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Ch. 14: Acids and Bases

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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  1. Ch. 14: Acids and Bases Dr. Namphol Sinkaset Chem 152: Introduction to General Chemistry

  2. I. Chapter Outline • Introduction • Acid/Base Definitions • Titrations • Strong vs. Weak • The pH Scale • Buffers

  3. I. Introduction

  4. 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.

  5. I. Two Acids

  6. 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.

  7. 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.

  8. 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.

  9. 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)

  10. II. An Arrhenius Base • Sodium hydroxide is an example of an Arrhenius base.

  11. 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.

  12. 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!

  13. 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.

  14. 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.

  15. 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-.

  16. 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.

  17. II. Conjugate Acid-Base Pairs • The formulas of conjugate pairs differ by only one H+!

  18. 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)

  19. 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 [ ].

  20. 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.

  21. 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.

  22. 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?

  23. 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.

  24. IV. Strong Acids and Bases • Strong acids and bases are strong electrolytes; they dissociate completely.

  25. IV. Strong Acids and Bases

  26. IV. Weak Acids and Bases • Weak acids and bases are weak electrolytes; they do not dissociate completely.

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

  28. V. Water Reacts w/ Itself! • We said before that water is amphoteric; it can also react with itself in an acid/base reaction.

  29. 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.

  30. 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.

  31. 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?

  32. 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.

  33. 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…

  34. V. Sig Figs for log

  35. 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.

  36. 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.

  37. VI. Example of a Buffer

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