Acids and Bases: A Brief Review Acids: taste sour and cause dyes to change color. Bases: taste bitter and feel soapy. Arrhenius: acids increase [H+] bases increase [OH-] in solution. Arrhenius: acid + base salt + water. The Arrhenius definition only works for aqueous solutions. Protons in Solution H+ is simply a proton. (Try to remember to call it that!) H+(aq) is a hydrated proton.
In water, the H+(aq) form clusters. The simplest cluster is H3O+(aq), referred to as the hydronium ion. Larger clusters are H5O2+ and H9O4+. Generally we use H+(aq) and H3O+(aq) interchangeably. Brnsted-Lowry Acids and Bases
Brnsted-Lowry definition focuses on the H+(aq). Acid - something that donates H+ (H+ donor) Base something that accepts H+. (H+ acceptor) The Brnsted-Lowry definition is more general than Arrhenius base because it includes bases other than OH-. Acid/Base Reaction = Proton Transfer Reaction Consider HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq): HCl donates a proton to water. Therefore, HCl is an acid. H2O accepts a proton from HCl. Therefore, H2O is a base. Acid/Base reaction molecular
models Water as an Acid Water reacts with ammonia as an acid: Water can behave as either an acid or a base. Amphoteric substances can behave as acids and bases. Conjugate Acid-Base Pairs HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) Whatever is left of the acid after the proton is donated is called its conjugate base. (It acts as a base in the reverse
reaction) The base after it has accepted a proton is the conjugate acid. (It acts as an acid in the reverse reaction) Conjugate acid-base pairs differ by only one proton. Relative Strengths of Acids and Bases HA + B A- + BH+ HA + B A- + BH+
The more toward the right the equilibrium lies, the stronger the acid. (Greater Keq = stronger acid) The stronger the acid, the weaker the conjugate base. Acid/base equilibria are favored in the direction of stronger acid weaker acid. H+ is the strongest acid that can exist in aqueous solution. OH- is the strongest base that can exist in aqueous solution. Table of Acids and Bases
Relative Strengths of Acids and Bases Any acid or base that is stronger than H+ or OH- simply reacts to produce H+ and OH-. The conjugate base of a strong acid (e.g. Cl-) has negligible acid-base properties. (is neutral) Similarly, the conjugate acid of a strong base has negligible acid-base properties. The Autoionization of Water In pure water the following equilibrium is established H2O(l) + H2O(l)
H3O+(aq) + OH-(aq) at 25 C [H3O ][OH- ] K eq 2 [ H 2O ] K eq [H 2O]2 [H3O ][OH- ] K w [H3O ][OH- ] 1.0 10 14 KW is the ion product of water.
The pH Scale In most solutions [H+(aq)] is quite small. We define pH log[H3O ] log[H ] pOH log[OH- ] In neutral water at 25 C, pH = pOH = 7.00. In acidic solutions, [H+] > 1.0 10-7, so pH < 7.00.
In basic solutions, [H+] < 1.0 10-7, so pH > 7.00. The higher the pH, the lower the pOH, the more basic the solution. Most pH and pOH values fall between 0 and 14. There are no theoretical limits on the values of pH or pOH. (e.g. pH of 2.0 M HCl is -0.301.) pHs of common solutions Other p Scales
In general for a number X, p X log X For example, pKw = -log Kw. K w [H ][OH- ] 1.0 10 14 pK w log [H ][OH- ] 14 log[H ] log[OH - ] 14 pH pOH 14
Measuring pH Most accurate method to measure pH is to use a pH meter. However, certain dyes change color as pH changes. These are indicators. Indicators are less precise than pH meters. Many indicators do not have a sharp color change as a function of pH. Most indicators tend to be red in more acidic solutions. Indicators
Strong Acids The strongest common acids are HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4. Strong acids are strong electrolytes. All strong acids ionize completely in solution: HNO3(aq) + H2O(l) H3O+(aq) + NO3-(aq) Since H+ and H3O+ are used interchangeably, we write HNO3(aq) H+(aq) + NO3-(aq) In solutions the strong acid is usually the only source of H+. (If the molarity of the acid is less than 10-6 M then the autoionization of water needs to be considered.)
