Arrhenius Model. Hydronium Ion. Chapter 13 Acids and Bases

Transcription

1 Chapter 13 Acids and Bases What are Acids and Bases? Strong and Weak Acids and Bases Relative Strengths of Weak Acids Acidic, Basic, and Neutral Solutions The ph Scale Buffered Solutions 131 Copyright The McGrawHill Companies, Inc. Permission required for reproduction or display. Arrhenius Model Arrhenius Model Acids and Bases Proposed by Svante Arrhenius in late 1800s An acid in aqueous solution produces hydrogen ions, H + H HCl(g) 2 O H + (aq) + Cl (aq) A base in aqueous solution produces hydroxide ions, OH H 2 O 2 Hydronium Ion A fundamental problem with the Arrhenius model is the treatment of the behavior of the hydrogen ion, H + Hydrogen ions are better represented as hydronium ions, H 3 O +, in solution NaOH(s) Na + (aq) + OH (aq) Also explains neutralization of acids and bases; when H + reacts with OH, water is formed H + (aq) + OH (aq) H 2 O(l)

2 BronstedLowry Theory An acid is any substance that can donate an H + ion to another substance A base is any substance that can accept an H + ion from another substance Inclusive of all Arrhenius acids and bases Practice BronstedLowry Acids and Bases For each reaction, identify the BronsteadLowry acid and base reactants. 1. OCl (aq) + H 2 O(l) HOCl(aq) + OH (aq) 2. H 2 SO 4 (aq) + F (aq) HSO 4 (aq) + HF(aq) 3. NH 4+ (aq) + H 2 O(l) NH 3 (aq) + H 3 O + (aq) Practice Solutions Bronsted Lowry Acids and Bases For each reaction, identify the BronsteadLowry acid and base reactants. 1. OCl (aq) + H 2 O(l) HOCl(aq) + OH (aq) The water (H 2 O) donates an H + to the OCl to form HOCl. The water is an acid. The OCl accepts a H + from the water (H 2 O), therefore it is the base

3 Practice Solutions Bronsted Lowry Acids and Bases 2. H 2 SO 4 (aq) + F (aq) HSO 4 (aq) + HF(aq) H 2 SO 4, sulfuric acid, is the acid and it donates an H + to the F. F accepts H + from the sulfuric acid and is therefore the base. 3. NH 4+ (aq) + H 2 O(l) NH 3 (aq) + H 3 O + (aq) NH 4+ donates an H + to the water and thus is the acid. H 2 O accepts the H + and is the base. Conjugate AcidBase Pairs When an acid donates an H + to a base, the two products differ from the reactants by one H + ion. Conjugate acid The product that forms as a result of gaining an H + ion Conjugate base The product that forms as a result of losing an H + ion HCl(g) + H 2 O(l) H 3 O + (aq) + Cl (aq) acid base conjugate conjugate acid base Practice Conjugate Acids and Bases Identify the conjugate acid for each base. a)f b)hco 3 c)h 2 O

4 Practice Solutions Conjugate Acids and Bases Identify the conjugate acid for each base. The conjugate acid will have one more H + ion than the base has. a) F The conjugate acid for F is HF. b) HCO 3 The conjugate acid for HCO 3 is H 2 CO 3. c) H 2 O The conjugate acid for H 2 O is H 3 O +. Amphoteric Substances A substance that can act as either an acid or a base Water is the most common amphoteric substance. Another common amphoteric substance is the bicarbonate ion, HCO 3 : HCO 3 (aq) + OH (aq) CO3 2 (aq) + H2 O(l) acid base conjugate conjugate base acid HCO 3 (aq) + H3 O + (aq) H 2 CO 3 (aq) + H 2 O(l) base acid conjugate conjugate acid base Acidic Hydrogen Atoms If an acid has more than one hydrogen atom, we need to determine which hydrogen atoms are acidic. Typically, in oxoacids, any hydrogen atoms bonded to oxygen atoms are acidic. Figure

