Section 19.1 Acid – Base Theories Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories

Acids Bases vinegar  citrus fruits carbonated drinks  car battery lemon juice  tea Bases calcium hydroxide in mortar  antacids household cleaning agents

Properties of Acids Give foods a tart or sour taste e.g. lemon & vinegar for example Aqueous solutions of acids are electrolytes (conduct electricity) Acids cause certain chemical indicators to change color. Acid + Base Salt + water

Properties of Bases Bases have a bitter taste e.g. soap Bases have a slippery feel Aqueous solutions of bases are electrolytes (conduct electricity) Bases cause certain chemical indicators to change color. Acid + Base Salt + water

Arrhenius Acids & Bases In 1887, Swedish chemist Svante Arrhenius proposed a revolutionary way of defining and thinking about acids and bases Acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution. Bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solution

Monoprotic acids: acids that contain one ionizable hydrogen HNO3 – nitric acid Diprotic acids: acids that contain two ionizable hydrogens H2SO4 – sulfuric acid Triprotic acids: acids that contain three ionizable hydrogens H3PO4 – phosphoric acid

Not all compounds that contain hydrogen are acids e.g. CH4 – methane has weak polar C – H bonds and no ionizable hydrogens. Not an acid. Not all hydrogens in an acid may be released as hydrogen ions. Only hydrogens in very polar bonds are ionizable. In the case where hydrogen is joined to a very electronegative element. e.g. HCl hydrogen chloride very polar covalent molecule

When HCl dissolves in water, it releases hydrogen ions because the hydrogen ions are stabilized by solvation. H2O H – Cl (g) H+ (aq) + Cl- (aq) Hydrogen Hydrogen Chloride chloride ion ion Ionizes to form an aqueous solution of hydronium ions and chloride ions HCl + H2O H3O+ + Cl-

Ethanoic acid CH3COOH is a monoprotic acid due to its structure H O H C C O H H The three H attached to the carbon are in weak polar bonds. They do not ionize. Only the H bonded to the highly electronegative O can be ionized

NaOH (s) Na+ (aq) + OH- (aq) Sodium hydroxide dissociates into sodium ions and hydroxide ions in aqueous solution. H2O NaOH (s) Na+ (aq) + OH- (aq) Sodium Sodium Hydroxide Hydroxide Ion ion Potassium hydroxide dissociates into potassium ions and hydroxide ions in aqueous solution. H2O KOH (s) K+ (aq) + OH- (aq) Potassium Potassium Hydroxide Hydroxide Ion ion

Arrhenius Bases Group one, the alkali metals, react with water to produce solutions that are basic. Group one metals are very soluble in water and can produce concentrated solutions. Group two metals are not very soluble in water. Their solutions are always very dilute.

Bronsted – Lowry Acids and Bases The Bronsted – Lowry theory defines Acid: a hydrogen-ion donor Base: a hydrogen-ion acceptor All acids and bases included in the Arrhenius theory are also acids and bases according to the Bronsted-Lowry theory.

Ammonia is the hydrogen-ion acceptor therefore it is a base. NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq) base acid conjugate conjugate acid base Ammonia is the hydrogen-ion acceptor therefore it is a base. Water is the hydrogen-ion donor and therefore it is an acid. Hydrogen ions are transferred from water to ammonia, which causes the hydroxide-ion concentration to be greater than it is in pure water. When ammonia dissolves and reacts with water, NH4+ is the conjugate acid of the base NH3. OH- is the conjugate base of acid H2O

Conjugate Acids and Bases HCl (g) + H2O (l) H3O+ (aq) + Cl - (aq) acid base conjugate conjugate acid base HCl is the hydrogen-ion donor – thus it is an acid. Water is the hydrogen-ion acceptor – thus it is a base

Conjugate Acid – Base Pair Conjugate acid: the particle formed when a base gains a hydrogen ion Conjugate base: the particle that remains when an acid has donated a hydrogen ion. Conjugate acids and bases are always paired with a base or an acid, respectively. Conjugate acid-base pairs consists of two substances related by the loss or gain of a single hydrogen ion.

