Chapter 14 Acids and Bases

Chapter 14 Section 1 – Properties of Acids and Bases Section 2 – Acid Base Theories Section 3 – Acid Base Reactions

14.1 Properties of Acids and Bases List five general properties of aqueous acids and bases. Name common binary acids and oxyacids, given their chemical formulas. List five acids commonly used in industry and the laboratory, and give two properties of each. Define acid and base according to Arrhenius’s theory of ionization. Explain the differences between strong and weak acids and bases.

Properties of: Acids Bases Sour taste Conducts electricity Turns litmus paper red Reacts with bases to produce salts and water Reacts with some metals and releases hydrogen gas Bitter taste Feels slippery Conducts electric current Turns litmus paper blue Reacts with acids to produce salts and water

Binary Acids Contains only two different elements Nomenclature: Hydrogen & an electronegative, nonmetal Nomenclature: hydro - _________ - ic acid

Diatomic Nomenclature

Oxyacid Contains hydrogen, oxygen, and a third element (hydrogen with a polyatomic ion) Nomenclature:

Acid Names

Oxyacids

Common Industrial Acids Sulfuric Acid Sulfuric acid is the most commonly produced industrial chemical in the world. Nitric Acid Phosphoric Acid Hydrochloric Acid Conc. HCl is commonly referred to as muriatic acid. Acetic Acid Pure acetic acid is a clear, colorless, and pungent-smelling liquid known as glacial acetic acid.

Arrhenius Acids and Bases Increases concentration of H+ ions in solution Arrhenius Bases: Increases concentration of OH- ions in solution

Arrhenius Acid Base Video

Acid/Base Strength Strong acid: Weak acid: Ionizes completely in solution and is an electrolyte Higher the KA, the greater the strength as an acid K reveals a greater extent of ionization Example: HCl, HClO4, HNO3 Weak acid: Releases few hydrogen ions in solution Hydronium ions, anions and dissolved acid molecules present Examples: HCN, Organic acids – HC2H3O2

Dissociation Constants Strong vs. Weak Base Strong bases ionizes completely in solution and is a strong electrolyte KB = dissociation constant of a base Higher the KB , the greater the strength of a base

Aqueous Acids

Base Strength Strong bases: Weak bases: Ionic compounds containing metal cation and hydroxide ion (OH-) Dissociates in water Weak bases: Molecular compounds do not follow Arrhenius definition: Ammonia (NH3) Produces hydroxide ions when it reacts with water molecules

Base Strength

Acidic solution has greater [H3O+] Basic solution has greater [OH–]

14.2 Acid Base Theories Define and recognize Brønsted-Lowry acids and bases. Define a Lewis acid and a Lewis base. Name compounds that are acids under the Lewis definition but are not acids under the Brønsted-Lowry definition.

Bronsted-Lowry Acid Bronsted-Lowry Acid: Proton (H+) donor Hydrogen chloride acts as a Bronsted-Lowry acid when it reacts with ammonia. Water can also act as a Bronsted-Lowry acid

Bronsted-Lowry Base Bronsted-Lowry Base: Proton acceptor Ammonia accepts a proton from hydrochloric acid.

Bronsted-Lowry Acid Base Reactions Protons are transferred from one reactant (the acid) to another (the base) acid base

Conjugate Acid – Base acid conjugate base Conjugate Base: The species that remains after a Bronsted-Lowry acid has given up a proton Conjugate Acid: The species that remains after a Bronsted-Lowry base has accepted a proton acid conjugate base

Conjugate Acid Base Pairs Match up the acid-base pairs (proton donor-acceptor pairs) acid1 base2 conjugate base1 conjugate acid2

Strength of Acid Base Pairs The stronger the acid, the weaker the conjugate base The stronger the base, the weaker the conjugate acid strong acid base acid weak base

Proton transfer favors the production of the weaker acid and base. stronger acid stronger base weaker acid weaker base weaker acid weaker base stronger acid stronger base

Acid Base Strength

Amphoteric Any species that can react as either an acid or a base Example: water acid1 base2 acid2 base1 base1 acid2 acid1 base2

Amphoteric Water Video

Other Amphoteric Compounds Covalently bonded –OH group in an acid is referred to as a hydroxyl group Molecular compounds with hydroxyl groups can be acidic or amphoteric The behavior of the compound is affected by the number of oxygen atoms bonded to the atom connected to the –OH group *The more oxygen’s in a polyatomic formula, the greater the strength of polyatomic as an acid

Oxyacids of Chlorine

Brønsted-Lowry Acid Base Video

Monoprotic Acids Can donate only one proton (hydrogen ion) per molecule One ionization step

Monoprotic and Diprotic Acids

Polyprotic Acids Donates more than one proton per molecules Multiple ionization steps Diprotic – donates 2 protons Ex: Triprotic – donates 3 protons Ex: Sulfuric acid solutions contain H3O+, HSO4-, SO4- ions 1. 2.

