Chapter 11: Acids and Bases

11.9 Buffers

-A buffer solution maintains the pH by neutralizing small amounts of added acid or base. -An acid must be present to react with any OH− added, and a base must be present to react with any H3O+ added.

Components of a Buffer

-A buffer solution: -contains a combination of acid-base conjugate pairs, a weak acid and a salt of its conjugate base, such as HC2H3O2(aq) and C2H3O2−(aq) -has equal concentrations of a weak acid and its salt.

Strong and Weak Acids

-A strong acid completely ionizes (100%) in aqueous solutions. HCl(g) + H2O(l)---> H3O+(aq) + Cl−(aq) -A weak acid dissociates only slightly in water to form a few ions in aqueous solutions. H2CO3(aq) + H2O(l)--->H3O+(aq) + HCO3− (aq)

Acids, Carbonates, and Bicarbonates

-Acids react with carbonates and hydrogen carbonates to produce carbon dioxide gas, a salt, and water: 2HCl(aq) + CaCO3(s)---> CO2(g) + CaCl2(aq) + H2O(l) HCl(aq) + NaHCO3(s)---> CO2(g) + NaCl(aq) + H2O(l)

Reactions of Acids

-Acids react with: -metals to produce salt and hydrogen gas. -bases to produce a salt and water. -bicarbonate and carbonate ions to produce carbon dioxide gas. -A salt is an ionic compound that does not have H+ as the cation or OH− as the anion.

Naming Acids

-Acids with a hydrogen ion (H+) and a nonmetal (or CN−) ion are named with the prefix hydro and end with ic acid. -HCl(aq) hydrochloric acid -Acids with a hydrogen ion (H+) and a polyatomic ion are named by changing the end of the name of the polyatomic ion from -ate to ic acid or ite to ous acid -ClO3− chlorate -ClO2− chlorite -HClO3 -chloric acid -HClO2 -chlorous acid

Basic Solutions

-Adding a base to pure water: -increases the [OH−] -causes the [OH−] to exceed 1.0 × 10−7 M -decreases the [H3O+] [H3O+] < [OH−] The solution is basic.

Acidic Solutions

-Adding an acid to pure water: -increases the [H3O+]. -causes the [H3O+] to exceed 1.0 × 10−7 M. -decreases the [OH−]. [H3O+] > [OH−] The solution is acidic.

Chemistry Link to Health: Antacids

-Antacids are substances that: -are used to neutralize excess stomach acid. -are made of aluminum hydroxide and magnesium hydroxide mixtures. -These hydroxides are not very soluble in water, so the levels of available OH− are not damaging to the intestinal tract.

Writing Dissociation Constants

-As with other dissociation expressions, -the molar concentration of the products is divided by the molar concentration of the reactants. -water is a pure liquid with a constant concentration and is omitted. -the expression is called acid dissociation constant, Ka.

Endpoint of titration

-At the endpoint of the titration: -the moles of base are equal to the moles of acid in the solution. -the concentration of the base is known. -the volume of the base used to reach the endpoint is measured. the molarity of the acid is calculated using the neutralization equation for the reaction.

Calculating the pH of a Buffer

-Because Ka is a constant at a given temperature: -the [H3O+] is determined by the [HC2H3O2]/[C2H3O2−] ratio. -the addition of small amounts of either acid or base changes the ratio of [HC2H3O2]/[C2H3O2−] only slightly. -the changes in [H3O+] will be small and the pH will be maintained. -the addition of a large amount of acid or base may exceed the buffering capacity of the system.

pH scale and [H3O+]

-Because pH is a log scale: -a change of one pH unit corresponds to a tenfold change in [H3O+]. -pH decreases as the [H3O+] increases. pH 2.00 is [H3O+] = 1.0 × 10−2M pH 3.00 is [H3O+] = 1.0 × 10−3 M pH 4.00 is [H3O+] = 1.0 × 10−4 M

Dissociation of a Weak Acid

-Because the dissociation of strong acids in water is essentially complete, the reaction is not considered to be an equilibrium process. -Weak acids partially dissociate in water as the ion products reach equilibrium with the undissociated weak acid molecules. -Formic acid is a weak acid that dissociates in water to form hydronium ion, H3O+, and formate ion, CHO2−.

Buffers and pH Changes

-Buffers can be prepared from conjugate acid-base pairs such as H2PO4−/HPO42− and HPO42−/PO43−, HCO3−/CO32−, or NH4+/NH3. -The pH of the buffer solution will depend on the conjugate acid-base pair chosen.