Common Conjugate Acid – Base Pairs Conjugate Base HCl Cl- H2SO4 HSO4- H3O+ H2O SO42- CH3COOH CH3COO- H2CO3 HCO3- CO32- NH4+ NH3 OH-

Bronsted – Lowry Acids and Bases A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion (H3O+) Amphoteric – a substance that can act as both an acid and a base e.g. water H2SO4 + H2O H3O+ + HSO4- NH3 + H2O NH4+ + OH-

Lewis Acids and Bases Gilbert Lewis proposed a third Acid Base theory Acid – accepts a pair of electrons during a reaction Base – donates a pair of electrons during a reaction Concept is more general than either the Arrhenius theory or the Bronsted-Lowry theory.

Lewis Acids and Bases Lewis Acid – a substance that can accept a pair of electrons to form a covalent bond. Lewis Base – a substance that can donate a pair of electrons to form a covalent bond. .. H+ : O – H :O: .. H H Lewis Lewis Acid Base

Electron-pair acceptor Acid Base Definitions Type Acid Base Arrhenius H+ producer OH- producer Bronsted - Lowry H+ donor H+ acceptor Lewis Electron-pair acceptor Electron-pair donor

End of Section 19.1

Section 19.2 Hydrogen Ions & Acidity

Hydrogen Ions From Water A water molecule that loses a hydrogen ion becomes a negatively charged hydroxide ion A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion H2O (l) OH- (aq) + H+ (aq) Hydroxide ion Hydrogen ion Self ionization of water – the reaction in which water molecules produce ions

Self Ionization of Water Hydrogen ions in aqueous solution have several names. Some chemists call them protons Some chemists call them hydrogen ions or hydronium ions. For our purposes, either H+ or H3O+ will represent hydrogen ions in aqueous solution. H2O + H2O H3O+ + OH-

In pure water at 25˚C, the equilibrium concentration of hydrogen ions and hydroxide ions are each only 1 x 10-7. In other words the concentration of OH- and H+ are equal in pure water Any aqueous solution in which H+ and OH- are equal is a neutral solution. When [H+] increases [OH-] decreases When [H+] decreases [OH-] increases

For aqueous solutions, the product of the hydrogen ion concentration and the hydroxide ion concentration equals 1.0 x 10-14 [H+] x [OH-] = 1 x 10-14 1 x 10-7 x 1 x 10-7 = 1 x 10-14 This equation is true for all dilute aqueous solutions at 25˚C. Ion-Product Constant for Water (Kw) the product of the concentrations of the hydrogen ions and hydroxide ions in water Kw = [H+] x [OH-] = 1.0 x 10-14

But not all solutions are neutral When some substances dissolve in water, they release hydrogen ions. When hydrogen chloride dissolves in water, it forms hydrochloric acid. H2O HCl (g) H+ (aq) + Cl- (aq)

Ion Product Constant for Water In the previous HCl solution, the hydrogen-ion concentration is greater than the hydroxide-ion concentration. Acidic Solution: A solution in which [H+] is greater than [OH-]. The [H+] of an acidic solution is greater than 1 x 10-7 M

NaOH(s) Na+(aq) + OH-(aq) When sodium hydroxide dissolves in water, it forms hydroxide ions in solution. H20 NaOH(s) Na+(aq) + OH-(aq) In the above solution, the hydrogen-ion concentration is less than the hydroxide-ion concentration. Basic Solution: A solution in which [H+] is less than [OH-] The [H+] of a basic solution is less than 1 x 10-7 Basic solutions are also known as alkaline solutions.

