Chapter 14 Acids and Bases

Homework Assigned Questions and Problems (odd only) Section 14.1 (14.1 to 14.5) Section 14.2 (14.7 to 14.15) Section 14.3 (14.17 to 14.25) Section 14.4 Section 14.5 (14.31 to 14.39) Section 14.6 (14.41 to 14.49) Section 14.7 (14.51 to 14.55) Section 14.8 (14.57 to 14.63) Additional Questions and Problems and Challenge Questions (except 14.99, 14.105)

Acids Among the most common and important compounds Aqueous solutions are important in biological systems and in chemical industrial processes Characteristics Cause sour taste of lemons and vinegar Digest food in stomach Dissolves some metals generating hydrogen gas Chemical produced in the largest quantity in the USA Sulfuric acid H2SO4

Acids Acids When dissolved in water will generate H+ and an anion The H+ that is generated will give a sour taste Vinegar (acetic acid) Lemon (citric acid) Most acids are oxo acids (will also gen. H+)

Naming Binary Acids Use the prefix hydro- before the root name of the element Add the suffix -ic and the word acid to the root name for the element Example: HCl hydrochloric acid Example: HI hydroiodic acid Hydrogen chloride as a gas Hydrogen iodide as a gas

Naming Oxo Acids Produce H+ and a polyatomic ion when dissolved in water Composed of hydrogen, oxygen, and another nonmetal Use the root name of the polyatomic ion If it ends in -ate use the suffix -ic acid If it ends in -ite use the suffix -ous acid Example: H2SO4 (from SO42- , sulfate ion) sulfuric acid Example: H2SO3 (from SO32- , sulfite ion) sulfurous acid It looks like a hydrogen compound of a polyatomic ion

Bases Bases When dissolved in water will generate OH- and a metal ion. Sometimes called alkalis The OH- that is generated will give a bitter taste Slippery feel (like soap) Most bases that are generated are composed from group 1A and 2A metals Others such as aluminum and iron III hydroxide are insoluble

Naming Bases Most bases are usually ionic compounds The hydroxide ion has a charge of (-1) and is combined with a positively charged ion (group IA or IIA metal ion) Hydroxides (bases) are named by their cation first, then the word “hydroxide” is added to the metal cation Most common bases: NaOH (sodium hydroxide) KOH (potassium hydroxide) Ca(OH)2 (calcium hydroxide) NH4OH (ammonium hydroxide)

Arrhenius First person to recognize the essential nature of acids and bases Acids: Produce hydrogen ions (H+) in water Bases Produce hydroxide ions (OH-) in water

Brønsted-Lowry Acids and Bases Limitations to Arrhenius model Only for aqueous solutions Free H+ does not exist in water New model by Brønsted and Lowry Acid Any substance that can donate a proton (H+) to another substance: Proton donor Base Any substance that can accept a proton from another substance: Proton acceptor

Water as a Base H+ does not exist in water due to the strong attraction to the polar water molecule H3O+ is called a hydronium ion Real way that protons exist in water It is covalently bonded to the water molecule

Brønsted-Lowry Model The reaction involves a proton transfer Acid Base Conjugate Acid Conjugate Base The reaction involves a proton transfer Acid can lose its proton to form a conjugate base (something that could accept a proton back again) Base accepts a proton to form a conjugate acid (something that could donate a proton)

Brønsted-Lowry Model Acid Base Conjugate Acid Conjugate Base The HCl is the Brønsted-Lowry acid because it is donating proton (H+) to water molecule The water molecule is the Brønsted-Lowry base since it accepts a proton

Conjugate Acid/Base Pairs Related to each other by donating/accepting a single proton (H+) The acid of the pair has the proton The base of the pair does not Every acid has a conjugate base Every base has a conjugate acid “Conjugate” is given to the part of the pair that is produced in the reaction

Conjugate Acid/Base Pairs Each acid is related to a base on the opposite side If on the reactants side (left side) substances are called “acid” or “base” If on the products side (right side) substances are called “conjugate acid” or “conjugate base” Base Acid Conjugate Acid Conjugate Base Each acid is related to a base on the opposite side

Conjugate Acid/Base Pairs Which of the following represent conjugate acid-base pairs? H2O, H3O+ OH-, HNO3 H2SO4, SO42- HC2H3O2, C2H3O2-

Conjugate Acid/Base Pairs H2O, H3O+ OH-, HNO3 H2SO4, SO42- HC2H3O2, C2H3O2-

Strength of Acids When an acid dissolves in water, it gives its proton to water to form the conjugate base of the acid and a hydronium ion (H3O+) The strength of the acid is determined by the amount of H3O+ that is produced

Strength of Acids Strong acids completely dissociate and form H3O+ If a weak acid, there is nothing stopping the conjugate base and the hydronium ion from reacting to re-form the acid and water Forward Reverse

Strength of Acids There is a competition for the proton between the conjugate base and water Water wins: Acid dissociates completely. It is a strong acid Conjugate base wins: Acid doesn’t dissociate a lot. It is a weak acid

Strong Acids Dissociate 100% (almost) Have very weak conjugate bases (weak reverse reaction) Conjugate base does not react readily with H3O+ Forward reaction predominates

Weak Acids Don’t dissociate very much Have stronger conjugate bases Conjugate base does react readily with H3O+ Reverse reaction predominates

