ACIDS AND BASES CHAPTER 14

14.1 THE NATURE OF ACIDS AND BASES
1. bases are also alkalis, are characterized by their bitter taste and slippery feel.
2. Arrhenius concept: acids produce hydrogen ions in aqueous solutions, while bases produce hydroxide ions..
3. Hydronium ion: the H3O+ formed when a proton transfers from the HCL molecule to the water molecule.
4. Bronsted-Lowry model: an acid is a proton [H+] donor, and a base is a proton accepter.
5. Acid dissociating constant: Ka, where Ka=[H+][A-]/[HA]. (equilibrium constant, K, is explained in chapter 13.
6. Conjugate base: everything that remains of the acid molecule after a proton is lost. Conjugate acid: formed when a proton is transferred to the base. Conjugate acid-base pair: two substances related to each other by the donating and accepting of a single proton.
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14.2 ACID STRENGTH
1. a strong acid is one acid in which its equilibrium position of its dissociation reaction lies far to the right. al acids dissolve in water. it yields a weak base (weaker than water), which has low affinity for protons. a weak acid is one that its equilibrium lies far to the left. most of HA do not dissociate. a weak acid yields strong conjugate bases. the water molecule is not very good at pulling H+ ions from its conjugate base. the common strong acids include sulfuric acid, hydrochloric acid, nitric acid, and perchloric acid (H2SO4, HCl, HNO3, HClO4, respectively. sulfuric acid, H2SO4, is actually a diprotic acid because two of its acidic protons dissociates.
2. most acids, and common weak acids, are oxyacids, in which the acidic proton is attached to an oxygen atom. organic acids (see DEFINITION) are usually weak (acetic acid). the remainder of the hydrogens in these molecules are not acidic; they do not form H+ in water.
3. water is the most common amphoteric substance. its autoionization process involves the transfer of a proton from one water molecule to another ro produce a hydroxide ion and a hydronium ion. this lead to Kw, which is ion product constant, or the dissociation constant for water ([H3O+][OH-]=1.0x10^-14).
4. autoionization also occurs in ammonia.
14.3 THE pH SCALE
1. pH=-log[H+] pOH=-log[OH-]
2. the number of decimal places in the log (pH value) is equal to the number of significant figures in the original number([H+].
3. the pH changes by 1 for every power of 10 change in [H+]. the pH decreases as [H+] increases.
4. the pH of a solution is usually measured using a pH meter, an electronic device with a probe that can be inserted into a solution of unknown pH. the probe contains an acidic aqueous solution enclosed by a special glass membrane that allows migration of H+ ions. if the unknown solution has a different pH from the solution in the probe, an electric potential results, which is registered on the meter.
5. pKw=pH+pOH=14.00
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14.4 CALCULATING THE pH OF STRONG ACID SOLUTIONS
1. when we deal with acid-base equilibria, we must focus on the solution components and their chemistry. if it is a strong acid, then it completely dissociates with no HCl molecules left. the next step is to focus on the major species, those solution components present in relatively large amount.
2. example question
  1. Example question
3. if the acid concentration is too small, even if it is a strong acid, water prevails, making the pH 7.00
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14.5 CALCULATING THE pH OF WEAK ACID SOLUTIONS
1. the first step is always to write down the major species. for weak acids, they only dissociate to a small extent. usually, it is the weak acid that dominates the dissociation of H+. you first list out the initial concentrations, then determine the changes required to reach equilibrium, which is X because only some of H+ is dissociated. then substitute equilibrium concentrations into equilibrium expressions of Ka (RICE table). this can be solved by using a quadratic equation, or approximate. only approximate when molarity/Ka is greater than 500. at last use the log relationship between [H+] and pH to determine the pH
2. percent dissociation= amount dissociated(mol/L) x 100%/initial concentration (mol/L)
3. for a given weak acid, the percent dissociation increases as the acid becomes more dilute. for solutions of any weak aid HA, [H+] decreases as [HA]o decreases, but the percent dissociation increases as [HA]o decreases.
4. in order to calculate the pH of weak acid mixtures, through Ka, determine the acid that contributes to [H+] the most, and focus on the equilibrium expression of this acid alone. there is only one kind of H+ in these problems, and since the Ka of other acids are so small, the contribution to H+ are negligible.
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14.6 BASES
1. Arrhenius base is a substance that produces OH= ions in aqueous solution. Bronsted-Lowry base is one that accepts protons
2. NaOH and KOH are strong bases that dissociate completely. other strong bases include hydroxides with Group 1A
3. alkaline earth hydroxides are very soluble.
