Chapter 8 Lecture General Organic and Biological Chemistry

Chapter 8 Lecture General, Organic, and Biological Chemistry: An Integrated Approach Chapter 8 Laura Frost, Todd Deal and Karen Timberlake Acids, Bases, and Buffers in the Body by Richard Triplett © 2011 Pearson Education, Inc.

Chapter Outline 8. 1 Acids and Bases—Definitions 8. 2 Strong Acids and Bases 8. 3 Chemical Equilibrium 8. 4 Weak Acids and Bases 8. 5 p. H and the p. H Scale 8. 6 p. Ka 8. 7 Amino Acids: Common Biological Weak Acids 8. 8 Buffers: An Important Property of Weak Acids and Bases © 2011 Pearson Education, Inc.

Introduction • Citrus fruits taste sour because they contain acid. • Our stomach produces acid to aid in digestion, and our muscles produce lactic acid when we exercise. • An acid can be neutralized by a base. Soaps are mild bases, and, like other bases, feel slippery to the touch. © 2011 Pearson Education, Inc.

Introduction, Continued • p. H refers to the acidity of a solution. • Amino acids, which are the building blocks of proteins, will change form if the acidity of a solution changes. Proteins change their shape and functionality if the p. H of a solution is changed. • Our bodily fluids contain compounds that maintain p. H. These compounds are called buffers. © 2011 Pearson Education, Inc.

8. 1 Acids and Bases—Definitions Acids • The Swedish chemist, Svante Arrhenius, first described acids as substances that dissociate, producing hydrogen ions (H+) when dissolved in water. • The hydrogen ions give acids their sour taste and allow acids to corrode some metals. © 2011 Pearson Education, Inc.

8. 1 Acids and Bases—Definitions, Continued • In the early twentieth century, Brønsted and Lowry redefined acids as a compound that donates a proton. • H+ is the same as a proton since hydrogen has lost its electron. • A free proton rarely exists in an aqueous solution. The proton is attracted to the partial negative charge on the oxygen atom of water. © 2011 Pearson Education, Inc.

8. 1 Acids and Bases—Definitions, Continued The attraction is strong and the oxygen atom in water forms a third covalent bond, giving water a positive, ionic charge creating the hydronium ion, H 3 O+. © 2011 Pearson Education, Inc.

8. 1 Acids and Bases—Definitions, Continued Bases • Arrhenius described bases as compounds that dissociate to form a metal ion and a hydroxide ion (OH-) when dissolved in water. Arrhenius bases are formed from Group 1 A and 2 A metals. Hydroxide bases are characterized by a bitter taste and a slippery feel. • The Brønsted–Lowry definition of a base is a compound that accepts a proton. © 2011 Pearson Education, Inc.

8. 1 Acids and Bases—Definitions, Continued Sodium hydroxide, Na. OH, is an example of

8. 1 Acids and Bases—Definitions, Continued Acids and Bases Are Both Present in Aqueous Solution • The Brønsted–Lowry definitions of acids and bases imply that a proton is transferred in an acidic or basic solution. • Water can act as an acid or base by donating or accepting a proton. For example, when a hydrochloric acid solution is prepared, water accepts a proton, and is acting as a base. © 2011 Pearson Education, Inc.

8. 1 Acids and Bases—Definitions, Continued In another reaction, ammonia (NH 3) reacts with

8. 2 Strong Acids and Bases • Acids and bases are classified by their ability to donate or accept protons, respectively. • Strong acids completely dissociate into ions when placed in water, forming hydronium ions and anions. Examples of strong acids are as follows: © 2011 Pearson Education, Inc.

8. 2 Strong Acids and Bases, Continued Weak acids only partially dissociate when placed in water. Acetic acid, the main component of vinegar, is an example of a weak acid. © 2011 Pearson Education, Inc.

8. 2 Strong Acids and Bases, Continued • Sodium hydroxide, Na. OH, is a strong base. It is used in household products such as oven cleaners and drain openers. • Na. OH and other bases like Li. OH, KOH, and Ca(OH)2 are strong bases that completely dissociate in water. They dissociate to give a metal ion and a hydroxide ion. • Weak bases, like weak acids, only partially dissociate in water. © 2011 Pearson Education, Inc.