The pH Concept The pH scale was proposed by Danish Scientist Soren Sorensen in 1909. The pH scale is used to express [H+] 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Strongly Neutral Strongly Acidic Basic

Calculating pH pH = - log [H+] The pH of a solution is the negative logarithm of the hydrogen-ion concentration. pH = - log [H+]

In neutral solution, the [H+] = 1 x 10-7M pH = - log [H+] pH = - log (1 x 10-7) pH = 7

Classifying Solutions A solution in which [H+] is greater than 1 x 10-7 has a pH less than 7 is called acidic. A solution in which [H+] is less than 1 x 10-7 has a pH greater than 7 is called basic. The pH of pure water or a neutral aqueous solution is 7 Acidic solution: pH < 7.0 [H+] > 1 x 10-7 M Neutral solution: pH = 7.0 [H+] = 1 x 10-7 M Basic solution: pH > 7.0 [H+] < 1 x 10-7 M

pH can be read from the value of [H+] if it is written in scientific notation and has a coefficient of 1. Then the pH of the solution equals the exponent, with the sign changed from minus to plus e.g. [H+] = 1 x 10-2 has a pH of 2 e.g. [H+] = 1 x 10-13 has a pH of 13

If the pH is an integer, it is also possible to directly write the value of [H+]. pH = 9 ; then [H+] = 1 x 10-9 M pH = 4 ; then [H+] = 1 x 10-4M

A neutral solution has a pOH of 7 The pOH of a solution equals the negative logarithm of the hydroxide-ion concentration pOH = - log [OH-] A neutral solution has a pOH of 7 Acidic solution: pOH > 7 [OH-] < 1 x 10-7 M Neutral solution: pOH = 7 [OH-] = 1 x 10-7 M Basic solution: pOH < 7 [OH-] > 1 x 10-7 M

pH and pOH Relationship pOH + pH = 14 pH = 14 – pOH pOH = 14 – pH

Problem Example Colas are slightly acidic. If the [H+] in a solution is 1.0 X 10 - 5 M , is the solution acidic, basic or neutral. What is the [OH-] of this solution? [H+] = 1.0 X 10 - 5 M which is greater than 1.0 X 10 -7 M so the solution is acidic Kw = [OH-] x [H+] = 1.0 X 10 - 14 [OH-] = 1.0 X 10 - 14 ÷ [H+] [OH-] = 1.0 X 10 - 14 ÷ 1.0 X 10 - 5 [OH-] = 1.0 X 10 - 9 M

Problem Example What is the pH of a solution with a hydrogen-ion concentration of 4.2 x 10 - 10 M? pH = - log [H+] pH = - log (4.2 x 10 - 10) pH = 9.38

Using calculator find the anti log of - 6.35 Problem Example pH of an unknown solution is 6.35. What is its hydrogen-ion concentration? pH = -log [H+] 6.35 = -log [H+] - 6.35 = log [H+] Using calculator find the anti log of - 6.35 [H+] = 4.5 x 10 - 7 M

Problem Example What is the pH of a solution if the [OH-] = 4.0 X10 – 11 M? Kw = [H+] x [OH-] = 1 x 10 -14 [H+] = 1 x 10 -14 ÷ [OH-] [H+] = 1 x 10 -14 ÷ 4.0 x 10 -11 [H+] = 2.5 x 10 - 4 M

Problem Example (con’t) What is the pH of a solution if [OH-] = 4.0 X 10 - 11 M? pH = - log [H+] pH = - log (2.5 x 10 - 4) pH = 3.60

Acid – Base Indicators Indicator - is an acid or a base that undergoes dissociation in a known pH range An indicator is a valuable tool for measuring pH because its acid form and base form have different color in solution. For each indicator, the change from dominating acid form to dominating base form occurs in a narrow range of approximately two pH units. Within this range, the color of the solution is a mixture of the colors of the acid and the base forms. Knowing the pH range over which this color change occurs, can give you a rough estimate of the pH of the solution.

