3 H 2 O H Nh 3 example essay topic

CHEMISTRY 122 CHAPTER 17 NOTES CHAPTER 17 EXERCISES (THESE ARE NOT TO BE TURNED IN, BUT QUIZ QUESTIONS WILL BE TAKEN FROM THESE EXERCISES SO IT IS TO YOUR BENEFIT TO DO THEM): 2, 3, 5, 6, 8, 10, 13, 15, 16, 18, 19, 21, 23, 25, 27, 28, 29, 31, 33, 37, 39, 41, 43, 44, 47, 48, 50, 53, 57, 60, 63, 64, 67, 70, 71, 73, 75, 78, 80, 87, 95,102,112 Chemistry of Acids and BasesSvante Arrhenius (1884) developed a theory of acid-base reactions Arrhenius acid = substance that contains H and produces H+ ions in aqueous solutions Arrhenius base = substance that contains OH and produces OH- ions in aqueous solutions Neutralization = H+ (aq) + OH- (aq) -- "^3 H 2 O (l) Hydrated hydrogen ion H+ does not exist independently in aqueous solutions. It is associated with one or more water molecules. Therefore H+ is said to be hydrated and exists as H+ (H 2 O) n where n is a small integer. The H+ ions are attracted to the f^O- part of the water molecule centered on the O atom. The exact value of n is not known for most solutions but one way to represent it is as H 3 O+ which is called the ion.

For our purposes H+ (aq) and H 3 O+ are the same. Bronsted-Lowry theory expanded on the Arrhenius theory of acids and bases. Bronsted-Lowry acid = a proton (H+) donor Bronsted-Lowry base = a proton (H+) acceptor This means that an acid-base reaction is due to the transfer of H+ from an acid to a base. Example: HCl (aq) + NaOH (aq) -- "^3 Na Cl (aq) + H 2 O (l) Arrhenius description: Bronsted-Lowry description: Bronsted-Lowry acid -base reactions can also be described in terms of conjugate acid-base pairs. The substance left when an acid loses H+ is known as its conjugate base. Acid Conjugate Base HCl Cl- HNO 3 NO 3- CH 3 COOH CH 3 COO- H 2 SO 4 HSO 4- HSO 4- SO 42- H 3 O+ H 2 O H 2 O OH- The substance produced when a base gains H+ is known as its conjugate acid.

Base Conjugate acid OH- H 2 O H 2 O H 3 O+ CO 32- HCO 3- HCO 3- H 2 CO 3 NH 3 NH 4+ Reactions can be described in terms of what the conjugate acid-base pairs are. HCl (aq) + NaOH (aq) -- "^3 Na Cl (aq) + H 2 O (l) The stronger the acid the weaker its conjugate base and the stronger the base the weaker its conjugate acid. Water can act like an acid in the presence of a base or like a base in the presence of an acid. NH 3 (aq) + H 2 O (l) "^2"^3 NH 4+ (aq) + OH- (aq) CH 3 COOH (aq) + H 2 O (l) "^2"^3 H 3 O+ (aq) + CH 3 COO- (aq) Autoionization of water Autoionization = reaction in which water slightly ionizes to produce hydrated hydrogen and hydroxide ions H 2 O + H 2 O "^2"^3 H 3 O+ (aq) + OH- (aq) Or H 2 O "^2"^3 H+ (aq) + OH- (aq) One water molecule donates a H+ ion to the other water molecule Water is said to be because it can both accept and donate protons. Amphoterism is a property of a substance to act as either an acid or a base. Amphiprotism is a more specific type of in which transfer of H+ is involved.