5 Strong Acids and Bases An acid or a base that is a strong electrolyte and completely ionizes or dissociates in water Strong acid examples: HCl(aq) H 2 SO 4 (aq) HNO 3 (aq) Strong base examples: KOH(aq) Ca(OH) 2 (aq) HCl Example of a Strong Acid An example of a strong acid is HCl: HCl(aq) + H 2 O(l) H 3 O + (aq) + Cl (aq) Table 13.1 Common Strong Acids Formula HCl HBr HI HNO 3 HClO 3 HClO 4 H 2 SO 4 Name hydrochloric acid hydrobromic acid hydroiodic acid nitric acid chloric acid perchloric acid sulfuric acid

6 NaOH Example of a Strong Base Table 13.2 Common Strong Bases Formula LiOH NaOH KOH Mg(OH) 2 Ca(OH) 2 Ba(OH) 2 Name lithium hydroxide sodium hydroxide potassium hydroxide magnesium hydroxide calcium hydroxide barium hydroxide Weak Acids and Bases An acid or base that is a weak electrolyte and therefore, only partially ionizes in water If an acid or base is not strong, then it is weak

7 CH 3 COOH Example of a Weak Acid NH 3 Example of a Weak Base Table 13.3 Some Common Weak Acids Formula CH 3 CO 2 H H 2 CO 3 H 3 C 6 H 5 O 7 HF HOCl HC 3 H 5 O 3 HC 4 H 4 O 5 H 2 C 2 O 4 H 3 PO 4 H 2 C 4 H 4 O 6 Name Acetic acid Carbonic acid Citric acid Hydrofluoric acid Hypochlorous acid Lactic acid Malic acid Oxalic acid Phosphoric acid Tartaric acid Occurrence Vinegar, sour wine Soda, blood Fruit, soda Glass etching Sanitize pool and drinking water Milk Fruit Nuts, cocoa, parsley Soda, blood Candy, wine, grapes

8 Formula NH 3 CaCO 3 Ca(ClO) 2 CH 3 NH 2 (CH 3 ) 3 N Table 13.4 Some Common Weak Bases Name Ammonia Calcium carbonate Calcium hypochorite Methylamine Trimethylamine Occurrence Glass cleaners Antacids, minerals Chlorine source for swimming pools Herring brine Rotting fish Practice Strong vs. Weak Acids and Bases Identify each of the following as a strong acid, weak acid, strong base, or weak base. Write an equation to describe its reaction in water. 1. HI(aq) 2. NaCH 3 CO 2 (aq) 3. NH 4+ (aq) 4. NH 3 (aq) Practice Solutions Strong vs. Weak Acids and Bases 1. HI(aq) Hydroiodic acid is a strong acid (it is on the list of strong acids in Table 13.1). Therefore, it completely dissociates in water. HI(aq) + H 2 O(l) H 3 O + (aq) + I (aq) 2. NaCH 3 CO 2 (aq) Sodium acetate is an ionic compound that partially dissociates in water to form Na + (aq) and CH 3 CO 2 (aq). The CH3 CO 2 ion is the conjugate base of the weak acid, HC 2 H 3 O 2, so it is a weak base and does not completely dissociate in water: CH 3 CO 2 (aq) + H2 O(l) CH 3 CO 2 H(aq) + OH (aq)

9 Practice Solutions Strong vs. Weak Acids and Bases 3. NH 4+ (aq) Ammonium ion is the conjugate acid of the weak base, ammonia (NH 3 ), so ammonium is therefore a weak acid. It does not completely dissociate in water. NH 4+ (aq) + H 2 O(l) NH 3 (aq) + H 3 O + (aq) 4. NH 3 (aq) Ammonia is a weak base, shown in Table 13.4, and does not dissociate completely in water. NH 3 (aq) + H 2 O(l) NH 4+ (aq) + OH (aq) Practice Acids and Bases on the Molecular Level One of the diagrams below represents HClO 4, and the other represents an aqueous solution of HSO 4. Which is which? Explain your reasoning. Practice Solutions Acids and Bases on the Microscopic Level The picture on the left (A) shows no acid molecules and therefore shows an acid that has completely dissociated. Diagram A would then be a strong acid and of the two choices, perchloric acid (HClO 4 ) is the strong one. Diagram B (the picture on the right) shows acid molecules that have not completely dissociated and are therefore the weak acid, HSO