End of Section 19.2

Section 19.3 Hydrogen Ions & Acidity

Strong Acids Acids are classified as strong or weak depending on the degree to which they ionize in water. In general, strong acids are completely ionized in aqueous solution. HNO3 - nitric acid HCl - hydrochloric acid H2SO4 - sulfuric acid HClO4 - perchloric acid HBr - hydrobromic acid HI - hydroiodic acid HCl(g) + H2O(l) H3O +(aq) + Cl ¯(aq)

Weak Acids Weak acids ionize only slightly in aqueous solution. Some Weak Acids Acetic Acid CH3COOH Boric Acid H3BO3 (all three are weak) Phosphoric Acid H3PO4 (all three are weak) Sulfuric Acid HSO4- (first ionization is strong) CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO ¯ (aq) ethanoic acid water hydronium ethanoate ion ion

Acid Strength A strong acid completely dissociates in water ([H3O+] is high). A weak acid remains largely undissociated. ([H3O+] is low).

Equilibrium Constant (Keq) Write the equilibrium-constant expression from the balanced chemical equation. CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO ¯(aq) Keq = [H3O+] x [ CH3COO ¯] [H3COOH] x [H2O] [H2O] constant in dilute solutions

Acid Dissociation Constant (Ka) Ka = Ratio of the concentration of the dissociated form of an acid to the concentration of the undissociated form. CH3COOH (aq) + H2O (l) H3O +(aq) + CH3COO ¯(aq) Acid Dissociation Constant Ka = [H3O+] x [ CH3COO¯] [CH3COOH]

Acid Dissociation Constant (Ka) Acid dissociation constant reflects the fraction of an acid in the ionized form. (Ka sometimes called ionization constant) If the value of the Ka is small, then the degree of dissociation or ionization of the acid in the solution is small. Weak acids – small Ka values Stronger the acid – larger the Ka

Acid Dissociation Constant (Ka) Nitrous acid (HNO2) has a Ka of 4.4 x 10¯4 Acetic acid (CH3COOH) has a Ka of 1.8 x 10¯5 Nitrous acid is more ionized in solution and a stronger acid

Acids Strong Acids Have high [H3O+] Large dissociation constant (Ka) Weak Acids Have low [H3O+] Small dissociation constant (Ka)

Acids Diprotic and triprotic acids lose their hydrogens one at a time. Each ionization reaction has a separate dissociation constant. H3PO4 – 3 separate dissociation constants.

A 0. 012 M solution of formic acid HCOOH is partially ionized A 0.012 M solution of formic acid HCOOH is partially ionized. [H+] is 1.02 x 10-4 M. what is the acid dissociation constant (Ka) of formic acid? Concentration [HCOOH] [H+] [HCOO¯ ] Initial 0.012 Change – 1.02 x 10 - 4 1.02 x 10 - 4 1.02 x 10 – 4 Equilibrium 0.011989 Substitute the equilibrium values into the expression for Ka : [H+] [HCOO¯ ] (1.02 x 10 – 4) (1.02 x 10 – 4) Ka = = = 8.68 x10 - 7 [HCOOH] 0.011989 Classwork page: 610 # 22 & 23

Strong Bases Bases are classified as strong or weak depending on the degree to which they ionize in water. In general, strong bases are completely ionized in aqueous solution. Ca(OH)2 - calcium hydroxide NaOH - sodium hydroxide KOH - potassium hydroxide NaOH(s) + H2O(l) Na +(aq) + OH ¯(aq)

Weak bases ionize only slightly in aqueous solution. Some Weak bases Ammonia NH3 Methylamine CH3NH2 NH3 (aq) + H2O(l) NH4+(aq) + OH ¯ (aq) ammonia water ammonium ion hydroxide ion

Base Dissociation Constant (Kb) Strong bases dissociate completely into metal ions and hydroxide ions in aqueous solution. Some strong bases are not very soluble in water (calcium hydroxide and magnesium hydroxide) Small amounts that do not dissolve dissociate completely Weak bases react with water to form the hydroxide ion and the conjugate acid of the base. NH3(aq) + H2O(l) NH4+ (aq) + OH¯ (aq) Ammonia Water Ammonium Ion Hydroxide ion

Base Dissociation Constant (Kb) NH3(aq) + H2O(l) NH4+(aq) + OH ¯(aq) Ammonia Water Ammonium Ion Hydroxide ion Only about 1% of ammonia is present as NH4+ Base Dissociation Constant (Kb) Kb = [NH4+] x [OH¯ ] [NH3]

Concentration and Strength The words concentrated and dilute indicate how much of an acid or base is dissolved in solution. Number of moles of the acid or base in a given volume. The words strong and weak refer to the extent of ionization or dissociation of an acid or base. How many of the particles ionize or dissociate into ions A sample of HCl added to a large volume of water becomes more dilute, but it is still a strong acid. Vinegar is a dilute solution of a weak acid.

