Acids and Bases. Chapter 11 - PDF Free Download

Acids and Bases Chapter 11

Acids and Bases in our Lives Acids and bases are important substance in health, industry, and the environment. One of the most common characteristics of acids is their sour taste. Lemons and grapefruits taste sour because they contain acids such as citric and ascorbic acid (vitamin C). Vinegar tastes sour because it contains acetic acid.

Acids and Bases in our Lives We produce lactic acid in our muscles when we exercise. Acid from bacteria turns milks sour in the products of yogurt and cottage cheese. We have hydrochloric acid in our stomachs to help digest food and we take antacids, which are bases such as sodium bicarbonate, to neutralize the effects of too much stomach acid.

Acids and Bases in our Lives In the environment, the acidity or ph of rain, water, and soil can have significant effects. When rain becomes too acidic, it can dissolve marble statues and accelerate the corrosion of metals. In lakes and ponds, the acidity of water can affect the ability of plants and fish to survive. The acidity of soil around plants affect their growth. It can stop the plant from taking up nutrients through the roots

Acids and Bases in our Lives The lungs and kidneys are the primary organs that regulate the ph of body fluids, including blood and urine. Major changes in the ph of the body fluids can severely affect biological activities within the cells. Buffers are present to prevent large fluctuations.

Chapter 11 Acids and Bases 11.1 Acids and Bases 11.2 Brønsted-Lowry Acids and Bases 11.3 Strengths of Acids and Bases 11.4 Dissociation Constants for Acids and Bases 11.5 Dissociation of Water 11.6 The ph Scale 11.7 Reactions of Acids and Bases 11.8 Acid-Base Titration 11.9 Buffers

11.1 - Acids and Bases Describe and name acids and bases.

Acids The term acid comes from the Latin word acidus which means sour. In 1887, the Swedish chemistry Svante Arrhenius was the first to describe acids as substances that produce hydrogen ions (H + ) when they dissolve in water.

Acids are Electrolytes Because acids produce ions in water, they are also electrolytes (can conduct electricity). Hydrogen chloride dissociates in water to give hydrogen ions, H +, and chloride ions, Cl - : It is the hydrogen ions that give acids a sour taste.

Naming Acids Acids have two common formats: Binary acids: H n X H n = some number of H s x=nonmetals Examples: HCl, HBr, H, H 2 S Polyatomic acids: H n XO m XO m = polyatomic ion Examples: H 2 SO 4, H 3 PO 4, HClO 4

Naming Acids Binary acids: H n X hydro[nonmetal ic] acid Change the ending of the nonmetal to ic and insert into the brackets. hydro and acid do not change. HCl HBr H 2 S

Polyatomic Ion Review More O s = -ate SO 4 2- Less O s = -ite SO 3 2- Chlorine can form 4 polyatomic ions with oxygen: ClO 4 - ClO 3 - ClO 2 - ClO -

Naming Acids Polyatomic Acids: H n XO m [Polyatomic ion] acid -ate to ic -ite to ous H 2 SO 4 H 3 PO 4 HClO 3

Bases You may be familiar with some household bases such as antacids, drain cleaners, and oven cleaners. According to the Arrhenius theory, bases are ionic compounds that dissociate into cations and hydrogen ions (OH-) when they dissolve in water. They are electrolytes.

Bases Most Arrhenius bases are formed from a metal from Groups 1 or 2 and one or more hydroxides (OH - ) M(OH) n M=metal (OH) n = 1 or more hydroxide group Examples: LiOH, Ca(OH) 2 The hydroxide ions give bases common characteristics such as a bitter taste or slippery feel.

Naming Bases Bases have the same names that we used for ionic compounds. LiOH NaOH Ca(OH) 2 Al(OH) 3

11.2 Brønsted-Lowry Acids and Bases Identify the conjugate acid-base pairs for Brønsted- Lowry acids and bases.

Arrhenius Acids and Bases The definitions we gave in section 11.1 for acids and bases were first described by Arrhenius. So we call acids and bases described by H + and OH - as Arrhenius acids and bases. Arrhenius acid: substances that produce H + in water. Arrhenius base: substances that produce OH - in water.