Some of the insoluble metal hydroxides are amphoteric because they dissolve in both acids and bases. Example: Al (OH) 3 Al (OH) 3 + 3 HCl -- "^3 Al Cl 3 + 3 H 2 O Al (OH) 3 + NaOH -- "^3 Na Al (OH) 4 (aq) Autoionization of Water Pure water ionizes to a very slight extent: H 2 O (l) + H 2 O (l) "^2 -- "^3 H 3 O+ (aq) + OH- (aq) Kw = ion product for water = [H 3 O+] [OH-] = 1.0 x 10-14 at 250 For pure water [H 3 O+] = [OH-] = 1.0 x 10-7 The Kw expression [H 3 O+] [OH-] = 1.0 x 10-14 at 250 C holds for dilute aqueous solutions as well as for pure water The value of Kw does vary with temperature but unless otherwise stated assume that the temperature is 250 CExample: Calculate the concentrations of OH- and H 3 O+ in a solution of 0.00050 M Ca (OH) 2 (aq) Acidic solution Basic solution Neutral solution pH and pH scales pH = -log [H 3 O+] or [H 3 O+] = 10-pH p OH = -log [OH-] or [OH-] = 10-p OH p Kw = -log Kw = pH + p OH = 14.00 at 250 CExample: Determine the [H 3 O+], [OH-], pH, and p OH for an aqueous solution of 0.00035 M HCl Example: Determine the [H 3 O+] of a solution that has a pH of 5.26 at 250 C. At 250 Acidic solution Basic solution Neutral solution pH of solutions can be estimated by using acid-base indicators which are substances that change color depending upon the pH of the solution. Usually the indicator is a weak acid which is one color and its conjugate base which is a different color. The pH at which the color change occurs depends upon the equilibrium constant for the reaction. pH meters are used to obtain more precise values of pH. The meter is a voltmeter connected to a pH electrode which has a membrane sensitive to the movement of H 3 O+ ions and develops a potential that depends upon the concentration of H 3 O+ ions. Weak acids and bases are partly dissociated or ionized in solution Consider acetic acid CH 3 COOH (aq) + H 2 O (l) "^2 -- "^3 CH 3 COO- (aq) + H 3 O+ (aq) The molecular acid is in equilibrium with its conjugate base and H 3 O+The equilibrium constant for this reaction is Ka = acid ionization constant The larger Ka the greater the extent of ionization (the stronger the acid) The smaller the Ka the less the extent of ionization (the weaker the acid) Table 17.3 lists Ka values for some acids and Kb values for their conjugate bases.

The larger the Ka the stronger the acid. Strong acids have Ka values that are essentially infinity. For a conjugate acid / base pair Ka Kb = Kw so strong acids have very weak conjugate bases In general for acid HA Example: The pH of a 0.10 M solution of a weak acid HA is measured as 2.97. Find the Ka value for the acid. If the initial concentration and Ka for a weak acid are known, the equilibrium concentrations and pH of the solution can be calculated.

Example: Find the pH of a 0.150 M solution of CH 3 COOH (aq). Ka for CH 3 COOH is 1.8 x 10-5 Rule of thumb: When the variable x is added to or subtracted from a number that is large compared to Ka and when Ka is less than or equal to 10-4, x may be disregarded. This approximation is considered valid if x ends up being less than 5% of the starting concentration. Otherwise the equation has to be solved exactly usually using the quadratic formula. Example: Determine the pH of a 0.100 M solution of a weak acid HA that has Ka = 2.3 x 10-3 Weak bases partially dissociate in water to produce the conjugate acid and the hydroxide ion General reaction: B (aq) + H 2 O (aq) "^2 -- "^3 HB+ (aq) + OH- (aq) The equilibrium constant is labeled Kb for base dissociation constant and has the form Kb = The most common weak bases are ammonia, NH 3, and related compounds called amines in which one or more of the hydrogen's have been replaced organic groups (CH 3 NH 2 is methyl amine in which a H is replaced with the methyl group, CH 3) NH 3 (aq) + H 2 O (l) "^2 -- "^3 NH 4+ (aq) + OH- (aq) Kb = The larger the value of Kb the stronger the base Example: Calculate the pH of an aqueous solution of 0.15 M NH 3.

Kb for NH 3 = 1.8 x 10-5 Polyprotic Acids are acids that have more than one acidic H per formula unit, e.g. H 2 SO 4, H 2 SO 3, H 2 CO 3, H 3 PO 4 The loss of H+ ions can be viewed as occurring stepwise (one H+ at a time) and a Ka expression can be written for the loss of each H+Example: H 3 PO 4 Ka 1 is always much larger than Ka 2 and Ka 2 is always much larger than Ka 3 Example: Calculate the concentrations of all species present in a 0.100 M solution of arsenic acid, H 3 AsO 4. Ka 1 = 2.5 x 10-4, Ka 2 = 5.6 x 10-8, and Ka 3 = 3.0 x 10-13 Solvolysis = the reaction of a substance with the solvent in which it is dissolved Hydrolysis = the reaction of a substance with water For a conjugate acid / base pair Ka Kb = Kw or p Ka + p Kb = pKwThe stronger an acid the weaker its conjugate base and vice versa The stronger a base the weaker its conjugate acid and vice versa Conjugate bases of weak acids undergo hydrolysis with water A- (aq) + H 2 O (l) "^2 -- "^3 HA (aq) + OH- (aq) Example: CH 3 COO- (aq) + H 2 O (l) "^2 -- "^3 CH 3 COOH (aq) + OH- (aq) The presence of the hydroxide ion makes the solution basic Conjugate bases of strong acids do not undergo hydrolysis Conjugate acids of weak bases undergo hydrolysis Example: NH 4+ (aq) + H 2 O (l) "^2 -- "^3 NH 3 (aq) + H 3 O+The presence of H 3 O+ ions make the solution acidic Conjugate acids of strong bases do not undergo hydrolysis pH of salt solutions Salts of strong acids and strong soluble bases contain the conjugate base of a strong acid and the conjugate acid of strong base neither of which hydrolyze so the salt solution should be neutral (pH = 7) Example: Na Cl, KBr, Ca (NO 3) 2 solutions are all neutral Salts of strong soluble bases and weak acids The conjugate acid does not hydrolyze but the conjugate base does The Kb for the conjugate base is called the hydrolysis constant for the anion Example: Determine the pH of a 0.100 M solution of NaCH 3 COO (aq). Ka for CH 3 COOH is 1.8 x 10-5 Salts of weak bases and strong acids The conjugate base does not hydrolyze but the conjugate acid does The Ka for the conjugate acid is called the hydrolysis constant Example: Determine the pH of a 0.200 M solution of NH 4 Cl (aq). Kb for NH 3 is 1.8 x 10-5 Salts of weak bases and weak acids Both the conjugate acid and conjugate base hydrolyze Whether the solution is acidic or basic depends on whether the hydrolysis is larger for the conjugate acid or the conjugate base If Ka for the conjugate acid is larger than Kb for the conjugate base the solution is acidic If Kb for the conjugate base is larger than Ka for the conjugate acid the solution is basic If the Ka and Kb are equal the solution is neutral Example: Determine the pH of a solution 0.100 M NH 4 ClO (aq). Ka for HClO is 3.5 x 10-8 and Kb for NH 3 is 1.8 x 10-5 Salts that contain small highly charged cations Cations related to insoluble metal hydroxides react with water producing H 3 O+ Example: Al 3+ reacts with water to give Example: Calculate the pH of a solution of 0.10 M Be (NO 3) 2 (aq).