8. 2 Strong Acids and Bases, Continued Neutralization • When a strong acid and strong base are mixed, both completely dissociate to form ions in water. • For example, when HCl and Na. OH are mixed, sodium ions and chloride ions are formed, as well as hydroxide ions and hydronium ions. • The protons in the hydronium ion are attracted to the hydroxide ion to form water molecules. © 2011 Pearson Education, Inc.

8. 2 Strong Acids and Bases, Continued • As a result of this chemical reaction, a lot of heat is produced and is considered to be an exothermic (exo means “to give off”; thermo means “heat”) reaction. • The sodium and chloride ions remain in solution. • The reaction of a strong acid with a strong base produces a salt and water. © 2011 Pearson Education, Inc.

8. 2 Strong Acids and Bases, Continued The formation of water and a salt

8. 2 Strong Acids and Bases, Continued Antacids • Antacids are substances that neutralize excess stomach acid. Some antacids are bases that are not very soluble in water. • Aluminum hydroxide and magnesium hydroxide are examples of antacids that are not very soluble in water. They are used in combination to prevent unpleasant side effects. • Calcium carbonate is an antacid that will cause an increase in serum calcium. It is not recommended for people with peptic ulcers or for those that have a tendency to form kidney stones. © 2011 Pearson Education, Inc.

8. 2 Strong Acids and Bases, Continued • Sodium bicarbonate is an antacid that will affect the acidity level of blood and will elevate sodium levels in bodily fluids. • The following table shows some antacid preparations. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium • After forming products, some chemical reactions will reverse and reform reactants. These type of reactions are reversible reactions. • The generation of ammonia is a reversible reaction. Once ammonia is formed, the reaction will reverse and reform nitrogen and hydrogen. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued • The rate of ammonia formation and reforming nitrogen and hydrogen will eventually become equal. This is called chemical equilibrium. • An equilibrium arrow is used in this type of reaction to indicate that both the forward and reverse reactions take place simultaneously. • Because the rates of forward and reverse reactions are equal, there is no net change in the amounts of product or reactants. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued The Equilibrium Constant, K In this reaction, if we measure the concentrations of ammonia, nitrogen, and hydrogen, the fraction of products to reactants would be a constant value. This constant value is called the equilibrium constant, K. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued • K is defined as: • The brackets, [ ], mean “concentration of. ” • This equation states that the equilibrium constant is equal to the molar concentration of the products divided by the molar concentrations of the reactants. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued • The expression for K for the generation of ammonia is: • The superscripts in the expression come from the coefficients found in the chemical equation. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued • For an equilibrium reaction of the form: •

8. 3 Chemical Equilibrium, Continued • Only substances whose concentrations change appear in the equilibrium expression. • Substances like solid and pure liquids have constant concentrations, so they do not appear in the expression. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued • Values for K vary greatly depending on temperature and reaction. • If K has a value of 1, the ratio of products: reactants is 1: 1 or the [products] = [reactants]. • A K value greater than (>) 1 means the amount of products is larger than the amount of reactants or [products] > [reactants]. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued • A K value less than (<) 1 means that the amount of reactants is larger than the amount of products or [products] < [reactants]. • The following table summarizes the interpretations of K values. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued Effect of Concentration on Equilibrium— Le Châtelier’s Principle • Reconsider the production of ammonia. • What would happen if more nitrogen was injected in the reaction vessel? © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued • According to Le Châtelier’s principle, if stress is applied to the equilibrium, the rates of the forward and reverse reaction will change to relieve the stress, and equilibrium will be regained. • If more N 2 is added, the rate of the forward reaction will increase, shifting the equilibrium to produce more products. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued If one of the reactants, like H 2, is removed, the reverse reaction will increase faster than the forward reaction, allowing H 2 to be replenished. The equilibrium will shift to the left, forming more of the reactants. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued Effect of Temperature on Equilibrium • What affect would the change in temperature have on the ammonia reaction? • The reaction is known as an exothermic reaction (produces heat) as opposed to an endothermic reaction (absorbs heat from the surrounding). © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued • The reaction can be written as follows to signify that heat is produced: • If the temperature of the reaction is raised, the rate of the reverse reaction increases to offset the stress of adding heat. This causes the equilibrium to shift to the left. • If the reaction is cooled, the rate of the forward reaction will increase to replenish the heat produced, shifting the equilibrium to the right. © 2011 Pearson Education, Inc.

