Chapter 14 Acids and Bases

Homework Assigned Questions and Problems (odd only) Section 14.1 (optional) Section 14.2 Read this section Section 14.3 Read this section Section 14.4 33, 35, 37, 39 Section 14.5 Section on Neutralization Reactions, 43 Section 14.6 51, 53, 55, 97, 103 Section 14.8 65, 67, 69, 109, 119 Section 14.9 71, 73, 75, 77, 79, 81, 83, 85, 87, 107, 121

14.2 Acids: Properties and Examples Among the most common and important compounds known Aqueous solutions are important in biological systems chemical industrial processes Acetic Acid HC2H3O2(aq) Early known characteristics Causes a sour taste (lemons and vinegar) Causes litmus (dye) to change from blue to red Dissolves some metals generating hydrogen gas (e.g., zinc and iron) 2Al(s)+ 6HCl(aq)→ 2AlCl3(aq)+3H2(g) Early known characteristic based on what they did rather than their chemical composition

14.2 Acids: Properties and Examples Common Acids (common laboratory reagents) Sulfuric acid (H2SO4): most produced chemical in U.S. Used to produce phosphoric acid and phosphate fertilizers, rust removal in iron and steel production, paper production Hydrochloric acid (HCl): steel production, organic compound production, pH control and neutralization, main component of stomach acid Nitric Acid (HNO3): mainly used in the production of fertilzers, explosives, and used to dissolve metals See table 14.1

Section 5.9 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

Section 5.9 Naming Oxyacids Produce a hydrogen ion (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

14.3 Bases: Properties and Examples Also recognized as an important group of compounds Aqueous solutions are important in biological systems chemical industrial processes Early known characteristics Causes a bitter taste (when dissolved in water) Causes litmus (dye) to change from red to blue Dissolves fats that are placed in base solutions Feel slippery like soap

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)

14.4 The Arrhenius Definition Arrhenius (1884) first person to recognize the essential nature of acids and bases Defined them in terms of chemical composition Arrhenius defined an acid as a substance that produces hydrogen ions (H+) in water Acids are covalent compounds that ionize in water to produce H+ Acids: hydrogen chloride gas

14.4 Molecular Definitions of Acids and Bases Arrhenius 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 oxyacids (will also gen. H+)

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

Types of Acids Binary Acids Oxyacids Molecular compounds in which hydrogen is attached to a second nonmetallic element Formed from the pure compound which has been dissolved in water Oxyacids Molecular compounds composed of hydrogen, oxygen, and a nonmetal (e.g. S, C, N, Cl, or P) Molecular compounds which become acids when dissolved in water Acid hydrogen is attached to an oxygen Formulas are like polyatomic ions with hydrogen(s) attached

Types of Acids H C O H C O O- H+ Acetate Anion Acetic Acid Organic Acids Organic compounds (carbon-containing compounds) with acidic properties. Most common type are called carboxylic acids. The acidic portion of the molecule is called a carboxyl group For example, acetic acid (the active component of vinegar) contains a carboxyl group H 3 C O H 3 C O O- H+ Acetate Anion Acetic Acid

14.3 Bases: Properties and Examples Arrhenius Bases Arrhenius defined a base as a substance which contains hydroxide and produces hydroxide ions (OH-) in water Bases are ionic compounds that produce hydroxide ions (OH-) in water Bases:

14.3 Bases: Properties and Examples Arrhenius Bases When dissolved in water will generate OH- and a metal ion. Two common examples: Bases are 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

14.4 The Brønsted-Lowry Definition of Acids and Bases Arrhenius definition is widely used, but model has limitations Only for aqueous solutions Does not explain how compounds such as ammonia (does not contain OH-) produce a basic water solution New model (1923) by Brønsted and Lowry proposed a broader definition New model applied to both aqueous and non-aqueous solutions Explained how compounds without OH- could produce a basic solution when added to water

14.4 The Brønsted-Lowry Definition of Acids and Bases Brønsted-Lowry Acid Any substance that can donate a proton (H+) to another substance: a proton donor Brønsted-Lowry Base Any substance that can accept a proton (H+) from another substance: a proton acceptor Proton donation does not occur unless a proton acceptor is present Arrhenius Acids/Bases vs. Br.-Lowry Acids/Bases All acids in the Arrhenius definition are also acids according to the Bronsted/Lowry model The Br.-Lowry definition of bases is quite different from that of Arrhenius H+ ions produced by Br.-Lowry acids also react with water Some B/L bases are not considered Arrhenius acids

Water as a Base For example: HCl is a Bronsted/Lowry acid because it produces H+ in water H+ does not exist in water due to the strong attraction to the polar water molecule In an aqueous solution the H+ ion generated reacts with water to form H3O+ H+ acceptor Hydronium ion It is covalently bonded to the water molecule H+ donor

14.4 Identifying Brønsted-Lowry Bases and Their Conjugates Acid Base Conjugate Acid Conjugate Base Any chemical reaction that involves a Bronsted-Lowry acid must also involve a Br.-Lowry 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)

14.4 Identifying Brønsted-Lowry Bases and Their Conjugates Acid Base Conjugate Acid Conjugate Base HCl is the Brønsted-Lowry acid because it is donating a proton (H+) to the water molecule The water molecule is the Brønsted-Lowry base since it accepts a proton (H+)

14.4 Conjugate Acid/Base Pairs Acids/bases and conj acids/conj bases are 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

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

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

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

14.5 Reactions of Acids and Bases Neutralization Reactions When Arrhenius acids and bases are mixed, they will react with each other The acidic properties and the basic properties are destroyed and this means both substances are neutralized A neutralization is the reaction between an acid and a hydroxide base which forms a salt and water