10 Relative Strengths of Weak Acids Acid strength depends on the relative number of acid molecules that ionize when dissolved in water the degree of ionization. Remember that K eq describes the relative amounts of products over reactants. If the value of K eq is larger, then more products exist at equilibrium and the acid has a larger percentage of molecules which have been ionized. Acid Ionization Constants The acid ionization constant, K a, describes the equilibrium that forms when an acid reacts with water. The larger the K a value, the stronger the acid. When using K a values to determine the strengths of conjugate acids and cases, use this rule of thumb: The stronger the acid, the weaker the conjugate base Table 13.5 Weak Acids & K a Values Strongest Acids Acid HF HNO 2 HCO 2 H CH 3 CO 2 H HOCl K a Value 6.3 x x x x x 10 8 Conjugate Base F NO 2 HCO 2 CH 3 CO 2 OCl Weakest Bases Weakest Acids NH 4 + HCN 5.6 x x NH 3 CN Strongest Bases

11 Practice Significance of K a Which solution has the greater concentration of H 3 O +, 0.10 M HOCl [K a = 4.0 x 10 8 ] or 0.10 M HCN [K a = 6.2 x ]? Practice Solutions Significance of K a Which solution has the greater concentration of H 3 O +, 0.10 M HOCl or 0.10 M HCN? Because we have equal concentrations of each acid, we first need to determine which acid is stronger. The stronger acid will dissociate more completely in water, thus producing more H 3 O + ion. Practice Solutions Significance of K a Which solution has the greater concentration of H 3 O +, 0.10 M HOCl or 0.10 M HCN? Looking at the K a s of the two acids from Table 13.5: HOCl K a = 4.0 x 10 8 HCN K a = 6.2 x Since HOCl has the larger K a, it is the stronger acid and will therefore have a greater concentration of H 3 O + ion

12 Polyprotic Acids An acid that contains more than one acidic hydrogen and can thus donate more than one H + ion The acid donates one H + ion at a time in steps The K a values for polyprotic acids are often labeled to indicate the particular step in the overall ionization process (K a1, K a2, K a3, etc.) Figure Table 13.6 K a for Polyprotic Acids Name Carbonic acid Citric acid Hydrosulfuric acid Oxalic acid Phosphoric acid Sulfuric acid Tartaric acid Formula H 2 CO 3 H 3 C 6 H 5 O 7 H 2 S H 2 C 2 O 4 H 3 PO 4 H 2 SO 4 H 2 C 4 H 4 O 6 K a1 4.5 x x x x x 10 3 Strong 1.0 x 10 3 K a2 4.7 x x x x x x x 10 5 K a3 4.0x x10 13 Practice Polyprotic Acids in Water Oxalic acid, H 2 C 2 O 4, occurs in plants and foods such as parsley, rhubarb, almonds, and green beans. Its K a values are listed in Table a) Write equations that show the ionization of oxalic acid in water. b) Besides water, which ion or molecule has the highest concentration in solution when oxalic acid is added to water?

13 Practice Solutions Polyprotic Acids in Water a) Write equations that show the ionization of oxalic acid in water. H 2 C 2 O 4 (aq) + H 2 O(l) H 3 O + (aq) + HC 2 O 4 (aq) K a1 = 5.6 x 10 2 HC 2 O 4 (aq) + H2 O(l) H 3 O + (aq) + C 2 O 2 4 (aq) K a2 = 1.5 x 10 4 b) Besides water, which ion or molecule has the highest concentration in solution when oxalic acid is added to water? The K a1 value for H 2 C 2 O 4 shows that it is a weak acid and ionizes only to a small extent. That means that most of the H 2 C 2 O 4 and is present in water in the highest concentration. Acidic, Basic, and Neutral Solutions Acidic solution The H 3 O + ion concentration is greater than the OH ion concentration. Basic solution The OH ion concentration is greater than the H 3 O + ion concentration. Neutral solution Equal concentrations of OH and H 3 O + Neither acidic nor basic IonProduct Constant of Water Water reacts with itself in a process called selfionization, in which an H + ion is transferred from one water molecule to another: H 2 O(l) + H 2 O(l) OH (aq) + H 3 O + (aq) The equilibrium constant for this process, called the ionproduct constant of water, K w, is: K w = [OH ][H 3 O + ] = 1.0 x (at 25 C) and in pure water, the concentrations would be equal, so: [OH ] = [H 3 O + ] = 1.0 x