8. 3 Chemical Equilibrium, Continued The effects of temperature changes on ammonia production are

8. 3 Chemical Equilibrium, Continued Consider an endothermic reaction in which heat is absorbed

8. 3 Chemical Equilibrium, Continued The effects of changes on equilibrium are summarized here.

8. 4 Weak Acids and Bases Equilibrium • The principles of equilibrium apply to weak acids and bases because they only partially dissociate into ions. • The dissociation of acetic acid into acetate ions and hydronium ions is shown as: © 2011 Pearson Education, Inc.

8. 4 Weak Acids and Bases, Continued • The equilibrium constant expression for this reaction would be: • Pure liquids like water are present in large amounts, do not change significantly, and are not included in the equilibrium expression. © 2011 Pearson Education, Inc.

8. 4 Weak Acids and Bases, Continued The Equilibrium Constant, Ka • Weak acids dissociate much less than 100% and have an equilibrium constant called an acid dissociation constant, Ka. • The strength of a weak acid can be determined from the Ka value. The larger the Ka value, the stronger the acid (the more protons dissociated). © 2011 Pearson Education, Inc.

8. 4 Weak Acids and Bases, Continued This table shows Ka values for substances

8. 4 Weak Acids and Bases, Continued Two common organic functional groups that act

8. 4 Weak Acids and Bases, Continued Conjugate Acids and Bases • The Brønsted–Lowry theory states that the reaction between an acid and base involves a proton transfer. If a weak acid is mixed with water, water will act as a base and accept a proton from the weak acid. • Consider the dissociation of acetic acid. © 2011 Pearson Education, Inc.

8. 4 Weak Acids and Bases, Continued • Acetic acid, CH 3 COOH, donates a proton to a molecule of water to form the hydronium ion, H 3 O +. • After the donation of a proton, CH 3 COO-, a carboxylate called acetate anion remains and is called the conjugate base of CH 3 COOH. • The hydronium ion, H 3 O+, is the conjugate acid of water, which is acting as a base. © 2011 Pearson Education, Inc.

8. 4 Weak Acids and Bases, Continued • Molecules or ions related by the loss or gain of a proton are referred to as conjugate acid–base pairs. • The functional groups carboxylic acid and carboxylate are conjugate acid–base pairs of the same functional group. • Weak acids are usually designated HA and their conjugate bases as A-. © 2011 Pearson Education, Inc.

8. 4 Weak Acids and Bases, Continued • A weak base, such as an amine, will accept a proton to form a protonated amine. In this case, water acts as an acid. • The protonated amine and amine are conjugate acid–base pairs of the same functional group. © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale The Autoionization of Water, Kw • Water can act as either a weak acid or a base depending on whether a base or acid is present in solution. • In pure water, the water molecules spontaneously react with each other as shown. © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale, Continued • This reaction is called the autoionization of water. • The equilibrium constant expression for water, Kw, can be written as: Kw = [OH-][H 3 O+] • Pure water will not appear in this expression. • The Kw for water is 1 x 10 -14, which means there are such small amounts of ions in pure water that water will not conduct electricity. © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale, Continued [H 3 O+], [OH-], and p. H • Pure water has an equal number of hydroxide and hydronium ions, so [H 3 O+] = [OH-]. At 25 o. C both these values are 1 x 10 -7 M. • When these concentrations are equal, the solution is said to be neutral. • If acid is added to water, there is an increase in [H 3 O+] and a decrease in [OH-], which makes the solution acidic. © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale, Continued If a base is

8. 5 p. H and the p. H Scale, Continued • Most aqueous solutions are not neutral, meaning they have unequal concentrations of H 3 O+ and OH-. • The range of hydronium ion in an aqueous solution can range from 18 M to 1 x 10 -14 M. • Because of this large range, it is more useful to compare [H 3 O+] by a log scale because it gives values that fall between 0 and 14. © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale, Continued • Values between 1 and 14 describes the p. H scale. • Mathematically, p. H can be determined from the [H 3 O+] using the following expression: p. H = - log [H 3 O+] © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale, Continued The relationship between p.