Types of Acids Multiprotic Acids Can donate more than one proton H2SO4 H3PO4 Strong Acid Weak Acid

Types of Acids Oxo Acids Acid hydrogen is attached to an oxygen Acids of Polyatomic Ions

Types of Acids Organic Acids Acids with a carbon backbone (carboxyl group) Acetic Acid (CH3COOH or HC2H3O2)

Ionization of Water Water can act as an acid or a base Amphoteric substance Base Acid Conj Acid Conj Base

Ionization of Water In sample of pure water a small percentage has dissociated to produce ions It involves a proton transfer Results in an equal amount of H+ and OH- Base Acid Conj Acid Conj Base

Ionization of Water Autoionization Ionizing of water to form hydronium and hydroxide ions at 25 ºC Doesn’t occur to a large extent Square brackets around a compound mean “concentration of”

Ion-Product Constant for Water At any temperature, the product of the concentration of H+ and OH- is always a constant Valid in pure water or water with solutes Kw (Ion-Product Constant of Water) Conc. of ion expressed in moles/liter Is 1x10-14

Ion-Product Constant for Water If the [H+] is increased by the addition of acidic solute, the [OH-] must decrease until the expression is 1.0 × 10-14 again Or, if the OH- ions are added to the water, the H+ must likewise decrease

Ion-Product Constant for Water An acid is a substance that will increase the H+ ions in solution All acidic solutions have a higher [H+] than [OH-] An base is a substance that will increase the OH- ions in solution All basic solutions have a higher [OH-] than [H+]

Ion-Product Constant for Water Neutral Solution Acidic Solution Basic Solution In all cases:

Example Calculate [H+] in a solution in which [OH-] = 2.0x10-2 M. Is this solution acidic, basic or neutral?

The pH Scale H+ concentrations range from very high values to extremely small valves Difficult and inconvenient to work with numbers over such a large range i.e. [H+] of 10 M is 1000 trillion times greater than 10-14 The pH scale of a solution was proposed as an easier and more practical way to handle such large numbers

p Scale Used to express very small numbers Based on log 10 If N is a number then pN is: Take the log of the number and then multiply by (-1) to change the sign

The pH Scale The pH scale is defined as the negative log of the molar hydrogen ion concentration Logarithms are exponents The negative powers of 10 in the concentrations are converted to positive numbers

Calculating the pH of Solutions The negative log of the H+ concentration To determine the number of significant Figures for logs: The number of decimal places for the log is equal to the number of sig. figs. in the concentration (original number)

pH Examples Calculate the pH for each of the following solutions A solution in which [H+]=1.0x10-3 M A solution in which [OH-]=5.0x10-5 M

pH Examples

pH Scale pH is a log scale A change in one unit on the pH scale means a tenfold increase or decrease in [H+] Every time exponent changes by one, the pH changes by one lowering the pH increases the [H+] Small pH = acidic solution Large pH = basic solution

Measuring pH Use a pH meter Use pH paper Electronic device that measures the pH of the solution Use pH paper Paper has a chemical in it that changes to different colors depending on the pH of the solution

pOH Scale Same as pH scale, but associated with the [OH-] Low [OH-] means high pOH High [OH-] means low pOH

pH and pOH In an aqueous solution, the sum of the pH and pOH is always 14 The value 14 corresponds to the negative log of Kw

pH and pOH [H+] > [OH-] [H+] = [OH-] [H+] < [OH-]

Example 1 A sample of rain in an area with severe air pollution has a pH of 3.5. What is the pOH of this rain water?

Example 2 The pH of rainwater in a polluted area was found to be 3.50. What is the [H+] for this rainwater?

Example 3 The pOH of a liquid drain cleaner was found to be 10.50. What is the [OH-] for this cleaner?

Calculating pH for Strong Acids Strong acids dissociate 100% Solution contains only H+ and the anion of the acid Find the concentration of H+ Find the pH

pH of Strong Acids Calculate the pH of a solution of 5.0x10-3 M HCl complete dissociation

Reactions of Acids and Bases Neutralization: The reaction between an acid and a base to form a salt and water H+ from the acid combines with the OH- from the base to form water The properties of the both reactants are neutralized The salt contains the positive ion from the base and the negative ion from the acid

Reactions of Acids and Bases The reaction hydrochloric acid and sodium hydroxide: A double replacement reaction Ionic Equation Net Ionic Equation acid base salt water

Reactions of Acids and Bases Balancing Neutralization Reactions Write the reactants and products if HCl reacts with Ba(OH)2 Balance the H+ and OH- in the acid and base

Reactions of Acids and Bases Balance the H2O with the H+ and the OH- Write the salt from the remaining acid and base

Acid-Base Titration Determining the acid or base concentration in a solution is a routing laboratory practice The pH only determines the amount of dissociated H+ in solution Titration will give info about the total number of acid or base molecules present

Acid-Base Titration Add one solution to another until the solute in the first solution has reacted completely with the solute in the second solution To determine the concentration of an acid solution add a solution of base of known concentration to the flask by a buret

Acid-Base Titration molesacid = molesbase Base addition continues until all the acid has completely reacted with the added base: endpoint To determine endpoint, an indicator is used to detect when the acid-base reaction is complete If you know original volume of the acid the volume and concentration of the base the concentration of the acid can be calculated molesacid = molesbase

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