How buffers work

-Buffers work because: -they resist changes in pH from the addition of an acid or a base. -in the body, they absorb H3O+ or OH− from foods and cellular processes to maintain pH. -they are important in the proper functioning of cells and blood. -they maintain a pH close to 7.4 in blood. -A change in the pH of the blood affects the uptake of oxygen and cellular processes.

Calculating [H3O+] from pH

-Given the pH of a solution, we can reverse the calculation to obtain the [H3O+]. -For whole number pH values, the negative pH value is the power of 10 in the [H3O+] concentration. [H3O+] = 10−pH -For pH values that are not whole numbers, the calculation requires the use of the 10x key, which is usually a 2nd function key.

Function of a Weak Acid in a Buffer

-If a small amount of base is added to this same buffer solution, it is neutralized by the acetic acid, HC2H3O2, which shifts the equilibrium in the direction of the products acetate ion and water. HC2H3O2(aq) + OH−(aq)---> C2H3O2−(aq) + H2O(l) Equilibrium shifts in the direction of the products.

Using Kw to Calculate [H3O+] and [OH−]

-If we know the [H3O+] of a solution, we can use the Kw to calculate the [OH−]. -If we know the [OH−] of a solution, we can use the Kw to calculate the [H3O+]. .

Acids and Hydroxides: Neutralization

-In a neutralization reaction: -an acid reacts with a base to produce salt and water. -the salt formed is the anion from the acid and the cation from the base. HCl(aq) + NaOH(aq)--->NaCl(aq) + H2O(l)

Strong and Weak Acid Dissociation

-In an HCl solution, the strong acid HCl dissociates 100% to form H+ and Cl−. -A solution of the weak acid HC2H3O2 contains mostly molecules of HC2H3O2 and a few ions of H+ and C2H3O2−.

Conjugate Acid-Base Pairs

-In any acid-base reaction, there are two conjugate acid-base pairs: -Each pair is related by the loss and gain of H+. -One pair occurs in the forward direction. -One pair occurs in the reverse direction. Acid and conjugate base pair 1 HA + B A− + BH+ Base and conjugate acid pair 2

Acids and Hydroxides: Neutralization

-In neutralization reactions, one H+ always reacts with one OH−. -If we write the strong acid and strong base as ions, HCl(aq) + NaCl(aq)--->NaOH(aq) + H2O(l) -we see that H+ reacts with OH− to form water, leaving the ions Na+ and Cl− in solution: H+(aq) + Cl−(aq) + Na+(aq) + OH−(aq)---> Na+(aq) + Cl−(aq) + H2O(l) -The overall reaction occurs as the H+ from the acid and OH− from the base form water: H+(aq) + OH−(aq)---> H2O(l) Net ionic equation

Chemistry Link to Health: Buffers in Blood Plasma

-In the body, the concentration of carbonic acid is closely associated with the partial pressure of CO2, PCO2: -If the CO2 level rises, increasing H2CO3, the equilibrium shifts to produce more H3O+, which lowers the pH. This condition is called acidosis. -A lowering of the CO2 level leads to a high blood pH, a condition called alkalosis.

How Buffers Work

-In the buffer with acetic acid (HC2H3O2) and sodium acetate (NaC2H3O2) -the salt produces acetate ions and sodium ions. NaC2H3O2(aq)--->C2H3O2−(aq) + Na+(aq) -the salt is added to provide a higher concentration of the conjugate base C2H3O2− than from the weak acid alone. HC2H3O2(aq) + H2O(l)---> C2H3O2−(aq) + H3O+(aq) Large amount and Large amount

Writing the Dissociation Constant, Kw

-In the equation for the dissociation of water, there is both a forward and a reverse reaction. -In pure water, the concentrations of H3O+ and OH− at 25 °C are each 1.0 × 10−7 M. .

Conjugate Acid-Base Pairs

-In the reaction of NH3 and H2O, -one conjugate acid-base pair is NH3/NH4+. -the other conjugate acid-base pair is H2O/H3O+.

Conjugate Acid-Base Pairs

-In this acid-base reaction: -the first conjugate acid-base pair is HF, which donates H+ to form its conjugate base, F−. -the other conjugate acid-base pair is H2O, which accepts H+ to form its conjugate acid, H3O+. -each pair is related by a loss and gain of H+.