Brønsted-Lowry Acids and Bases In 1923, a pair of scientists, J.N. Brønsted and T.M. Lowry expanded the definitions of acids and bases. The shortcoming of the Arrhenius definitions was that there were many molecules that didn t have OH- groups that acted like bases. A new set of definitions describing Brønsted-Lowry acids and bases included a greater number of molecules.

Brønsted-Lowry Acids and Bases Brønsted-Lowry acid: a substance that donates a hydrogen ion, H + Brønsted-Lowry base: a substance that accepts a hydrogen ion, H + Arrhenius acid: produces H + Arrhenius base: produces OH -

H + = H 3 O + A free hydrogen, H+, does not actually exist in water. Its attraction to polar water molecules is so strong that the H+ bonds to a water molecules and forms a hydronium ion, H 3 O +

Brønsted-Lowry acid: donates H + Brønsted-Lowry base: accepts H + Brønsted-Lowry Acids HCl donates its H + to water producing H 3 O + and Cl - By donating the H +, HCl is acting as the acid in this reaction. By accepting the H +, water is acting as a base in this reaction.

Brønsted-Lowry acid: donates H + Brønsted-Lowry base: accepts H + Brønsted-Lowry Bases Water gives an H + to NH 3 forming NH 4+ and OH - NH 3 acts as the base by accepting the H + Water acts as the acid by donating the H+

Water: a B-L acid and base Water can act as both a Bronsted-Lowry acid or base depending on what it reacts with.

Brønsted-Lowry acid: donates H + Brønsted-Lowry base: accepts H + Practice Identify the reactant that is a Bronsted-Lowry acid and the reactant that is a Bronsted-Lowry base: HBr(aq) + H 2 O(l ) H 3 O + (aq) + Br - (aq)

Brønsted-Lowry acid: donates H + Brønsted-Lowry base: accepts H + Practice Identify the reactant that is a Bronsted-Lowry acid and the reactant that is a Bronsted-Lowry base: CN - (aq) + H 2 O(l ) HCN(aq) + OH - (aq)

Conjugate Acid-Base Pairs According to Bronsted-Lowry theory, a conjugate acid-base pair consists of molecules or ions related by the loss of one H + by an acid, and the gain of one H + by a base. Every acid-base reaction contains two conjugate acid-base pairs because an H + is transferred in both the forward and reverse directions.

Conjugate Acid-Base Pairs When an acid such as HF loses one H +, it becomes F -. HF is the acid, and F - is its conjugate base. * The conjugate is always what is formed by donating or accepting H +. So it is always on the products side.

Conjugate Acid-Base Pairs When the base H 2 O gains an H +, its conjugate acid, H 3 O + is formed.

Conjugate Acid-Base Pairs Now if we combine the two previous examples:

Conjugate Acid-Base Pairs

Amphoteric Substances Water can act like an acid when it donates H + or as a base when it receives H + Substances that can act as both acids and bases are amphoteric. Water is the most common amphoteric substance and its behavior depends on the other reactant. Water will donate H + when mixed with a base and will accept H + when mixed with an acid.

Amphoteric Substances Another example of an amphoteric substance is bicarbonate, HCO 3-. With a base, HCO 3- acts as an acid and donates H + to give CO 3-. With an acid, HCO 3- acts as a base and accepts H + to give H 2 CO 3

Practice Identify the conjugate acid-base pairs in the following reaction: HBr(aq) + NH 3 (aq) Br - (aq) + NH 4+ (aq)

11.3 Strengths of Acids and Bases Write equations for the dissociation of strong and weak acids; identify the direction of reaction.

Strong vs Weak In the process called dissociation, an acid or base separates into ions in water. The strength of an acid is determined by the moles of H 3 O + that are produced for each mole of acid that dissolves. The strength of a base is determined by the moles of OH - that are produced for each mole of base that dissolves. Strong acids and bases dissociate completely in water. Weak acids and bases dissociate only slightly, leaving most of the initial acid or base undissociated.

Strong Acids Strong acids are examples of strong electrolytes because they donate H + so easily that their dissociate in water is essentially complete. When HCl (a strong acid) dissociates in water, H + is transferred to H 2 O. The resulting solution contains essentially only H 3 O + and Cl -. Thus one mole of a strong acid dissociates in water to yield one mole of H3O+ and one mole of its conjugate base. We write the equation for a strong acid, such as HCl, with a single arrow.