Therefore, the pH of the solution is the initial molarity of Strong Bases Most ionic hydroxides are strong bases (e.g. NaOH, KOH, and Ca(OH)2). These are strong electrolytes and dissociate completely in solution. The pOH (and hence pH) of a strong base is given by the initial molarity of the base. Be careful of stoichiometry. In order for a hydroxide to be a base, it must be soluble. Bases do not have to contain the OH- ion: O2-(aq) + H2O(l) 2OH-(aq)
H-(aq) + H2O(l) H2(g) + OH-(aq) N3-(aq) + H2O(l) NH3(aq) + 3OH-(aq) Weak Acids, Ka Weak acids are only partially ionized in aqueous solution. There is a mixture of ions and unionized acid in solution. Therefore, weak acids are in+ equilibrium: HA(aq) + H2O(l) H3O (aq) + A-(aq) [H3O ][A - ] Ka
[HA] HA(aq) H+(aq) + A-(aq) [H ][A - ] Ka [HA] Examples of weak acids Ka is the acid dissociation constant. The larger the Ka the stronger the acid (i.e.
the more ions are present at equilibrium relative to unionized molecules). If Ka >> 1, then the acid is completely ionized and the acid is a strong acid. Calculating Ka from pH These problems are simply equilibrium calculations. The pH gives the equilibrium concentration of H+. Using Ka, the concentration of H+ (and hence the pH) can be calculated. Write the balanced chemical equation clearly showing the equilibrium.
Write the equilibrium expression. Find the value for Ka. Write down the initial and equilibrium concentrations for everything except pure water. We usually assume that the change in concentration of H+ is x. Substitute into the equilibrium constant expression and solve. Remember to turn x into pH if necessary. Percent ionization Percent ionization is another measure of acid strength. Percent ionization relates the equilibrium H+ concentration, [H+]eqm, to the initial HA concentration, [HA]0.
HA(aq) + H2O(l) H3O+(aq) + A-(aq) [H3O ]eqm % ionization 100 [HA]0 The higher percent ionization, the stronger the acid. Percent ionization of a weak acid decreases as the molarity of the solution increases. For acetic acid, 0.05 M solution is 2.0 % ionized whereas Plot of % ionization vs M
Polyprotic Acids Polyprotic acids have more than one ionizable proton. The protons are removed in sequence (not all at once) H2SO3(aq) H+(aq) + HSO3-(aq) Ka1 = 1.7 x 10-2 HSO3-(aq) H+(aq) + SO32-(aq) Ka2 = 6.4 x 10-8
It is always easier to remove the first proton in a polyprotic acid than the second. Therefore, Ka1 > Ka2 > Ka3 etc. Table of Polyprotic Acids Polyprotic Acids Weak Bases, Kb Weak bases remove protons from substances. There is an equilibrium between the base and the resulting ions: Weak base + H2O
conjugate acid + OH Example: NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) The base dissociation constant, Kb, is defined as [BH ][OH - ] Kb [B] [ NH 4 ][OH- ] Kb [ NH3 ]
Types of Weak Bases Bases generally have a non-bonding electron pair (lone pair). This pair is used to from the bond to the H! A negative charge usually indicates a base. i.e., Most anions are basic. The most common neutral weak bases contain nitrogen. Amines are related to ammonia and have one or more NH bonds replaced with N-C bonds (e.g., CH3NH2). Anions of weak acids are also weak bases. Example: OCl- is the conjugate base of HOCl (weak acid): ClO-(aq) + H2O(l)
HClO(aq) + OH-(aq) Kb = 3.3 x 10-7 Relationship Between Ka and Kb When two reactions are added together, the Keq for the combined net reaction is the product of the Keqs K3 K1 K 2 reaction 1 + reaction 2 = reaction 3 For a conjugate acid-base pair HA + H2O A- + H3O+ Ka A- + H2O
HA + OH- Kb - pK pK K a KbH3O+ pK H2O K+wH2O + OH K w
aw b Therefore, the larger the Ka, the smaller the Kb. That is, Ka and Kb for some acid-base pairs Acid-Base Properties of Salt Solutions Soluble salts dissolve by dissociating into the individual hydrated ions.
Acid-base properties of salts are a consequence of the reaction of their ions in solution. The reaction in which ions produce H+ or OH- in water is called hydrolysis. An anion that is the conjugate base of a weak acid is basic. An anion that is the conjugate base of a strong acid is neutral (neither basic nor acidic). An Anions Ability to React with Water Anions, X-, can be considered conjugate bases from
acids, HX. If X- comes from a strong acid, then it is neutral. If X- comes from a weak acid, then X-(aq) + H2O(l) HX(aq) + OH-(aq) The pH of the solution can be calculated using equilibrium! A Cations Ability to React with Water Polyatomic cations with ionizable protons can be considered conjugate acids of weak bases.
NH4+(aq) + H2O(l) NH3(aq) + H3O+(aq) Therefore, they react as weak acids and lower the solution pH. Some metal ions (all except the group I and II metals) can also react in solution to lower pH. Combined Effect of Cation and Anion in Solution An anion from a strong acid has no acid-base properties. An anion that is the conjugate base of a weak acid will cause an increase in pH.
A cation that is the conjugate acid of a weak base will cause a decrease in the pH of the solution. Metal ions will cause a decrease in pH except for the alkali metals and alkaline earth metals. When a solution contains both cations and anions from weak acids and bases, use Ka and Kb to determine the final pH of the solution.
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