4. calcium hydroxide, often called slaked lime, is widely used in industry because it is inexpensive and plentiful. for example, slaked lime is used in scrubbing stack gasses to remove sulfur dioxide from the exhaust of power plants and factories. in the scrubbing process a suspension of slaked lime is sprayed into the stack gases to react with sulfur dioxide gas. H2O (liquid) and CaSO2 (solid) forms. slaked lime can also be used to soften hard water, which involves the removal of ions such as Ca2+ and Mg2+ that hampers the actions of detergents. . the softening method most often employed in water treatment plants is the Lime-soda process, in which lime, CaO and soda ash, Na2CO3 are added to the water. Co3 2- ion reacts with water to produce the HCO3- ion. when the lime is added to the water, it forms slaked lime, which then reacts with the HCO3- ion from the added soda ash and the Ca+ ion in the hard water to produce calcium carbonate
5. solving for pH of strong bases and weak bases share similar strategies and procedures with solving for pH of strong acids and weak acids, respectively, when doing base problems, be sure to use corresponding Kb to write the equilibrium expression. instead of [H+] being X, [OH-] is X, so in order to find the pH, use 14-pOH, where pOH is -log([OH-])
6. there are many types of proton acceptors that do not contain hydroxide ions, but they increase the concentrations of hydroxide ion because of their reaction with water. for example, ammonia reacts with water to make ammonium and hydroxide ions. the ammonia molecule accepts a proton and thus function as a base. water is the acid in this reaction. other bases that are like ammonia usually have at least one unshared pair of electrons that is capable of forming a bond with a proton. in most of theses bases, the lone pair is located on a nitrogen atom. some examples are methylamine, dimethylamine, trimethylamine, ethylamine, and pyridine.
7. this is the general reaction between a base B and water, where Kb always refers to the reaction of a base with water to form the conjugate acid and the hydroxide ion. the B in the equation compete with OH-, a very strong base, for H+ ion. thus, a Kb value tend to be small when a base is a weak base.
14.7 POLYPROTIC ACIDS
1. include H2SO4, H3PO4, triprotic
2. the successive Ka values for the dissociation equilibria are designated Ka1 and Ka2. for a typical acid, Ka1 is larger than Ka2 is larger than Ka3. it indicates than the loss of a second or a third proton occurs less readily than loss of the first proton. as the negative charge on the acid increases, it becomes more difficult to remove the positively charged proton.
3. phosphoric acid:H3PO4 (Ka1)>H2PO4- (Ka2)>PO4 3- (Ka3)
4. Sulfuric acid
  1. Sulfuric acid, unique case.JPG
14.8 ACID-BASE PROPERTIES OF SALTS
1. salt is another name for ionic compound. when a salt dissolves in water, we assume that it breaks up into its ions, which move independently
2. because the conjugate base of a strong acid has virtually no affinity for protons in water, strong acids completely dissociate in aqueous solutions. so Cl- dissolves in water they don't combine with H+ and have no affect on pH. salt that consist of the cations of strong bases and anions of strong acids have no effect on the [H+] when dissolved in water. aqueous solutions of salts such as KCl, Nacl, KNO3 are neutral with a pH of 7
3. for any salt whose cation has neutral properties, such as Na+ or K+, and whose anion is the conjugate base of a weak acid, the aqueous solution will be basic. the Kb value for the anion can be obtained from the relationship Kb=Kw/Ka, where Kw=1.00x10^-14
4. salts in which the anion is not a base and the cation is the conjugate acid of a weak base produce acidic solutions. salts can also produce an acidic solution when it contains a highly charged metal ion. for example, when solid aluminum chloride dissolves in water, the solution is acidic. the higher the charge on the metal ion, the stronger the acidity of the hydrated ion
5. if the Ka value for the acidic ion is larger than the Kb value for the basic ion, the solution will be acidic. if the Kb value is larger than the Ka value, the solution will be basic. equal Ka and Kb value means a neutral solution.
14.9 THE EFFECT OF STRUCTURE ON ACID-BASE PROPERTIES
1. molecules containing C-H bonds, such as chloroform, and nitromethane, do not produce acidic aqueous solutions because a C-H bond is both strong and nonpolar and thus there is no tendency to donate protons. on the other hand, although the H-Cl bond in gaseous hydrogen chloride is slightly stronger than a C-H bong, it is much more polar, and this molecule readily dissociates when dissolved in water.
2. although one might expect HF to be a very strong acid due to its molarity, it is actually a weak acid. The H-F bond is unusually strong, thus difficult to dissociate.