Weak Acids Weak acids are weak electrolytes because they dissociate slightly in water, forming only a small amount of H 3 O + ions. When acetic acid dissociates in water, it donates the H+ to water. However, only part of the acetic acid molecules dissociate into ions. Most remain as molecules. Thus one mole of a weak acid partially dissociates in water to give less than a mole of H 3 O + and C 2 H 3 O2 - We write the equation for a weak acid in aqueous solutions with a double arrow to indicate that the forward and reverse reactions are at equilibrium.

Strong and Weak Acids There are only 6 common strong acids: Hydroiodic acid HI Heavily regulated Hydrobromic acid HBr Used to make other molecules and extracting ore Perchloric acid HClO 4 Rocket fuel ingredient Hydrochloric acid HCl Stomach acid Sulfuric acid H 2 SO 4 Drain cleaner, lead-acid batteries Nitric acid HNO 3 Explosives ingredient

Strong and Weak Acids The rest are weak acids.

Diprotic Acids Some weak acids, such as carbonic acid, are diprotic acids that have two H +, that dissociate one at a time. For example, carbonated soft drinks are prepared by dissolving CO 2 in water to form carbonic acid, H 2 CO 3. H 2 CO 3 dissociates partially into HCO 3- and H + in water: H 2 CO 3 (aq) + H 2 O(l) H 3 O + (aq) + HCO 3- (aq) HCO 3- is also a weak acid and will partially dissociate into CO 2-3 and H + HCO 3- (aq) + H 2 O(l) H 3 O + (aq) + CO 2-3 (aq)

Diprotic Acids Sulfuric acid, H 2 SO 4, (a strong acid) is also a diprotic acid. H 2 OS 4 will dissociate completely into H + and HSO 4- : H 2 SO 4 (aq) + H 2 O(l) H 3 O + (aq) + HSO 4- (aq) HSO 4- is a weak acid and dissociates only partially: HSO 4- (aq) + H 2 O(l) H 3 O + (aq) + SO 4 2- (aq)

Acid Summary A strong acid in water dissociates completely into ions. A weak acid in water dissociates only slightly into a few ions but remains mostly as molecules. Strong acid: HI(aq) + H 2 O(l) Weak acid: HF(aq) + H 2 O(l) H 3 O+(aq) + I - (aq) H 3 O+(aq) + F - (aq)

Bases As strong electrolytes, strong bases dissociate completely in water. KOH(s) K+(aq) + OH-(aq) Weak bases are weak electrolytes that are poor H+ acceptors and produce very few ions in solution. NH 3 (g) + H 2 O(l) NH 4+ (aq) + OH - (aq)

Bases in household products

Direction of Reaction There is a relationship between the components of each conjugate acid-base pair: Strong acids have weak conjugate bases. As the strength of the acid decreases, the strengths of the 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. H 3 SO 4 (aq) + H 2 O(l) H 3 O + (aq) + HSO 4- (aq)

Practice By comparing their relative strengths, we can determine the direction of a reaction. H 2 SO 4 (aq) + H 2 O(l) H 3 O + (aq) + HSO 4- (aq)

Practice Which direction will the reaction favor? CO 3 2- (aq) + H 2 O(l) HCO 3- (aq) + OH - (aq)

Practice Which direction will the reaction favor? HF(aq) + H 2 O(l) H 3 O + (aq) + F - (aq)

11.4 Dissociation Constants for Acids and Bases Write the expression for the dissociation constant of a weak acid or weak base.

As we have seen, acids have different strengths depending on how much they dissociate in water. Because the dissociation of strong acids in water is essentially complete, the reaction is not considered to be an equilibrium situation. However, because weak acids in water dissociate only slightly, the ion products reach equilibrium with the undissociated weak acid molecules.

