ACIDS & BASES Kenneth E. Schnobrich

ACIDS - OPERATIONAL DEFINITION Have a sour taste (not recommended) React with bases to form water and a salt (Neutralizationn) React with some metals to form a salt and H2 gas (see Table J) Reacts with indicators Litmus Blue -> Red Phenolphthalein - colorless

ACIDS - OPERATIONAL DEFINITION In aqueous solution they are good electrolytes (conduct current) Can be very corrosive and can cause severe burns

BASES - OPERATIONAL DEFINITION Bases have a bitter taste (not recommended) They have a slippery feel (soapy) In aqueous solution they are good electrolytes (conduct electric current) React with acids to form a salt and water (neutralization) Cause indicators to change color Litmus Red -> Blue Phenolphthalein - pink

ACIDS - CONCEPTUAL DEFINITION (Arrhenius) ACIDS - are substances that will yield H+ as the only positive ion in an aqueous solution. HCl(aq) --> H+(aq) + Cl-(aq) H2SO4(aq) --> 2H+(aq) + SO4-2(aq) H3PO4(aq) --> 3H+1(aq) + PO4-3(aq) Monoprotic acid Diprotic acid Triprotic acid

BASES - CONCEPTUAL DEFINITION (Arrhenius) BASES - are substances that will yield OH- as the only negative ion in an aqueous solution. NaOH(s) --> Na+(aq) + OH-(aq) Ca(OH)2(s) --> Ca+2(aq) + 2OH-1(aq) Fe(OH)3(s) --> Fe+3(aq) + 3OH-1(aq)

Acid/Base Reference Tables These are considered strong acids because they ionize almost completely in an aqueous solution. They yield a high concentration of H+ or H3O+ ions in aqueous solutions. These are considered moderate to weak acids because they do not ionize completely in an aqueous solution. They yield a lower concentration of H+ or H3O+ ions in aqueous solutions. These are considered strong bases because they ionize (dissociate) almost completely in an aqueous solution. They yield a high concentration of OH - ions in aqueous solutions. These are considered weak bases because they ionize (dissociate) to a small degree in an aqueous solution. They yield a low concentration of OH - ions in aqueous solutions.

Nature of the Hydrogen Ion In an aqueous solution the hydrogen ion NEVER exists by itself - it forms the Hydronium Ion (H3O+1) H2O H+1 H3O+1 O x H +1 H+1 H O x H x x H H

Arrhenius Theory ACIDS - substances whose water solutions (aqueous) contain Hydrogen ions (Hydronium ions) as the only positive ion. CHART K - contains a list of common Arrhenius acids Listed by strength (strongest ->weakest) BASES - substances whose water solutions (aqueous) contain Hydroxide ions as the only negative ion. CHART L - contains a list of common Arrhenius bases

ACID/BASE STRENGTH The Strength of an acid or base is determined by how much it ionizes in an aqueous solution. HCl, HNO3, H2SO4 - are considered strong because they ionize almost 100% in an aqueous solution KOH, NaOH - are considered strong bases because they ionize (dissociate) almost 100% in a aqueous Solution Therefore, strong acids and bases are good electrolytes in an aqueous solution

BRONSTED THEORY HCl + H2O -> H3O+1 + Cl-1 The Bronsted-Lowry Theory is an expansion of the definition of acids and bases (not a replacement). ACIDS - substances that will donate a proton (all Arrhenius acids are also Bronsted acids) BASES - substances that will accept a proton (all Arrhenius bases are also Bronsted bases) HCl + H2O -> H3O+1 + Cl-1

REACTIONS Acids with Metals TABLE J of your Reference Tables for Chemistry is important in determining what metals will react with an acid - any metal above **H2 on Table J will react with an acid to form H2 gas and a salt. HCl(aq) + Mg(s) -> MgCl2(aq) + H2(g)

Activity Series All of the metals listed above **H2 on this activity series will react to replace hydrogen from an acid to produce H2 as one of the products of the reaction. 2HCl(aq) + Zn(s) -> Zn+2(aq) + 2Cl-(aq) + H2(g)

