjugate pairs differ by only one H+!

II. Practice Problem

  • Identify the conjugate acid-base pairs in the reactions below.

    • HNO3(aq) + H2O(l) H3O+(aq) + NO3-(aq)

    • C5H5N(aq) + H2O(l) C5H5NH+(aq) + OH-(aq)

III. Acid/Base Titration

  • When an acid reacts with a base, the product is always water and a salt.

  • We can use the stoichiometry of the reaction to figure out the concentration of one if we know the concentration of the other.

  • Titration is a technique in which a solution of known [ ] is reacted with another solution of unknown [ ].

III. A Typical Titration

  • Solution of known [ ] is added through a buret.

  • An indicator tells you when to stop.

  • At the equivalence point, moles acid = moles base.

III. Calculating the Unknown [ ]

  • Use the volume added from the buret and the concentration to find moles of known.

  • Use the balanced equation to convert moles of known to moles of unknown.

  • Divide moles of unknown by the sample volume.

III. Sample Problem

  • A 25.0-mL sample of sulfuric acid is titrated with a 0.225 M solution of sodium hydroxide. If it takes 21.27 mL to reach the endpoint, what is the molarity of the sulfuric acid solution?

IV. Acid/Base Strength

  • Different acids and bases have different strengths.

  • There are actually more weak acids and bases than strong acids and bases.

  • Acid/base strength is related to whether they are strong or weak electrolytes.

IV. Strong Acids and Bases

  • Strong acids and bases are strong electrolytes; they dissociate completely.

IV. Strong Acids and Bases

IV. Weak Acids and Bases

  • Weak acids and bases are weak electrolytes; they do not dissociate completely.

IV. Indicating Weakness

  • Equations showing weak acids or bases use a double arrow to indicate incomplete dissociation.

V. Water Reacts w/ Itself!

  • We said before that water is amphoteric; it can also react with itself in an acid/base reaction.

V. Water Ion Product Constant

  • In pure water at 25 °C, there’s always a little H3O+ and OH- in equal amount.

  • Specifically, [H3O+] = [OH-] = 1.0 x 10-7 M.

  • When these concentrations are multiplied, you get the ion product constant for water, Kw.

  • Kw = [H3O+][OH-]

  • At 25 °C, Kw = 1.0 x 10-14.

V. Acidic/Basic Solutions

  • In an acidic solution, additional H3O+ ions exist, increasing [H3O+].

  • In a basic solution, additional OH- ions exist, increasing [OH-].

  • However, in all aqueous solutions, the product of hydronium and hydroxide concentrations always equals Kw.

V. Sample Problem

  • Calculate the [H3O+] concentration of a solution that has [OH-] = 1.5 x 10-2 M at 25 °C. Is the solution acidic or basic?

V. The pH Scale

  • pH is simply another way to specify the acidity or basicity of a solution.

    • pH < 7 is acidic; pH = 7 is neutral; pH > 7 is basic.

V. pH is a log Scale

  • pH = -log [H3O+]

  • Since it’s a log scale, a one unit change is actually a 10x change.

  • log is a different type of math, so it has its own sig fig rule…

V. Sig Figs for log

V. Sample Problems

  • Perform the following calculations.

    • Calculate the pH of a solution in which the hydronium concentration is 4.2 x 10-3 M.

    • Calculate the pH of a solution in which the hydroxide concentration is 7.89 x 10-8 M.

    • If the pH of a solution is 4.67, calculate the concentration of hydronium.

VI. Resisting Changes in pH

  • The only things that affect the pH are free H3O+ and OH-.

  • If we can create a solution that “captures” any added H3O+ or OH-, then we can resist changes in pH.

  • A solution that can do this is called a buffer.

VI. Example of a Buffer

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