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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

Ch. 14: Acids and Bases

Dr. Namphol Sinkaset

Chem 152: Introduction to General Chemistry

i chapter outline
I. Chapter Outline
  • Introduction
  • Acid/Base Definitions
  • Titrations
  • Strong vs. Weak
  • The pH Scale
  • Buffers
i acids
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 bases
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
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
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
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
II. An Arrhenius Base
  • Sodium hydroxide is an example of an Arrhenius base.
ii br nsted lowry definitions
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
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 br nsted lowry acid
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 br nsted lowry 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
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
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 pairs1
II. Conjugate Acid-Base Pairs
  • The formulas of conjugate pairs differ by only one H+!
ii practice problem
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
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
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
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
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
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
IV. Strong Acids and Bases
  • Strong acids and bases are strong electrolytes; they dissociate completely.
iv weak acids and bases
IV. Weak Acids and Bases
  • Weak acids and bases are weak electrolytes; they do not dissociate completely.
iv indicating weakness
IV. Indicating Weakness
  • Equations showing weak acids or bases use a double arrow to indicate incomplete dissociation.
v water reacts w itself
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
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
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
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
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
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 sample problems
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
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.
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