Formic acid HCHO 2, the acid found in bee and ant stings, is a weak acid. It dissociates in water to form hydronium ion, H 3 O +, and formate ions CHO 2 -

Writing Dissociation Constant Expressions Because weak acids and bases reach an equilibrium when mixed in water, we can write an equilibrium constant expression (just like in ch. 10). aa + bb cc + dd K a = [Products] [Reactants] = [D]d [C] c [A] a [B] b K a is called the acid dissociation constant.

Practice Write the equilibrium expression. HCHO 2 (aq) + H 2 O(l) H 3 O + (aq) + CHO 2- (aq) * Only (aq) states are included in equilibrium expressions. (s) and (l) are ignored (including water).

Writing Dissociation Constants An equilibrium expression can also be written for weak bases: CH 3 -N 2 (aq) + H 2 O(l) CH 3 -NH 3+ (aq) + OH - (aq) Kb = [Products] [Reactants] = * Only (aq) states are included in equilibrium expressions. (s) and (l) are ignored (including water).

Dissociation Constants Just like in chapter 10, K s less than 1 indicate that there is more reactant than product. Which is in agreement of how we defined weak acids and weak bases. (Mostly molecules (reactants) and a small amount of ions (products)). Strong acids and bases have very large K s because its almost 100% dissociated. These K s are not usually bothered with.

11.5 Dissociation of Water Use the water dissociation constant expressions to calculate the [H 3 O + ] and [OH - ] in an aqueous solution.

In this section, we will use the dissociation constant expression and apply it to a very important equilibrium reaction: water reacting with itself.

In many acid-base reactions, water is amphoteric, which means tat it can act either as an acid or a base. In pure water, there is a forward reaction between two water molecules that transfers H + from one water molecule to the other. One molecule acts as an acid by losing H+ and the other water molecule that gains H + acts as the base. Water Every time H+ is transferred between 2 water molecules, the products are one H3O+ and one OH-, which reacts in the reverse direction to re-form two water molecules.

Water Dissociation Constant, K w H 2 O(l) + H 2 O(l) H 3 O + (aq) + OH - (aq) Kw = Experiments show that in pure water and 25 C, [H3O+] = [OH-] = *ignore (s) and (l) If we plug the concentrations back into Kw: Kw =

Neutral, Acidic, and Basic Solutions The K w applies to any aqueous solution at 25 C because all aqueous solutions contain H 3 O + and OH -. When [H 3 O + ] and [OH - ] in a solution are equal, the solution is neutral. However most solutions are not neutral; they have different concentrations of [H 3 O + ] and [OH - ].

Neutral, Acidic, and Basic Solutions If acid is added to water, there is an increase in [H 3 O + ] and a decrease in [OH - ], which makes it an acidic solution. If base is added to water, [OH - ] increases and [H 3 O + ] decreases, which gives a basic solution. However for any aqueous solution, whether it is neutral, acidic, or basic, [H 3 O + ][OH - ] = 1.0 x 10-14

Using K w to calculate [H 3 O + ] and [OH - ] If we know [H 3 O + ], we can use K w to calculate [OH - ] or if we know [OH - ] we can use K w to calculate [H 3 O + ]. Kw = [H 3 O + ][OH - ] [OH - ] = K w [H3O + ] [H 3 O + ] = K w [OH ]

Practice A vinegar solution has a [OH - ] = 5.0 x 10-12 M at 25 C. What is [H 3 O + ] of the vinegar solution? Is the solution acidic, basic, or neutral?

Practice What is the [H 3 O + ] of an ammonia cleaning solution with [OH - ] = 4.0 x 10-4 M? Is the solution acidic, basic, or neutral?

11.6 - The ph Scale Calculate ph from [H 3 O + ]; given the ph, calculate the [H 3 O + ] and [OH - ] of a solution.

ph Scale Although we have expressed H 3 O + and OH - as molar concentrations, it is more convenient to describe the acidity of solutions using the ph scale. On this scale, a number between 0 to 14 represents the [H 3 O + ] concentration for common solutions Acidic solution ph less than 7.0 Neutral solution ph = 7.0 Basic solution ph greater than 7.0

ph Scale When an acid is added to water, the [H 3 O + ] (acidity) of the solution increases, but the ph decreases. When a base is added to pure water, it becomes more basic. Which means the acidity decreases and the ph increases.