8. 5 p. H and the p. H Scale, Continued Measuring p. H The p. H of a solution can be measured electronically using a p. H meter. It can also be measured visually using p. H paper, which is embedded with indicators that change color based on the p. H of a solution. © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale, Continued Math Matters: Logarithms • The exponent of 10 is the logarithm, or log, of these numbers. For example, the log of 102 = 2; log of 103 = 3, etc. • The purpose of the log function is to show the number of tens places included in a really large or a really small number. © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale, Continued Math Matters: Logarithms, Continued Negative numbers do not have a log value. The log function is described as the log of base 10 because the logs of integers come from numbers that are whole number multiples of 10. © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale, Continued Math Matters: Inverse Logarithms • Suppose we want to solve for x in the following: 4 = log x • To solve for x, the equation must be rearranged. • To do this, we must take the inverse log of both sides of the equation, and applying the inverse log function, we can solve for x: © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale, Continued Calculating p. H • Calculate the p. H of a 0. 050 M HCl solution. Because strong acids completely dissociate in solution, [HCl] = [H 3 O+], • The number of significant figures in the [H 3 O+] will be the number of decimal places in the p. H value. © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale, Continued Calculating [H 3 O+] • If we measure the p. H of a solution to be 3. 00, how do we find the corresponding [H 3 O+] if we know that p. H = - log [H 3 O+]? • Multiply both sides of the equation by negative 1 and the inverse log function, INV log. The inverse log function is 10 x. © 2011 Pearson Education, Inc.

8. 5 p. H and the p. H Scale, Continued • Solving the equation for [H 3 O+], we have: INV log (-) p. H = [H 3 O+] or alternatively, 10 -p. H = [H 3 O+] • The solution for the [H 3 O+] of a p. H 3. 00 solution is: INV log (-) 3. 00 or 10 -3. 00 = [H 3 O+] 1. 0 x 10 -3 M = [H 3 O+] • The number of significant figures in the [H 3 O+] is: © 2011 Pearson Education, Inc.

8. 6 p. Ka • To determine which acids are strongest, the Ka values can be compared. • As seen with the p. H scale, it is easier to compare whole numbers than those in scientific notation. The p. Ka values are used to measure acid strength. • The lower the p. Ka value, the stronger the acid. © 2011 Pearson Education, Inc.

8. 6 p. Ka, Continued In this table, a comparison is shown between Ka and p. Ka. Those Ka values closer to 1 are the stronger acids. The p. Ka values for the stronger acids are low. © 2011 Pearson Education, Inc.

8. 6 p. Ka, Continued The Relationship between p. H and p. Ka • p. H and p. Ka are two numbers associated with weak acid solutions. • p. Ka gives us the fraction of acid molecules that will dissociate, that is, p. Ka gives the ratio of conjugate base and hydronium ion to weak acid. • p. H tells how much hydronium ion is present in solution. © 2011 Pearson Education, Inc.

8. 6 p. Ka, Continued • At constant temperature, p. Ka does not change if an acid or a base is added to a weak acid solution, however the p. H does change if an acid or a base is added to a weak acid solution. • Consider an acetic acid solution at three different p. H’s: 1. A p. H below the p. Ka 2. A p. H equal to p. Ka 3. A p. H above the p. Ka © 2011 Pearson Education, Inc.

8. 6 p. Ka, Continued A p. H Below the p. Ka • Extra H 3 O+ has been added to the equilibrium solution. • The equilibrium has shifted to the left, meaning there is more acetic acid present than its conjugate base, acetate. © 2011 Pearson Education, Inc.

8. 6 p. Ka, Continued p. H Equal to p. Ka • The Ka

8. 6 p. Ka, Continued p. H Is Above the p. Ka • Less equilibrium concentration of H 3 O+ is present. • The equilibrium has shifted to the right, producing more acetate than acetic acid. © 2011 Pearson Education, Inc.

8. 6 p. Ka, Continued The relationship between acid (HA), conjugate base (A-), p.

8. 7 Amino Acids: Common Biological Weak Acids • The molecule shown below is alanine, with functional groups identified as a carboxylate anion and a protonated amine. • Alanine belongs to a class of molecules called amino acids, which are the building blocks of proteins. © 2011 Pearson Education, Inc.