Strong acids

-In water, the dissolved molecules of HA, a strong acid: -dissociate into ions 100%. -produce large concentrations of H3O+ and the anion (A−). -The strong acid HCl dissociates completely into ions: HCl(g)+H2O(l)---> H3O+(aq) + Cl−(aq)

Weak Acids

-In weak acids, only a few molecules dissociate. -Most of the weak acid remains as the undissociated (molecular) form of the acid. -The concentrations of H3O+ and the anion (A−) are small. H2CO3 is a weak acid: H2CO3(aq) + H2O(l)---> H3O+(aq) + HCO3−(aq)

Chemistry Link to Health: Buffers in Blood Plasma

-Kidneys also supply more of the bicarbonate anion, HCO3−, setting up an important buffer system in the body fluid: CO2(g) + H2O(l)--->H2CO3(aq) HCO3−(aq) + H3O+(aq) -Excess H3O+ entering the body fluids reacts with the HCO3−, and excess OH− reacts with the carbonic acid. H2CO3(aq) + H2O(l)---> HCO3−(aq) + H3O+(aq) Equilibrium shifts in the direction of the reactants.

Chemistry Link to Health: Buffers in Blood Plasma

-Kidneys also supply more of the bicarbonate anion, HCO3−, setting up an important buffer system in the body fluid: CO2(g) + H2O(l)-->H2CO3(aq)---> HCO3−(aq) + H3O+(aq) -Excess H3O+ entering the body fluids reacts with the HCO3−: H2CO3(aq) + H2O(l)---> HCO3−(aq) + H3O+(aq) Equilibrium shifts in the direction of the reactants. -Excess OH- entering the body fluids reacts with the H2CO3: H2CO3(aq) + OH−(aq)--->H2O(l) + HCO3−(aq) Equilibrium shifts in the direction of the products.

Diprotic Acids: Sulfuric Acid

-Some strong acids, such as sulfuric acid, are diprotic acids that have two H+, which dissociate one at a time. H2SO4(aq) + H2O(l)---> H3O+(aq) + HSO4−(aq) -Because HSO4− is a weak acid, a second dissociation can take place to produce another H+ and the sulfate ion, SO42− HSO4−(aq) + H2O(l)---> H3O+(aq) + SO42−(aq)

Diprotic Acids: Carbonic Acid

-Some weak acids, such as carbonic acid, are diprotic acids that have two H+, which dissociate one at a time. H2CO3(aq) + H2O(l)---> H3O+(aq) + HCO3−(aq) -Because HCO3− is also a weak acid, a second dissociation can take place to produce another hydronium ion and the carbonate ion, CO32− HCO3−(aq) + H2O(l)---> H3O+(aq) + CO32−(aq)

Strong and Weak Bases

-Strong Bases: Lithium hydroxide LiOH Sodium hydroxide NaOH Potassium hydroxide KOH Rubidium hydroxide RbOH Cesium hydroxide CsOH Calcium hydroxide Ca(OH) 2* Strontium hydroxide Sr(OH) 2* Barium hydroxide Ba(OH)2* *Low solubility, but they dissociate completely -Bases in Household Products: Weak Bases Window cleaner, ammonia, NH3 Bleach, NaOCl Laundry detergent, Na2CO3, Na3PO4 Toothpaste and baking soda, NaHC3 Baking powder, scouring powder, Na2CO3 Lime for lawns and agriculture, CaCO3 Laxatives, antacids, Mg(OH) 2, Al(OH)3 Strong Bases Drain cleaner, oven cleaner, NaOH

Direction of Reaction

-Strong acids have weak conjugate bases that do not readily accept H+. -As the strength of the acid decreases, the strength of its conjugate base increases. -In any acid-base reaction, there are two acids and two bases. -However, one acid is stronger than the other acid, and one base is stronger than the other base. -By comparing their relative strengths, we can determine the direction of the reaction.

Strong Bases

-Strong bases as strong electrolytes: -are formed from metals of Groups 1A (1) and 2A (2). -include LiOH, NaOH, KOH, Ba(OH)2, Sr(OH)2, and Ca(OH)2. -dissociate completely in water. KOH(s)--->K+(aq) + OH−(aq) -are found in household products used to remove grease and unclog drains.

Amphoteric substances

-Substances that can act as both acids and bases are amphoteric or amphiprotic. -For water, the most common amphoteric substance, the acidic or basic behavior depends on the other reactant. -Water donates H+ when it reacts with a stronger base. -Water accepts H+ when it reacts with a stronger acid.