3. in groupings that have H-O-X bonds (oxyacids), the acid strength increases with an increase in the number of oxygen atoms attached to the central atom. this happens because the very electronegative oxygen atoms are able to draw electrons away from the central atom and the O-H bond. the net effect is to both polarize and weaken the OH bond. a proton is also most readily produced by the molecule with the largest number of attached oxygen atoms. this type of behavior is also observed for hydrated metal ions such as Al3+. the acidity of the water molecules attached to the metal ion is increased by the attraction of electrons to the positive metal ion. the greater the charge on the metal ion, the more acidic the hydrated ion becomes.
4. for acids containing the H-O-X grouping, the greater the ability of X to draw electrons toward itself, the greater the acidity of the molecule. since the electronegativity of X reflects its ability to attract the electrons involved in bonding, it is expected that acid strength also depends on the electronegativity of X.
14.10 ACID BASE PROPERTIES OF OXIDES
1. In the O-X bond, if X has a relatively high electronegativity, the O-X bond would be strong and covalent. when the compound containing H-O-X grouping is dissolved in water, the O-X bond would remain intact. it would be the polar and weak H-O bond that will tend to break, releasing a proton. for example, when a covalent oxide such as sulfur trioxide is dissolved in water, an acidic solution results because sulfuric acid is formed.
2. If X has very low electronegativity, the O-X bond will be ionic and subject to being broken in polar water. For example, NaOH and KOH dissolve in water to give the metal cation and the hydroxide ion.
3. when a covalent oxide dissolves in water, an acidic solution forms. these oxides are called acidic oxides.
4. most ionic oxides, such as those of Group 1A and 2A metals, produce basic solutions when they are dissolved in water. these are called basic oxides.
14.11 THE LEWIS ACID-BASE MODEL
1. Lewis acid is a electron pair acceptor while a lewis base is a electron pair donor
2. the electron deficiency of boron triofluoride makes it very reactive toward any electron pair donor. that is, it is a strong acid.
3. Lewis acids and Lewis bases encompass many acids and bases that do not belong to the Bronsted-Lowry type. an example of this is sulfur trioxide, which is a Lewis Acid, reacts with Water, the Lewis base, to form sulfuric acid.
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14.12 STRATEGY FOR SOLVING ACID-BASE PROBLEMS: A SUMMARY
1. when solving acid-base questions, it might be tempting to memorize a fixed way for a specific kind of problem, however, the best way to solve these types of problems include to list out the major species involved, determine if the reaction goes to completion and the dominating equilibrium. in the end, let the problem guide you and be patient.
MULTIPLE CHOICE EXAMPLES
1. 1. What is the pH of an aqueous solution at 25 degrees Celsius in which [OH-] is 0.0025 M? A) +2.60 B) -2.60 C) +11.40 D)-11.40 E) -2.25
2. 2. What is the pH of an aqueous solution at 25 degrees Celsius that contains 3.98 x 10-9 M hydronium ion? A) 7.000 B) 9.000 C) 8.400 D) 5.600 E) 3.980
3. 3. What is the concentration (in M) of hydronium ions in a solution at 25 degrees Celsius with pH = 4.282? A) 4.28 B) 1.92 ˛ 10-10 C) 1.66 ˛ 104 D) 9.71 E) 5.22 x 10-5
4. 4. Which solution below has the highest concentration of hydroxide ions? A) PH = 3.21 B) pH = 9.82 C) pH = 7.93 D) pH = 12.59 E) pH = 7.00
5. 5. What is the pH of a 0.015-M aqueous solution of barium hydroxide? A) 12.18 B) 1.52 C) 12.48 D) 1.82 E) 10.35
6. ANWERS WITH WORK SHOWN
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FREE RESPONSE EXAMPLES
1. KEY TO IMAGE 1
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2. KEY TO IMAGE2
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3. KEY TO IMAGE3
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DEFINTIONS
1. Arrhenius concept: acids produce hydrogen ions in aqueous solutions, while bases produce hydroxide ions.
2. Bronsted-Lowry model: an acid is a proton [H+] donor, and a base is a proton accepter.
3. Hydronium ion: the H3O+ formed when a proton transfers from the HCL molecule to the water molecule.
4. Conjugate base: everything that remains of the acid molecule after a proton is lost
5. Conjugate acid: formed when a proton is transferred to the base
6. Conjugate acid-base pair: two substances related to each other by the donating and accepting of a single proton.
7. Acid dissociating constant: Ka, where Ka=[H+][A-]/[HA]. (equilibrium constant, K, is explained in chapter 13.
8. Strong acid: equilibrium lies far to the right. Ka is large.
9. weak acid: equilibrium lies far to the left. Ka is small.