8. 7 Amino Acids: Common Biological Weak Acids, Continued • Why is alanine shown with a carboxylate group and a protonated amine rather than a carboxylic acid group and amine group, respectively? • The p. Ka value for the carboxylic acid is 2. 3, below physiological p. H, so the conjugate base form predominates. The p. Ka value for the amine group is 9. 7, above physiological p. H, so the conjugate acid predominates. • This ionic form with no net ionic charge is called a zwitterion. © 2011 Pearson Education, Inc.

8. 7 Amino Acids: Common Biological Weak Acids, Continued • Amino acids and other biological molecules contain more than one weak acid group and have more than one p. Ka value. • With more than one p. Ka value, these molecules exist in different acid/conjugate base forms depending on the p. H of the solution. • These molecules have a unique p. H at which the zwitterion is the predominate form. This point is called the isoelectric point (p. I). © 2011 Pearson Education, Inc.

8. 7 Amino Acids: Common Biological Weak Acids, Continued • At the p. I for alanine, the negative charge on the carboxylate group is balanced by the positive charge of the ammonium ion, and the net charge is zero. • The p. I for alanine is 6. 0 and is halfway between the p. Ka values for the protonated amine and carboxylic acid. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases • Our bodies operate under strict conditions of temperature, concentration, and p. H. • How do our bodies maintain p. H in our bloodstream when we consume a variety of foods at different p. H’s? • Our bodies contain solutions of weak acids, containing both acids and conjugate bases, to help neutralize incoming acids and bases. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued • A buffer is a solution that contains both a weak acid and its conjugate base, or a base and its conjugate acid. • A buffer will resist a change in p. H if small amounts of an acid or a base are added. • Buffers are what helps our body maintain the proper p. H in our bloodstream when we consume a variety foods at different p. H’s. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued • The bicarbonate buffer system is the main buffer system in our blood. • Dissolved CO 2, produced during cellular respiration, is equilibrated through carbonic acid into bicarbonate ions prior to exhalation at the lungs. • The intermediates, carbonic acid and water, are often omitted since they are short lived in the reaction. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued • The bicarbonate buffer system in our bloodstream is shown in the following figure. • The bicarbonate buffer system can be denoted by showing the acid and its conjugate base like this: H 2 CO 3/HCO 3 -. Sometimes the conjugate base is shown as an ionic compound like this: H 2 CO 3/Na. HCO 3. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued Maintaining Physiological p. H with Bicarbonate Buffer: Homeostasis • The bicarbonate system helps our bodies maintain its optimal physiological p. H. • The ability of an organism to maintain its internal environment by adjusting such factors as p. H, temperature, and solute concentration is called homeostasis. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued Changes in Ventilation Rate • In normal breathing, carbon dioxide is removed from the bloodstream and blood p. H is maintained. • A person who hypoventilates may fail to remove enough carbon dioxide due to shallow breathing causing carbon dioxide to build up in the bloodstream. • A buildup of carbon dioxide in the bloodstream makes the blood more acidic, a condition that is known as respiratory acidosis. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued Changes in Ventilation Rate, Continued • Individuals suffering this condition must be treated to raise blood p. H back to normal. • A bicarbonate solution can be administrated intravenously. This will drive the equilibrium to the left, when the excess bicarbonate present reacts with the excess acid. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued Changes in Ventilation Rate, Continued • A person who hyperventilates will exhale too much carbon dioxide from the lungs. • This will draw H 3 O+ from the bloodstream, making the blood more basic. This condition is known as respiratory alkalosis. • This condition necessitates getting more carbon dioxide back into the bloodstream, which can be done by having the person breathe into a paper bag. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued Changes in Ventilation Rate, Continued • This will enrich carbon dioxide in the bloodstream, shifting the equilibrium back to the right, thereby producing more H 3 O+. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued Changes in Metabolic Acid Production • Chemical reactions in the body change the p. H of blood by producing too much or too little H 3 O +. • Since diabetics use less glucose for energy production, they rely on fatty acids as an energy source. A by-product of fatty acid breakdown is acid production. • An imbalance caused by chemical reactions in the body is termed metabolic acidosis. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued Changes in Metabolic Acid Production, Continued • A treatment for this condition is to administer bicarbonate to neutralize the excess acid so that more carbon dioxide can be exhaled. • If the blood is basic, the body is losing too much acid, a condition known as metabolic alkalosis. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued Changes in Metabolic Acid Production, Continued • Metabolic alkalosis occurs under conditions of excess vomiting. • To lower the p. H back to normal, ammonium chloride (a weak acid) can be administrated. © 2011 Pearson Education, Inc.