End of section 19.3

Section 19.4 Neutralization Reactions

Acid – Base Reactions If you mix a solution of a strong acid containing hydronium ions with a solution of a strong base that has an equal number of hydroxide ions, a neutral solution results. Final solution has properties that are characteristic of neither an acidic nor a basic solution. HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l) H2SO4 (aq) + 2KOH (aq) K2SO4 (aq) + H2O (l)

Neutralization Reactions Reactions of weak acids and weak bases do not usually produce a neutral solution. In general, reactions with which an acid and a base react in an aqueous solution to produce a salt and water are called neutralization reactions.

Making Salts Prepare potassium chloride by mixing equal molar quantities of hydrochloric acid and potassium hydroxide. HCl + KOH KCl + H2O Heating the solution to evaporate the water will leave the salt potassium chloride. In general, the reaction of an acid with a base produced water and salt

Titration The number of moles of hydrogen ions provided by the acid are equivalent to the number of hydroxide ions provided by the base. HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l) 1 mole 1 mole 1 mole 1 mole H2SO4 (aq) + 2NaOH (aq) Na2SO4 (aq) + 2H2O (l) 1 mole 2 mole 1 mole 2 mole When an acid & base are mixed, the Equivalence point is when the number of moles of hydrogen ions equals the number of moles of hydroxide ions.

Sample Problem How many moles of sulfuric acid are required to neutralize 0.50 mol of sodium hydroxide? H2SO4 (aq) + 2NaOH (aq) Na2SO4 (aq) + 2H2O (l)

Practice Problem How many moles of potassium hydroxide are needed to completely neutralize 1.56 mol of phosphoric acid? H3PO4 (aq) + 3KOH (aq) K3PO4 (aq) + 3H2O (l)

Titration You can determine the concentration of acid or base in a solution by performing a neutralization reaction. You must use an appropriate acid-base indicator to show when neutralization has occurred. In the lab, typically phenolphthalein for acid base neutralization reactions. Solutions that contain phenolphthalein turn from colorless to deep pink as the pH of the solution changes from acidic to basic.

Titration Measured volume of an acid solution of unknown concentration is added to a flask

Titration Several drops of the indicator are added to the solution while the flask is swirled

Titration Measured volumes of the base of known concentration are mixed into the acid until the indicator just barely changes color.

Titration Titration – the process of adding a known amount of solution of known concentration to determine the concentration of another solution. Standard solution – the solution of known concentration End point – the point at which the indicator changes color, the point of neutralization You can also use titration to find the concentration of a base using a standard acid. Equivalence point – the point in a titration where the number of moles of hydrogen ions = number of moles of hydroxide ions.

A 25ml solution of H2SO4 is completely neutralized by 18ml of 1.0M NaOH. What is the concentration of the H2SO4 solution? H2SO4 (aq) + 2NaOH (aq) Na2SO4 (aq) + 2H2O (l)

Practice Problem How many milliliters of 0.45M HCl will neutralize 25.0ml of 1.00M KOH?

Practice Problem What is the molarity of H3PO4 if 15.0 ml is completely neutralized by 38.5 ml of 0.150 M Ca(OH)2?

End of section 19.4

Section 19.5 Salts in Solution

Salt Hydrolysis A salt consists of an anion from an acid and a cation from a base. The salt forms as a result of a neutralization reaction Although solutions of many salts are neutral, some are acidic and others are basic.