14.5 Reactions of Acids and Bases Neutralization Reactions A neutralization is a double-displacement reaction Whenever an acid is completely reacted with a base, a neutralization occurs The net ionic equation is water formation Acid Base Salt Water

14.5 Reactions of Acids and Bases Neutralization Reactions 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 both reactants are neutralized The salt contains the positive ion from the base and the negative ion from the acid

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

14.8 Water: Acid and Base in One (Ionization of Water) Substances that can act as acids or bases are termed amphoteric Water can act as an acid or a base Amphoteric substance H+ acceptor Hydronium ion Base Acid Conj. Acid Conj. Base H+ donor

14.8 Water: Acid and Base in One (Ionization of Water) Experiments show that in a sample of pure water, a very small percentage has dissociated to produce ions It involves a proton transfer (a Br/L acid-base rxn) The net result is an equal amount of H3O+ and OH- Equal amounts produced Base Acid Conj Acid Conj Base

14.8 Water: Acid and Base in One (Ionization of Water) The acid-base reaction of water with itself is called autoionization (self-ionization) It is an equilibrium reaction Ionization of water to form hydronium ion and hydroxide ion is balanced by the recombination of ions to form water At 25 ºC, the concentration of each ion: 1.0×10-7 M Square brackets around a compound denotes “concentration” in moles per liter The equilibrium concentration is 1.0 X 10-7

14.8 Ion-Product Constant for Water At any given temperature, the product of the concentration of H+ and OH- is always a constant This value can be calculated at 25 °C since the concentration (each ion) is known Kw (Ion-Product Constant for Water) is the product of the H3O+ and OH- ion molar concentrations in water Valid in pure water or water with solutes

14.8 Ion-Product Constant for Water If the [H3O+] is increased by the addition of acidic solute, the [OH-] must decrease until the expression is 1.0 × 10-14 is satisfied Or, if [OH-] is increased by the addition of a basic solute to the water, the [H3O]+ must similarly decrease

14.8 Ion-Product Constant for Water In an aqueous solution, neither [H3O+] or [OH-] is ever zero An acid is a substances that will increase the [H3O+] in solution All acidic solutions have a higher [H3O+] than [OH-] An base is a substance that will increase the [OH-] ions in solution All basic solutions have a higher [OH-] than [H3O+]

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

Using Kw in Calculations, Example 1 Calculate [H3O+] in a solution in which [OH-] = 2.0x10-2 M. Is this solution acidic, basic or neutral?

14.9 The pH and pOH Scale H3O+ (from H+) concentrations range from very high values to extremely small valves Difficult and inconvenient to work with numbers over such a large range For example, [H3O+] of 10 M is 1000 trillion times greater than 10-14 M The pH scale of a solution was proposed as an easier and more practical way to handle such large numbers

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

14.9 The pH Scale The pH scale uses a common logarithm based on powers of 10 Expressed mathematically: For a number expressed in scientific notation with a coefficient of 1, the logarithm of that number is equal to the value of the exponent Take the log of the number and then multiply by negative one (change the sign)

14.9 Calculating the pH from [H3O+] The negative log of the H3O+ 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) Given: 2 SF 2 D.P.

14.9 Calculating the pH from [H3O+] Calculate the pH for the following solutions A solution in which [H3O+]=1.0x10-3 M A solution in which [OH-]= 5.0x10-5 M

Calculating the pH from [H3O+] Calculating the pH from [OH-]

Calculating the [H3O+] from pH It is often necessary to calculate the hydronium ion concentration for a solution from its pH value The logarithm must be undone To do this requires determining the antilog of the pH value The antilog can be obtained on a calculator using the antilog function Use (-) pH value inv log

Calculating the [H3O+] from pH The pH of rainwater in a polluted area was found to be 3.50. What is the [H3O+] for this rainwater?

The pH Scale pH is a log scale A change in one unit on the pH scale means a tenfold increase or decrease in [H3O+] Every time an exponent changes by one, the pH changes by one Lowering the pH increases the [H3O+] low pH value= acidic solution high pH value= basic solution

Measuring pH Use a pH meter 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 Indicators A compound that exhibits different colors depending on the pH of its surroundings

The pOH Scale The negative log is also a way of expressing the magnitudes of other quantities The concentration of OH- can be expressed as pOH Same as pH scale, but associated with the [OH-] Low pOH value means high [OH-] High pOH value means low [OH-]

pH and pOH In an aqueous solution, the sum of the pH and pOH is always 14 The value 14 corresponds to the of Kw equation -log 1.0 × 10-14

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

Calculating pOH from pH 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?

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

14.6 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 Only dissociated molecules influence the pH value Titration will give info about the total number of acid or base molecules present (concentration)

14.6 Acid-Base Titration Titration involves the gradual addition of one solution to another until the solute in the first solution has reacted completely with the solute in the second solution First measure a known volume of acid (unknown molarity) to the flask To determine the concentration of an acid solution, add a solution of base of known concentration to the flask by a buret

14.6 Acid-Base Titration At endpoint, molesacid = molesbase Mbase Base addition continues until all the acid has completely reacted with the added base: endpoint or equivalence point 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 At endpoint, molesacid = molesbase Mbase The selected indicator is a compound, weak acid or weak base that has a CA or CB of a different color. It is selected if it will change color at a pH corresponding as close as possible to the pH of the solution when the titration is complete. Vbase endpoint Vacid

Macid Vacid = Mbase Vbase 14.6 Acid-Base Titration As the base (NaOH) is added, OH- will react with and neutralize HCl (and free H+), forming water At endpoint all of the acid has completely reacted with all of the base At endpoint, molesacid = molesbase At endpoint, Macid Vacid = Mbase Vbase

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