Lewis Acid Lewis acid: Lewis base: Atom, ion, or molecule that ACCEPTS an ELECTRON PAIR to form a covalent bond A proton (hydrogen ion) is a Lewis acid Lewis base: Atom, ion, or molecule that DONATES an ELECTRON PAIR to form a covalent bond

Lewis Acid A lewis acid might not include hydrogen Silver as a lewis acid:

Lewis Acid Base Video

Acid and Base Definitions

Acid Base Definitions Video

14.3 Acid Base Reactions Describe a conjugate acid, a conjugate base, and an amphoteric compound. Explain the process of neutralization. Define acid rain, give examples of compounds that can cause acid rain, and describe effects of acid rain.

Neutralization Reactions What does it mean to neutralize something? Neutralization reactions: Hydronium and hydroxide ions react to form water The left over cation and anion in solution produce a salt (ionic compound)

Neutralization Reactions

Neutralization Reaction Video

Acid Rain NO, NO2, CO2, SO2, and SO3 gases from industrial processes can dissolve in atmospheric water to produce acidic solutions. Very acidic rain is known as acid rain. Acid rain can erode statues and affect ecosystems.

Chapter 15 Acid Base Titration and pH

Chapter 15 Section 1 – Aqueous Solutions and the Concept of pH Section 2 – Determining pH and Titrations

15.1 Aqueous Solutions and pH Describe the self-ionization of water. Define pH, and give the pH of a neutral solution at 25°C. Explain and use the pH scale. Given [H3O+] or [OH−], find pH. Given pH, find [H3O+] or [OH−].

Self Ionization of Water Two water molecules produce a hydronium ion and hydroxide ion by proton transfer In water at 25°C, [H3O+] = 1.0 ×10−7 M and [OH−] = 1.0 × 10−7 M The ionization constant of water, Kw Kw = [H3O+][OH−]

Kw = [H3O+][OH−] = (1.0 × 10−7)(1.0 × 10−7) = 1.0 × 10−14 At 25OC Kw = [H3O+][OH−] = (1.0 × 10−7)(1.0 × 10−7) = 1.0 × 10−14 Kw = 1.0 x 10-14 Kw increases as temperature increases

Ion Concentration neutral acidic basic [H3O+] = [OH−] [H3O+] > 1.0 × 10−7 M [OH−] > [H3O+] basic [OH−] > 1.0 × 10−7 M

Calculating Concentration Strong acids and bases are considered completely ionized or dissociated in aqueous solutions. 1 mol 1 mol 1 mol 1.0 × 10−2 M NaOH therefore, [OH−] = 1.0 × 10−2 M [H3O+] is calculated using Kw

Example Problem 1 [H3O+] = ______________ Unknown: [OH-] = ? Given: [HCl] = 2.0 × 10−4 M [H3O+] = ______________ Unknown: [OH-] = ? Kw = [H3O+][OH−] = 1.0 × 10−14

pH = −log [H3O+] = −log(1 × 10−7) = −(−7.0) = 7.0 Definition of the pH of a solution: negative of the common logarithm of the hydronium ion concentration, [H3O+]. pH = −log [H3O+] Example: a neutral solution has a [H3O+] = 1×10−7 pH = −log [H3O+] = −log(1 × 10−7) = −(−7.0) = 7.0

pH Values as Specified [H3O+]

The pH Scale

pOH pOH = −log [OH–] pH + pOH = 14.0 The pOH of a solution is defined as the negative of the common logarithm of the hydroxide ion concentration, [OH−]. pOH = −log [OH–] pH + pOH = 14.0 Example: a neutral solution has a [OH–] = 1×10−7 the pH of this solution is?