Direction of Reaction: H2SO4

-Sulfuric acid, H2SO4, is a strong acid that readily gives up H+ to water. H2SO4(aq) + H2O(l)---> H3O+(aq) + HSO4−(aq) (stronger acid) + (stronger base) ---> (weaker acid) + (weaker base) -The hydronium ion H3O+ produced is a weaker acid than H2SO4. -The conjugate base HSO4− is a weaker base than water.

Chemistry Link to Health: Buffers in Blood Plasma

-The arterial blood plasma has a normal pH of 7.35 to 7.45. If changes in H3O+ lower the pH below 6.8 or raise it above 8.0, cells cannot function properly and death may result. -In our cells, CO2: -is continually produced as an end product of cellular metabolism. -is carried to the lungs for elimination, and the rest dissolves in body fluids such as plasma and saliva, forming carbonic acid, H2CO3. -As a weak acid, carbonic acid dissociates to give bicarbonate, HCO3−, and H3O+.

Working Buffers

-The buffer described here consists of about equal concentrations of acetic acid (HC2H3O2) and its conjugate base, acetate ion (C2H3O2−). -Adding H3O+ to the buffer reacts with the salt, C2H3O2−, whereas adding OH− neutralizes the acid HC2H3O2. -The pH of the solution is maintained as long as the added amounts of acid or base are small compared to the concentrations of the buffer components.

Direction of Reaction: CO32−

-The carbonate ion from carbonic acid, H2CO3, reacts with water. -Water donates one H+ to carbonate, CO32− to form HCO3− and OH−. -From Table 11.3, we see that HCO3− is a stronger acid than H2O. -We also see that OH− is a stronger base than CO32−. -To reach equilibrium, the strong acid and strong base react in the direction of the weaker acid and weaker base. CO32− (aq) + H2O(l)---> OH−(aq) + HCO3−(aq) (weaker acid) + (weaker base)---> (stronger base) + (stronger acid)

Dissociation Constant, Kw

-The ion product constant for water, Kw, is defined as: -the product of the concentrations of H3O+ and OH−. -equal to 1.0 × 10−14 at 25 °C (the concentration units are omitted). -When: -[H3O+] and [OH−] are equal, the solution is neutral. -[H3O+] is greater than the [OH−], the solution is acidic. -[OH−] is greater than the [H3O+], the solution is basic.

pH Measurement

-The pH of a solution can be determined using (a) a pH meter, (b) pH paper, and (c) indicators that turn different colors corresponding to different pH values. -On the pH scale, values below 7.0 are acidic, a value of 7.0 is neutral, and values above 7.0 are basic.

The pH Scale

-The pH of a solution is commonly measured using: -a pH meter in the laboratory. -pH paper, an indicator that turns specific colors at a specific pH value. -The pH of a solution is found by comparing the colors of indicator paper to a chart.

11.6 The pH Scale

-The pH scale is used to describe the acidity of solutions. -A dipstick is used to measure the pH of urine.

Calculating the pH of Solutions

-The pH scale: -is a logarithmic scale that corresponds to the [H3O+] of aqueous solutions. -is the negative logarithm (base 10) of the [H3O+]. pH = −log[H3O+] -To calculate the pH, the negative powers of 10 in the molar concentrations are converted to positive numbers. If [H3O+] is 1.0 × 10−2 M, pH = −log[1.0 × 10−2 ] = −(−2.00) = 2.00

pH: significant figures

-To determine the number of significant figures in the pH value: -the number of decimal places in the pH value is the same as the number of significant figures in the coefficient of [H3O+]. -the number to the left of the decimal point in the pH value is the power of 10.

Chemistry Link to Health: Buffers in Blood Plasma

-To maintain the normal blood plasma pH (7.35 to 7.45), -the ratio of [H2CO3]/[HCO3−] needs to be about 1 to 10. -concentrations of 0.0024 M H2CO3 and 0.024 M HCO3− work to maintain that pH.

Naming Bases

-Typical Arrhenius bases are named as hydroxides. -NaOH: sodium hydroxide -KOH: potassium hydroxide -Ba(OH)2: barium hydroxide -Al(OH)3: aluminum hydroxide

Buffers and pH Changes

-Using a common phosphate buffer for biological specimens, we can look at the effect of using different ratios of [H2PO4−/HPO42−] on the [H3O+] and pH. The Ka of H2PO4− is 6.2 × 10−8. -The equation and the [H3O+] are written as follows: H2PO4−(aq) + H2O(l)---> H3O+(aq) + HPO42−(aq)

Dissociation Constant of Water, Kw

-Water is amphoteric—it can act as an acid or a base. In water, -H+ is transferred from one H2O molecule to another. -one water molecule acts as an acid, while another acts as a base. -equilibrium is reached between the conjugate acid-base pairs. .