Salt Hydrolysis Salt Hydrolysis – the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ion to water. Hydrolyzing salts are usually derived from a strong acid and weak base or from a weak acid and a strong base. In general, salts that produce acidic solutions contain positive ions that release protons to water. Salts that produce basic solutions contain negative ions that attract protons from water.

Salt Hydrolysis CH3COONa (aq) CH3COO ¯ (aq) + Na+ (aq) Sodium ethanoate ethanoate ion sodium ion CH3COONa is the salt from a weak acid CH3COOH and a strong base NaOH In solution, the salt is completely ionized.

Salt Hydrolysis Salt Hydrolysis – the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ion to water. CH3COO¯ (aq) + H2O (l) CH3COOH (aq) + OH ¯(aq) BL base BL acid makes hydrogen-ion hydrogen-ion solution acceptor donor basic This process is called hydrolysis because it splits a hydrogen ion off a water molecule. Resulting solution contains a hydroxide-ion concentration greater than the hydrogen-ion concentration. Thus the solution is basic

Salt Hydrolysis NH4Cl (aq) NH4+ (aq) + Cl ¯ (aq) Ammonium Ammonium ion Chloride ion chloride NH4Cl is the salt from a strong acid (hydrochloric acid, HCl) and a weak base (ammonia, NH3) In solution the salt is completely ionized.

Salt Hydrolysis NH4+ (aq) + H2O (l) NH3 (aq) + H3O+ (aq) BL acid BL base Hydrogen-ion Hydrogen-ion donor acceptor This process is also called hydrolysis because it splits a hydrogen ion off a water molecule. Resulting solution contains a hydrogen-ion concentration greater than the hydroxide-ion concentration. Thus the solution is acidic

Salt Hydrolysis Equivalence Point Strong Acid Strong Base pH= 7 neutral Weak Acid Strong Base pH > 7 basic Strong Acid Weak Base pH < 7 acidic Equivalence point – the point in a titration where the number of moles of hydrogen ions = number of moles of hydroxide ions

Buffers Buffer – a solution in which the pH remains relatively constant when small amounts of acid or base are added. A buffer is a solution of a weak acid and one of its salts, or a solution of a weak base and one of its salts. A buffer solution is better able to resist drastic changes in pH than is pure water.

Buffers A solution of ethanoic acid (CH3COOH) and sodium ethanoate (CH3COONa) is an example of a typical buffer. CH3COOH and CH3COO - (source is the completely ionized CH3COONa) act as reservoirs of neutralizing power.

Buffers CH3COO - (aq) + H+ (aq) CH3COOH (aq) ethanoate ion Hydrogen ion ethanoic acid When an acid is added to the solution, the ethanoate ions act as a hydrogen-ion sponge. CH3COOH (aq) + OH - (aq) CH3COO - (aq) + H2O (l) Ethanoic acid hydroxide ion ethanoate ion water When a base is added to the solution, the ethanoic acid and the hydroxide ions react to produce water and the ethanoate ion.

Buffers The ethanoate ion is not strong enough base to accept hydrogen ions from water extensively. The buffer solution cannot control the pH when too much acid is added, because no more ethanoate ions are present to accept hydrogen ions. Buffer also become ineffective when too much base is added. No more ethanoic acid molecules are present to donate hydrogen ions.

Buffers When too much acid or base is added, the buffer capacity is exceeded. Buffer capacity – the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs.

Buffers When a base is added to a buffered solution, the acidic form removes hydroxide ions from the solution. When an acid is added to a buffered solution, the basic form removes hydrogen ions from the solution.

Buffers & Your Blood Your body function properly only when the pH of your blood lies between 7.35 and 7.45 Your blood contains buffers (hydrogen carbonate ions and carbonic acid) HCO3- (aq) + H+ (aq) H2CO3 (aq) Hydrogen Hydrogen ion Carbonic acid carbonate ion As long as there are hydrogen carbonate ions available, the excess hydrogen ions are removed, and the pH of the blood changes very little.

End of Section Chapter 19