Calculating [H3O+] from pH Finding the [H3O+] from pH requires taking the antilog of the negative pH [H3O+] = antilog (-pH) You can find the [OH−] by also taking the antilog of the negative pOH. [OH-] = antilog (-pOH)

The Circle of pH pH pOH [ H3O+] [ OH-] -log [H3O+] antilog (-pH) antilog (-pOH) -log [OH-] = 1.0x10-14 + pOH = 14

pOH Video

pH Values of Some Common Materials

Approximate pH Range of Common Materials

Comparing pH and pOH Video

pH of Weak Acids and Bases The pH of solutions of weak acids and weak bases must be measured experimentally. The [H3O+] and [OH−] can then be calculated from the measured pH values.

Significant Figures There must be as many significant figures to the right of the decimal as there are in the number whose logarithm was found. Example: [H3O+] = 1 × 10−7 one significant figure pH = 7.0

15.2 Determining pH and Titrations Describe how an acid-base indicator functions. Explain how to carry out an acid-base titration. Calculate the molarity of a solution from titration data.

Indicators Acid-base indicators: compounds whose colors are sensitive to pH. The pH range over which an indicator changes color is called its transition interval.

pH Meters pH meter determines the pH of a solution by measuring the voltage between the two electrodes that are placed in the solution. The voltage changes as the hydronium ion concentration in the solution changes. Measures pH more precisely than indicators

Color Ranges of Indicators

Color Ranges of Indicators

Color Ranges of Indicators

Antacids Video with Methyl Orange

H3O+(aq) + OH−(aq) 2H2O(l) Titration Neutralization occurs when hydronium ions and hydroxide ions are supplied in equal numbers by reactants. H3O+(aq) + OH−(aq) 2H2O(l) Titration: the controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concentration.

Titration Points equivalence point: point at which the two solutions used in a titration are present in chemically equivalent amounts end point: point in a titration at which an indicator changes color

Which indicator do I choose? pH less than 7 Indicators that change color at pH lower than 7 are used to determine the equivalence point of strong-acid/weak-base titrations. strong-acid/weak-base titration = acidic. pH at 7 Indicators that undergo transition at about pH 7 are used to determine the equivalence point of strong-acid/strong base titrations. strong acids/strong bases = salt solution with a pH of 7.

Which indicator do I choose? pH greater than 7 Indicators that change color at pH higher than 7 are used to determine the equivalence point of weak-acid/strong-base titrations. weak-acid/strong-base = basic

Titration Curve Strong Acid and a Strong Base Equivalence Point: pH at 7

Titration Curve Weak Acid and a Strong Base Equivalence Point: pH higher than 7

Titration Curve Strong Acid and a Weak Base Equivalence Point: pH less than 7

Titration Problems: * Can be used to determine concentration of unknown solution or volume of added standard Start with the balanced equation for the neutralization reaction Make amount of acid and base chemically equivalent to each other (1 to 1 mol ratio). Determine the molarity of the unknown solution. Equation: M1V1 = M2V2 1: starting solution 2: added standard

Molarity and Titration standard solution: solution that contains the precisely known concentration of a solute primary standard: highly purified solid compound used to check the concentration of the known solution The standard solution can be used to determine the molarity of another solution by titration.

Performing a Titration – Set up

Performing a Titration – Set up Acid

Performing a Titration – Starting Amount

Performing a Titration – Set up Base

Performing a Titration - Titrating

Performing a Titration – End Point

Molarity and Titration Determine the molarity of an acidic solution, 10 mL HCl, by titration Titrate acid with a standard base solution 20.00 mL of 5.0 × 10−3 M NaOH was titrated Write the balanced neutralization reaction equation. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) 1 mol 1 mol 1 mol 1 mol

Molarity and Titration Calculate the number of moles of NaOH used in the titration. 20.0 mL of 5.0 × 10−3 M NaOH is needed to reach the end point mol of HCl = mol NaOH = 1.0 × 10−4 mol Calculate the molarity of the HCl solution

Example Problem In a titration, 27.4 mL of 0.0154 M Ba(OH)2 is added to a 20.0 mL sample of HCl solution of unknown concentration until the equivalence point is reached. What is the molarity of the acid solution?

Example Problem Solution Given: 27.4 mL of 0.0154 M Ba(OH)2 Unknown: ? M HCl of 20.0 mL Solution: Write balanced equation: Ba(OH)2 + 2HCl BaCl2 + 2H2O 1 mol 2 mol 1 mol 2 mol

1. Calculate Moles of Given

2. Write a mole ratio: moles of base used to moles of acid produced

3. Calculate Unknown Molarity