Calculating the ph of Solutions The ph scale is a logarithmic scale that corresponds to the [H 3 O + ] of aqueous solutions. ph = -log[h 3 O + ]

Calculating the ph of Solutions ph = -log[h 3 O + ] Because ph is a logarithmic scale, a change of 1.0 ph unit corresponds to a 10x in [H 3 O + ].

Practice If a solution of aspirin (acetylsalicylic acid) has a [H 3 O + ] = 1.7 x 10-3 M, what is the ph of the solution? ph = -log[h 3 O + ]

Practice What is the ph of bleach with [H 3 O + ] = 4.2 x 10-12 M? ph = -log[h 3 O + ]

Practice ph can still be calculated if we are given [OH - ] instead of [H 3 O + ]. What is the ph of an ammonia solution with [OH - ] = 3.7 x 10-3 M ph = -log[h 3 O + ] K w = [H 3 O + ][OH - ] = 1.0 x 10-14

Practice Calculate the ph of a sample of bile that has [OH - ] = 1.3 x 10-6 M K w = [H 3 O + ][OH - ] = 1.0 x 10-14 ph = -log[h 3 O + ]

Calculating [H 3 O + ] from ph If we have ph, we can calculate [H 3 O + ]: ph = -log[h 3 O + ] [H 3 O + ] = 10 -ph

Practice If the ph of a solution is 3.0, what is [H 3 O + ]? [H 3 O + ] = 10 -ph

Practice Calculate [H 3 O + ] for a urine sample, which has a ph of 7.5. [H 3 O + ] = 10 -ph

Practice What are the [H 3 O + ] and [OH - ] of Diet Coke that has a ph of 3.17? [H 3 O + ] = 10 -ph K w = [H 3 O + ][OH - ] = 1.0 x 10-14

11.7 Reactions of Acids and Bases Write balanced equations for reactions of acids with metals, carbonate or bicarbonates, and bases.

Salt Salt: an ionic compound that does not have H + as the cation or OH - as the anion. Salts KCl NaCl CaCl 2 Not Salts NaOH HCl H 2 S FeS Ca(OH) 2

Acids react with Metals Acids react with certain metals to produce hydrogen gas (H 2 ) and a salt. Active metals include: K, Na, Ca, Mg, Al, Zn, Fe, and Sn. In these single replacement reactions, the metal ion replaces the hydrogen in the acid. Mg(s) + 2HCl(aq) H 2 (g) + MgCl 2 (aq) Zn(s) + 2HNO 3 (aq) H 2 (g) + Zn(NO 3 ) 2 (aq)

Acids react with Carbonates and Bicarbonates When an acid is added to a carbonate (CO 3 2- ) or bicarbonate (HCO 3- ), the products are carbon dioxide gas, water, and a salt. 2HCl(aq) + Na 2 CO 3 (aq) CO 2 (g) + H 2 O(l) + 2NaCl(aq) HBr(aq) + NaHCO 3 (aq) CO 2 (g) + H 2 O(l) + NaBr(aq)

Acids react with Carbonates and Bicarbonates 2HCl(aq) + Na 2 CO 3 (aq) CO 2 (g) + H 2 O(l) + 2NaCl(aq) The acid reacts with CO 3 2- or HCO 3- to produce carbonic acid, H 2 CO 3, which breaks down into CO 2 and H 2 O.

Acids and Hydroxides: Neutralization Neutralization: is a reaction between a strong or weak acid with a strong base combine to form water. HCl(aq) + NaOH(aq) H 2 O(l) + NaCl(aq) The H + from the acid and OH - from the base form H 2 O. The salt is the base s cation and acid s anion.

Balancing Neutralization Equations In a neutralization reaction, one H + always reacts with one OH -. Therefore, a neutralization may need coefficients to balance the H + from the acid with the OH - from the base. 2HCl(a) + Mg(OH) 2 (aq) 2H 2 O(l) + MgCl 2 (aq)

Practice Write the balanced equation for the neutralization of HCl(aq) and Ba(OH) 2 (s).

11.8 - Buffers Describe the role of buffers in maintaining the ph of a solution; calculate the ph of a buffer.