14 Table 13.7 Definitions of Neutral, Acidic, and Basic in Aqueous Solution Type of Solution Neutral Acidic Basic [OH ] = [H 3 O + ] [OH ] < [H 3 O + ] [OH ] > [H 3 O + ] [H 3 O + ] =1.0 x >1.0 x <1.0 x [OH ] =1.0 x <1.0 x >1.0 x K w 1.0 x x x Practice Calculating H 3 O + and OH Ion Concentrations Given the concentration of OH ion in each solution, calculate the concentration of H 3 O + in that solution. Identify each solution as acidic, basic, or neutral. a) [OH ] = 1.0 x 10 8 M b) [OH ] = M Practice Solutions Calculating H 3 O + and OH Ion Concentrations Given the concentration of OH ion in each solution, calculate the concentration of H 3 O + in that solution. Identify each solution as acidic, basic, or neutral. a) [OH ] = 1.0 x 10 8 M K w = [H 3 O + ][OH ] [H 3 O + ] = K w = 1.0 x = 1.0 x 10 6 M H 3 O + [OH ] 1.0 x 10 8 Acidic b) [OH ] = M K w = [H 3 O + ][OH ] [H 3 O + ] = K w = 1.0 x = 1.0 x M H 3 O + [OH ] 1.0 x 10 2 Basic Relative Concentration

15 The ph Scale The ph of a solution is the negative logarithm (base 10) of the H 3 O + concentration: ph = log [H 3 O + ] It is convenient to express the acidity of aqueous solutions on a ph scale (shown at right). Figure Practice Solutions Calculating ph What is the ph of each of the following solutions? Once you ve done the calculation, check your answer to make sure it makes sense. a) M HCl First, we need to find the [H 3 O + ]: [H 3 O + ] = M Next, calculate the ph: ph = log [H 3 O + ] = log ( ) = 3.1 Practice Calculating ph What is the ph of each of the following solutions? Once you ve done the calculation, check your answer to make sure it makes sense. b) M NaOH First, we need to find the [H 3 O + ]: K w = [OH ][H 3 O + ] [OH ] = M [H 3 O + ] = K w = 1.0 x = 1.0 x M [OH ] M Next, calculate the ph: ph = log [H 3 O + ] = log (1.0 x ) =

16 Practice Calculating ph What is the ph of each of the following solutions? Once you ve done the calculation, check your answer to make sure it makes sense. c) 1.0 M HNO 3 First, we need to find the [H 3 O + ]: [H 3 O + ] = 1.0 M Next, calculate the ph: ph = log [H 3 O + ] = log (1.0) = 0 Calculating poh poh is defined as the negative logarithm (base 10) of hydroxide ion concentration, [OH ]: poh = log [OH ] The relationship between ph and poh is: ph + poh = 14 Calculating Concentrations from ph and poh The equation to find ph is: ph = log [H 3 O + ] To find the H 3 O + ion concentration, we need to take the inverse log of the negative ph: [H 3 O + ] = 10 ph The equation to find poh is: poh = log [OH ] To find the OH ion concentration, we need to take the inverse log of the negative poh: [OH ] = 10 poh