10. Diprotic acid: an acid having two acidic protons
11. Oxyacids: The acidic proton is attached to an oxygen atom
12. Organic acids: acids with a carbon atom backbone
13. Carboxyl group:
14. Monoprotic acid: acids with only one acidic poton
15. Amphoteric substances: a substance that can behave either as an acid or as a base. For example, water.
16. Autoionization: a process that involves the transfer of a proton from one water molecule to another to produce a hydroxide ion and a hydronium ion
17. Ion product constant: Kw, where Kw=[H+][OH-]. it is also called the dissociation constant for water, and it refers to the autoionization of water.
18. pH scale: a compact way to represent solution acidity. the pH is a log scale based on 10, where pH=-log[H+]
19. Major species: the solution components present in relative large amounts. the focus when dealing with aqueous solutions.
20. Percent dissociation: (amount dissociated mol/L)(100%)/(initial concentration mol/L)
21. Strong bases: Equilibrium lies far to the right. Kb is large
22. Slaked lime: calcium hydroxide (Ca(OH)2), inexpensive and plentiful. it is widely used as scrubbing stack gases to remove sulfur dioxide from the exhaust of power plants and factories
23. Lime-soda process: lime, CaO, and soda ash, Na2CO3, are added to soften the water. CaO and H2O yields Ca(OH)2, which then reacts with the HCO3- ion from the added soda ash and the Ca2+ ion in the hard water to produce calcium carbonate. for evey mole of Ca(OH)2 consumed, 1 mole of Ca2+ is removed from the hard water, thereby softening it.
24. Weak bases: bases with small Kb values, and its equilibrium lies far to the left
25. Amine:
26. Polyprotic acid: acids, such as sulfuric acid and phosphoric acid, that can furnish more than one proton.
27. Triprotic acid: acids that can dissociate 3 protons. for example, phosphoric acid.
28. Salt: another name for ionic compounds. it can break into independently moving ions, which can behave as acids or bases. salts that consist of cations of strong bases and anions of strong acids have no effect on the pH when dissolved in water. for any salt whose cation has neutral properties, such as Na+, or K+, and whose anion is the conjugate base of a weak acid, the aqueous solution will be basic. salts in which the anion is not a base and the cation is the conjugate acid of a weak base produce acidic solutions.
29. Acidic oxides: the oxide of when a covalent oxide dissolves in water and an acidic solution forms
30. Basic oxides: the ionic oxides, such as the metals in group 1A and 2A,that produce basic solutions when they dissolve in water
31. Lewis acid: electron pair acceptor
32. Lewis base: electron pair donor
LEARNING OBJECTIVES
1. Learning Objective 2.1: students can predict properties of substances based on their chemical formulas and provide explanations of their properties based on particle views.
2. Learning objective 2,2: the student is able to explain the relative strengths of acids and bases based on molecular structure, interparticle forces, and solution equilibrium.
3. Learning objective 3.7: the student is able to identify compounds as Bronsted-Lowry acids, bases and/or conjugate acid-base pairs using proton transfer reactions to justify the identification.
4. Learning objective 6.1: the student is able to, given a set of experimental observations regarding physical, chemical, biological, or environmental processes that are reversible, construct an explanation that connects the observations to the reversibility of the underlying chemical reactions or processes.
5. Learning objective 6.11: the student can generate or use a particulate representations of an acid (strong or weak or polyprotic) and a strong base to explain the species that will have large versus small concentrations at equilibrium.
6. Learning objective 6.12: the student can reason about the distinction between strong and weak acid solutions with similar values of pH, including the percent ionization of the acid, the concentrations needed to achieve the same pH, and the amount of base needed to reach the equivalence point in a titration
7. Learning objective 6.14: the student can, based on the dependence of Kw on temperature, reason that neutrality requires [H+]=[OH-] as opposed to requiring pH=7, including especially the applications to biological systems
8. Learning objective 6.15: the student can identify a given solution as containing a mixture of strong acids and/or bases and calculate or estimate the pH(and concentrations of all chemical species) in the resulting solution.
9. Learning objective 6.16: the student can identify a given solution as being the solution of a monoprotic weak acid or base (including salts in which one ion is a weak acid or base), calculate the pH and concentration of all species in the solution, and/or infer the relative strengths of the weak acids or bases from given equilibrium concentrations.