11.3 Strengths of Acids and Bases

-Weak acids only partially dissociate in water. -Hydrofluoric acid, HF, is the only halogen that forms a weak acid.

Weak Bases

-Weak bases are weak electrolytes: -that are poor acceptors of H+ ions. -produce very few ions in solution. -include ammonia. NH3(g) + H2O(l)---> NH4+(aq) + OH−(aq) Ammonia and Ammonium hydroxide

Function of Conjugate Base in a Buffer

-When a small amount of acid is added, the additional H3O+ combines with the acetate ion, C2H3O2−, causing the equilibrium to shift in the direction of the reactants, acetic acid and water. -The acetic acid produced contributes to the available weak acid. HC2H3O2(aq) + H2O(l)---> C2H3O2− (aq) + H3O+(aq) Equilibrium shifts in the direction of the reactants.

Buffers

-When an acid or a base is added to water, the pH changes drastically. -In a buffer solution, the pH is maintained; pH does not change when acids or bases are added.

Acid Dissociation Constant, Ka

-When the value of the Ka -is small, the equilibrium lies to the left, favoring the reactants. -is large, the equilibrium lies to the right, favoring the products. -Weak acids have small Ka values, while strong acids have very large Ka values.

Base Dissociation Constant, Kb

-When the value of the Kb, -is small, the equilibrium lies to the left, favoring the reactants. -is large, the equilibrium lies to the right, favoring the products. -The stronger the base, the larger the Kb value. CH3—NH2(aq) + H2O(l)---> CH3—NH3+(aq) + OH−(aq) -The concentration of water is omitted from the base dissociation constant expression.

Write the acid dissociation constant expression for nitrous acid, HNO2.

-Write the acid dissociation constant expression for nitrous acid, HNO2. -The equation for the dissociation of nitrous acid, a weak acid, is HNO2 (aq) + H2O(l)---> H3O+(aq) + NO2−(aq) -The acid dissociation constant is the molar concentration of the products divided by the molar concentration of the reactants.

The Ka for acetic acid, HC2H3O2, is 1.8 × 10-5. What is the pH of a buffer prepared with 1.0 M HC2H3O2 and 1.0 M C2H3O2−? HC2H3O2(aq) + H2O(l)---> C2H3O2−(aq) + H3O+(aq)

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Study check

1. HBr Br−, bromide C. hydrobromic acid The name of an acid with a hydrogen ion (H+) and a nonmetal uses the prefix hydro and ends with ic acid. 2. H2CO3 CO32−, carbonate A. carbonic acid An acid with a hydrogen ion (H+) and a polyatomic ion ending in ate is called an ic acid. 3. HBrO2 BrO2−, bromite C. bromous acid An acid with a hydrogen ion (H+) and a polyatomic ion ending in ite is called an ous acid.

Study Check

1. Write the conjugate base for each of the following acids: A. HBr: H+ + Br− B. H2S: H+ + HS− C. H2CO3: H+ + HCO3− 2. Write the conjugate acid of each of the following bases: (Add an H+ to each base to get the conjugate acid.) A. NO2− +H+: HNO2 B. NH3 +H+: NH4+ C. OH− + H+: H2O

11.2 Brønsted-Lowry Acids and Bases

According to the Brønsted-Lowry theory: -an acid is a substance that donates H+. -a base is a substance that accepts H+. -In the reaction of ammonia and water: -NH3 acts as the base that accepts H+. -H2O acts as the acid that donates H+.

Acids and Metals

Acids react with metals to produce hydrogen gas and the salt of the metal. 2K(s) + 2HCl(aq)--->2KCl(aq) + H2(g) Zn(s) + 2HCl(aq)--->ZnCl2(aq) + H2(g)

Arrhenius Acids

Arrhenius acids: -produce hydrogen ions (H+) when they dissolve in water. HCl(g)--->H+(aq) + Cl−(aq) -are also electrolytes, because they produce H+ in water. -have a sour taste. -turn blue litmus red. -corrode some metals.