17 Calculating Concentrations from ph and poh Practice Calculating OH and H 3 O + concentration When the ph of water in a lake falls below about 4.5, the lake may be considered dead because few organisms can survive in such an acidic environment. What is the ph of a lake that has an OH concentration of 1.0 x 10 9 M? Would this lake be considered dead? Practice Solutions Calculating OH and H 3 O + concentration When the ph of water in a lake falls below about 4.5, the lake may be considered dead because few organisms can survive in such an acidic environment. What is the ph of a lake that has an OH concentration of 1.0 x 10 9 M? Would this lake be considered dead? If the ph falls below 4.5, then the lake is dead. [OH ] = 1.0 x 10 9 M K w = [H 3 O + ][OH ] [H 3 O + ] = K w = 1.0 x = 1.0 x 10 5 M [OH ] 1.0 x 10 9 ph = log [H 3 O + ] = log (1.0 x 10 5 ) = 5 Since the ph of the lake is 5, then the lake is NOT dead

18 Measuring ph ph meters and ph indicators are often used to determine the ph of a solution. Figure Buffered Solutions A buffer (also known as a buffer system) is a combination of a weak acid and its conjugate base (or a weak base and its conjugate acid) in about equal concentrations. The main buffer system in the blood is made of H 2 CO 3 /HCO 3 : Figure Practice Buffer Systems Which of the following systems, when added to water, can act as a buffer system? For each buffer system, write a balanced equation. a) HCl and NaOH b) CH 3 CO 2 H and NaCH 3 CO 2 c) HBr and KBr

19 Practice Solutions Buffer Systems Which of the following systems, when added to water, can act as a buffer system? For each buffer system, write a balanced equation. a) HCl and NaOH This is not a buffer system, because HCl is a strong acid. Strong acids cannot be components of buffers, because they ionize completely in water and are not in equilibrium with their conjugate bases. Practice Solutions Buffer Systems Which of the following systems, when added to water, can act as a buffer system? For each buffer system, write a balanced equation. b) CH 3 CO 2 H and NaCH 3 CO 2 This is a buffer system because the acid CH 3 CO 2 H is a weak acid and its conjugate base, CH 3 CO 2, forms when NaCH 3 CO 2 dissolves in water. The equilibrium that forms in water is: CH 3 CO 2 H(aq) + H 2 O(l) CH 3 CO 2 (aq) + H 3 O + (aq) Practice Solutions Buffer Systems Which of the following systems, when added to water, can act as a buffer system? For each buffer system, write a balanced equation. c) HBr and KBr This is not a buffer system, because HBr is a strong acid. Strong acids cannot be components of buffers, because they ionize completely in water and are not in equilibrium with their conjugate bases

20 Math Toolbox 13.1: Log and Inverse Log Functions Using Log Functions on your calculator: ph = log [H 3 O + ] Step 1: Press the +/ (change of sign) key Step 2: Press the log key Step 3: Enter the H 3 O + concentration, and then the ENTER or = key (On some calculators the steps may be reversed) Practice Using Log Functions Use your calculator to find the ph of the following solutions: 1. [H 3 O + ] = 1.0 x 10 8 M 2. [H 3 O + ] = 6.2 x 10 1 M 3. [H 3 O + ] = 5.0 x 10 4 M Practice Solutions Using Log Functions Use your calculator to find the ph of the following solutions: 1. [H 3 O + ] = 1.0 x 10 8 M ph = log [H 3 O + ] = log (1.0 x 10 8 ) = [H 3 O + ] = 6.2 x 10 1 M ph = log [H 3 O + ] = log (6.2 x 10 1 ) = [H 3 O + ] = 5.0 x 10 4 M ph = log [H 3 O + ] = log (5.0 x 10 4 ) =

21 Math Toolbox 13.1: Log and Inverse Log Functions Using Inverse Log Functions: [H 3 O + ] = 10 ph Step 1: Press the INV, SHIFT, or 2 nd button Step 2: Press the log button Step 3: Press the +/ (change of sign) key Step 4: Enter the ph (or poh), and then the ENTER or = key (On some calculators you may need to perform Step 4 first, then step 3, then steps 1 and 2) Practice Using Inverse Log Functions Use your calculator to find the [H 3 O + ] of the following solutions: 1. ph = ph = Practice Solutions Using Inverse Log Functions Use your calculator to find the [H 3 O + ] of the following solutions: 1. ph = 5.00 [H 3 O + ] = 10 ph = = 1.00 x 10 5 M 2. ph = [H 3 O + ] = 10 ph = = x M