Ka for Be 2+ = 1.0 x 10-5 Strengths of Acids Binary acids: the extent to which a binary acid loses a H+ ion depends on two factors 1) the strength of the H-X bond, the stronger the bond the fewer H+ ions lost 2) the stability of the resulting ions in solution bond strength depends on the polarity and the bond length consider the hydrogen halides HF, HCl, HBr, and HI even though F is the most electronegative of the halogens it is also the smallest and the shorter the bond the stronger the bond in terms of bond strength HF HCl HBr HI because HF has the strongest bond it is the weakest acid of the group (loses only a small portion of H+ ions in aqueous solution) out of the others HI should be the strongest acid and HCl the weakest acid but in aqueous solution they are all strong acids this is because water is basic enough that it does not distinguish between the three acids water is a leveling solvent for HI, HBr, and HCl because they are all completely ionized in water and have the same strength a more acidic solvent (e.g. anhydrous acetic acid) could differentiate between the strengths of the three acids Leveling effect of water: the strongest acid that can exist in water is the H+ (aq) or H 3 O+ (aq) and all acids stronger than this react completely with water to produce H+ (aq) or H 3 O+ (aq) and the conjugate base of the acid Example: HCl + H 2 O -- "^3 Cl- (aq) + H 3 O+ (aq) Water also shows a leveling effect with strong bases because the hydroxide ion is the strongest base that can exist in water Example: CaO + H 2 O -- "^3 Ca 2+ (aq) + 2 OH- (aq) For the Group VIA hydrides the order of bond strengths is H 2 O H 2 S H 2 Se H 2 Te So in this group H 2 O is the weakest acid and H 2 Te is the strongest acid For the binary acids, acid strength increases going down the group. Ternary acids For the ternary acids the H is bonded to O so the bond length is not as important a factor in determining the bond strength Instead the bond strength depends on the bond polarity which can be influenced by two factors 1) the electro negativity of the central element the greater the electro negativity of an element the stronger the acid example: HClO 3 is a stronger acid than Hbr 3 because Cl is more electronegative than Br this is because the more electronegative element draws the electrons farther away from H make the bond more polar and allowing H+ to be lost more easily 2) for acids with the same central element the acid strength depends on the number of oxygen's oxygen is very electronegative so the more O's the more the electrons are withdrawn from the H and the easier H+ is lost the greater the number of O's in the formula the stronger the acid example: HNO 3 is a strong acid and HNO 2 is a weak acid Lewis Theory of Acids and Bases (1923) Lewis acid = any substance that can accept a pair of electrons Lewis base = any substance that can donate a pair of electrons Coordinate covalent bond = a covalent bond in which both electrons come from the same atom or ion The product is sometimes called an acid-base adduct H+ + OH- -- "^3 H 2 O H+ + NH 3 -- "^3 NH 4+ Al Cl 3 + Cl- -- "^3 Al Cl 4- Sn 4+ + 6 Cl- -- "^3 Sn Cl 62-Ligand = a substance that donates a pair of electrons to form a covalent coordinate bond (a. k. a. Lewis base) Coordination complexes = molecules and ions formed by usually transition metal ions forming coordinate covalent bonds with ligands.