Arrhenius Bases

Arrhenius bases: -produce hydroxide ions (OH−) in water. -taste bitter or chalky. -are also electrolytes, because they produce hydroxide ions (OH−) in water. -feel soapy and slippery. -turn litmus indicator paper -blue and phenolphthalein indicator pink.

pH calculation

Aspirin, which is acetylsalicylic acid, was the first nonsteroidal anti-inflammatory drug used to alleviate pain and fever. If a solution of aspirin has a [H3O+] = 1.7 × 10−3 M, what is the pH of the solution? .

Calculating the pH of a Buffer

By rearranging the Ka expression to give [H3O+], we can obtain the ratio of the acetic acid/acetate buffer and calculate the pH. .

Acids and bases

Clinical laboratory technicians prepare specimens for the detection of cancerous tumors and type blood samples for transfusions. They must also interpret and analyze the test results, which are then passed on to the physician.

11.7 Reactions of Acids and Bases

Gastric acid contains HCl and is produced by parietal cells that line the stomach. When protein enters the stomach, HCl is secreted until the pH reaches 2, the optimum pH for digestion.

Identify each of the following as a strong or weak acid or base:

HBr: strong acid HNO2: weak acid NaOH: strong base H2SO4: strong acid Cu(OH)2: weak base

What is the molarity of an HCl solution if 18.5 mL of 0.225 M NaOH is required to neutralize 0.0100 L of HCl?

HCl(aq) + NaOH(aq)--->NaCl(aq) + H2O(l) .

Study Check

Identify each as a characteristic of A. an acid or B. a base. A-has a sour taste B-produces OH− in aqueous solutions B-has a chalky taste A, B- is an electrolyte A-produces H+ in aqueous solutions

study check

Identify each solution as acidic, basic, or neutral. HCl with a pH = 1.5: acidic pancreatic fluid, [H3O+] = 1 × 10−8M: basic Sprite soft drink, pH = 3.0: acidic pH = 7.0: neutral [OH−] = 3 × 10−10 M: acidic [H3O+ ] = 5 × 10−12: basic

Study check

Identify the sets that contain acid-base conjugate pairs. 1. HNO2, NO2− acid, conjugate base 2. H2CO3, CO32−not acid-base conjugate pair 3. HCl, ClO4− not acid-base conjugate pair 4. HS−, H2S base, conjugate acid 5. NH3, NH4+ base, conjugate acid

Pure Water Is Neutral

In pure water, the ionization of water molecules produces small but equal quantities of H3O+ and OH− ions. [H3O+] = 1.0 × 10−7 M [OH−] = 1.0 × 10−7 M [H3O+] = [OH−] Pure water is neutral.

Study check

Match the formulas of acids and bases with their names: 1. _D__ HNO2 D. nitrous acid 2. _E__ Ca(OH)2 E. calcium hydroxide 3. _B__ H2SO4 B. sulfuric acid 4. _A__ HIO3 A. iodic acid 5. _C__ NaOH C. sodium hydroxide

Strength versus Equilibrium Position

Table 11.5 summarizes the characteristics of acids and bases in terms of strength and equilibrium position.

11.5 Dissociation of Water

The equilibrium reached between the conjugate acid-base pairs of water produces both H3O+ and OH−.

Indicator

The indicator phenolphthalein: -is added to identify the endpoint. -turns pink when the solution is neutralized.

The pH Scale

The pH of a solution: -is used to indicate the acidity of a solution. -has values that usually range from 0 to 14. -is acidic when the values are less than 7. -is neutral at a pH of 7. -is basic when the values are greater than 7.

11.8 Acid-Base Titration

The titration of an acid. A known volume of an acid is placed in a flask with an indicator and titrated with a measured volume of a base solution, such as NaOH, to the neutralization endpoint.

Acid-Base Titration

Titration: -is a laboratory procedure used to determine the molarity of an acid. -uses a base such as NaOH to neutralize a measured volume of an acid. -requires a few drops of an indicator such as phenolphthalein to identify the endpoint. -In the following titration, a specific volume of acidic solution is titrated to the endpoint with a known concentration of NaOH.

study check

Using Table 11.3, identify the stronger acid in each pair. A. HNO2 or H2S HNO2: is the stronger acid. B. HCO3− or HBr: HBr is the stronger acid. C. H3PO4 or H3O+ H3O+: is the stronger acid.

11.4 Dissociation Constants for Acids and Bases

Write the expression for the dissociation constant of a weak acid or weak base.