Buffers The ph of water and most solutions changes drastically when a small amount of acid or base is added. However, when an acid or base is added to a buffer solution, there is little change in ph. A buffer solution maintains the ph of a solution by neutralizing small amounts of acids and base.

Buffers in the Blood In the human body, blood contains plasma, white blood cells and platelets, and red blood cells. The plasma contains buffers that maintain a consistent ph of about 7.4.

Buffers in the Blood If the ph of the blood plasma goes slightly above or below 7.4, changes in our oxygen levels and our metabolic processes can be drastic enough to cause death. Even though we obtain acids and bases from foods and cellular reactions, the buggers in the body absorb those compounds so effectively that the ph of our blood plasma remains essentially unchanged.

Buffers In a buffer, an acid is present to react with any OH - that is added, and a base is present to react with any H + (H 3 O + ) that is added. However, the acid and base must not neutralize each other. Therefore a combination of an acid-base conjugate pair (HA/A - ) is used in a buffer. Most buffer solutions consist of nearly equal concentrations of a weak acid and its conjugate base. Or a weak base and its conjugate acid Common buffers: HC 2 H 3 O 2 /C 2 H 3 O 2 - H 2 PO 4- /HPO 4 2- HPO 4 2- /PO 4 3- HCO 3- /CO 3 2- NH 4+ /NH 3

Preparing a Buffer A typical buffer can be made from a weak acid, such as acetic acid (HC 2 H 3 O 2 ) and its salt, sodium acetate (NaC 2 H 3 O 2, written C 2 H 3 O 2- ) As a weak acid, acetic acid dissociates slightly in water to form H 3 O + and a very small amount of C 2 H 3 O 2-. HC 2 H 3 O 2 (aq) + H 2 O(l) H 3 O + (aq) + C 2 H 3 O 2- (aq) For the buffer to work, more C 2 H 3 O 2 - is needed so NaC 2 H 3 O 2 is also added to the solution. NaC 2 H 3 O 2 Na + + C 2 H 3 O 2 -

Using a Buffer How the buffer maintains the [H 3 O + ] (to balance the ph) When a small amount of acid is added, the additional H 3 O + combines with the acetate ion, C 2 H 3 O 2- : HC 2 H 3 O 2 (aq) + H 2 O(l) H 3 O + (aq) + C 2 H 3 O 2- (aq) The new H 3 O + is used up to make more reactant which maintains ph.

Using a Buffer How the buffer maintains the [H 3 O + ] (balances the ph) If a small amount of base (OH - ) is added, it is neutralized by the acetic acid: HC 2 H 3 O 2 + OH - H 2 O + C 2 H 3 O 2 - [H 3 O + ] and thus ph of the solution remains the same.

Calculating the ph of a Buffer HC 2 H 3 O 2 (aq) + H 2 O(l) H 3 O + (aq) + C 2 H 3 O 2- (aq) K a = H 3O + [C 2 H 3 O 2 ] [HC 2 H 3 O 2 ] By solving for [H3O+] we can obtain the ratio of acetic acid/acetate buffer: [H 3 O + ] = K a x [HC 2H 3 O 2 ] [C 2 H 3 O 2 ]

Practice The K a for acetic acid (HC 2 H 3 O 2 ) is 1.8 x 10-5. What is the ph of a buffer prepared with 1.0M HC 2 H 3 O 2 and 1.0M C 2 H 3 O 2-? HC 2 H 3 O 2 (aq) + H 2 O(l) H 3 O + (aq) + C 2 H 3 O 2- (aq)

Practice One of the conjugate acid-base pairs that buffers the blood is H 2 PO 4- /HPO 4 2-, which has a K a of 6.2 x 10-8. What is the ph of a buffer that is prepared from 0.10 M H 2 PO 4- and 0.50 M HPO 4 2-

Buffering Capacity [H 3 O + ] = Ka x [HA] [A ] Because Ka is constant at a given temperature, [H 3 O + ] (and therefore ph) is determined by the [weak acid]/[conj. Base] ratio. As long as the addition of small amounts of either acid or base changes the ratio only slightly, the changes in [H 3 O + ] will be small and the ph will be maintained. If a large amount of acid or base is added, the buffering capacity of the system may be exceeded.