8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued A summary

Chapter Summary 8. 1 Acids and Bases—Definitions • Arrhenius defines acids as producing H+ and defines bases as producing OH-. • Brønsted–Lowry defines acids as H+ donors and bases as H+ acceptors. • H+, a proton, forms a hydronium ion with a water molecule. © 2011 Pearson Education, Inc.

Chapter Summary, Continued 8. 2 Strong Acids and Bases • Strong acids and bases completely dissociate in solution. • Water can act as either an acid or base. • Neutralization reactions occur when a strong acid combines with a strong base. Products are a salt and water. • Antacids are basic and neutralize stomach acid. © 2011 Pearson Education, Inc.

Chapter Summary, Continued 8. 3 Chemical Equilibrium • When forward and reverse reactions occur at the same rate, a chemical equilibrium is established. • K, equilibrium constant, defines the extent of a chemical reaction. • If a chemical reaction at equilibrium is stressed, the reaction regains equilibrium according to Le Châtelier’s principle. © 2011 Pearson Education, Inc.

Chapter Summary, Continued 8. 3 Chemical Equilibrium, Continued • Endothermic reactions require heat in order to occur. • Exothermic reactions give off heat when they occur. © 2011 Pearson Education, Inc.

Chapter Summary, Continued 8. 4 Weak Acids and Bases • Weak acids partially ionize in solution and have an equilibrium constant called the acid dissociation constant, Ka. • Weak acids produce a conjugate base when they dissociate, and weak bases produce a conjugate acid when they dissociate. These are referred to as conjugate acid–base pairs. • Strong acids have high Ka values. © 2011 Pearson Education, Inc.

Chapter Summary, Continued 8. 5 p. H and the p. H Scale • When water ionizes it produces hydronium ions and hydroxide ions. • When these ions are of equal concentration, a neutral solution exists. • An excess of hydronium ions produces an acidic solution and an excess of hydroxide ions produces a basic solution. © 2011 Pearson Education, Inc.

Chapter Summary, Continued 8. 5 p. H and the p. H Scale, Continued • The p. H scale is a measure of acidity with values between 0– 14. Neutral solutions have a p. H of 7, acidic solutions have values less than 7, and basic solutions have values greater than 7. • p. H can be measured with a p. H meter or using p. H paper. • Mathematically written, p. H = -log [H 3 O+]. © 2011 Pearson Education, Inc.

Chapter Summary, Continued 8. 6 p. Ka • p. Ka is a constant for a specific weak acid at a certain temperature. • If the p. H of a solution is the same as the p. Ka value of a weak acid, then the acid and conjugate base forms are present in equal amounts. • If p. H is higher than p. Ka, the conjugate base predominates in solution. © 2011 Pearson Education, Inc.

Chapter Summary, Continued 8. 6 p. Ka, Continued If p. H is lower than

Chapter Summary, Continued 8. 7 Amino Acids: Common Biological Weak Acids • Amino acids contain an acid and base functional group. • Amino acids are building blocks of proteins. • Since amino acids contain both carboxylate and protonated amine functional groups at physiological p. H, they have a net charge of zero and are called a zwitterion. • The isoelectric point is the p. H where the zwitterion exists. © 2011 Pearson Education, Inc.

Chapter Summary, Continued 8. 8 Buffers: An Important Property of Weak Acids and Bases • Buffers resist changes in p. H when an acid or a base is added to a solution. • Buffer solutions consist of weak acids and their conjugate base. Strong acids cannot be buffers. • If p. H of blood drops below the normal range of 7. 35– 7. 45 a condition called acidosis occurs. © 2011 Pearson Education, Inc.

Chapter Summary, Continued 8. 8 Buffers: An Important Property of Weak Acids and Bases, Continued • If blood p. H becomes elevated, a condition called alkalosis exists. • Acidosis can be caused by changes in breathing or by changes in metabolism, causing acid to build up in the bloodstream. Alkalosis is caused by acid being removed from the bloodstream. © 2011 Pearson Education, Inc.