REACTIONS Neutralization Neutralization - an Arrhenius acid reacts with an Arrhenius base to form a salt and water as the products of the reaction H+(aq) + Cl-(aq) Na+(aq) + OH-(aq) HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l) **Na+1 ions and Cl-1 ions are Spectator Ions

pH/Acid/Base - Strength pH = -log of the [H+] in a solution. The pH scale describes the concentration of [H+] in an aqueous solution. The scale is based on neutral water where the [H+] = [OH-] concentration. HOH(l) -> H+(aq) + OH-(aq) **Important note: remember H+ = H3O+ in aqueous solution

pH SCALE [H]+>[OH]- [H]+=[OH]- [H]+<[OH]- Neutral Becoming more Acidic Becoming more Basic 7 14 [H]+>[OH]- [H]+=[OH]- [H]+<[OH]- *Remember [ ] = concentration in mols/L

pH and Indicators Indicators tend to be weak organic acids – HIn(aq) H+(aq) + In-(aq) A change in color indicates a change in the pH of the solution by shifting the equilibrium above to decrease or increase the amounts of HIn or In- in solution. Color 1 Color 2 If you started with a pH of 2.0 and used bromcresol green as your indicator the solution would appear yellow. If you now added a base like NaOH, as the pH increases to something between 3.8 and 5.4, the solution would appear green in color (a combination of yellow and blue).

Titration using an Indicator Acid-Base titrations are very common in chemistry labs. What is a titration? A titration is a procedure in which a base is added to an acid (or vice versa), a neutralization reaction occurs; an indicator is used to determine the point at which you have equivalent quantities of H+(H3O+) and OH-. A commonly used acid-base indicator is Phenolphthalein. Using Phenolphthalein, you are looking for a faint pink color to indicate the so-called “end point” or “equivalence point”. At the “end point” even a trace amount of base will cause a drastic change in color because of the equilibrium shift. HIn (aq) H+(aq) + In-(aq) Colorless in acid solution Deep pink in base solution Phenolphthalein

Titration In the burette you would place a known quantity of base like NaOH (unknown molarity) – deliver it slowly using the stopcock You would add base until a very faint pink color appears in the flask – the equivalence point. In the Erlenmeyer flask you would have an acid solution with a known volume and molarity. You would also have 4-5 drops of Phenolphthalein indicator (it would appear colorless in the acid solution)

Titration Calculation Since we assume at the equivalence point that the [H+] = [OH-] we can say #mols H+ = #mols OH-. M(V) = # mols at the equivalence point MaVa = MbVb (see Table T) If we had 20 mL of a 1.0M acid and it took 40 mL of the base, what is the molarity of the base? (1.0 M)(20 mL) = (X)(40 mL) X = (1.0 M)(20 mL)/40 mL = 0.5M for the base

Titration Calculation #2 There are times when the titration uses a strong diprotic acid (2 mols of H+(H3O+)/mol of acid) like sulfuric acid. H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l) + + +  + + 2H3O+ SO4-2 2Na+ 2OH- 2Na+ SO4-2 2H2O

Titration Calculation #2 On the previous slide you notice it takes 2 mols of NaOH to neutralize the 2 mols of H+ (H3O+) using the diprotic acid, sulfuric acid (H2SO4). If we have 20.0 mL of 1.0M sulfuric acid in the flask and it takes 40 mL of the base, NaOH, what is the molarity of the base? MaVa = MbVb 2.0M(20.0 mL) = X(40 mL) X = 2.0M(20.0 mL)/40 mL) = 1.0MNaOH Notice the adjusted molarity of the acid to 2.0M because it is a diprotic acid. Sulfuric acid will yield 2 mols of hydrogen ions (hydronium ions) per mol of acid ionized.

Kw for Water The Kw for H2O = 1 x 10-14 at 25°C. This means that the [H+] and the [OH-] equals 1 x 10-7. This means that in an aqueous solution the [H+][OH-] = 1 x 10-14. In a 0.001M HCl solution the concentration of [H+] = 1 x 10-3 and the [OH-] = 1 x 10-11 because the product of the two concentrations must always equal 1 x 10-14 in an aqueous solution. The pH of this solution would be 3. pH = -log [H+] = -(log 1 x 10-3) = -(-3) = 3 At pH of 3 it would be considered a fairly strong acid solution HCl H2O H3O+ Cl-

Hydrolysis of a Salt Hydrolysis of a salt is the reverse of neutralization. In neutralization, an acid and a base react to form a salt and water. Salts can be formed by reactions of strong acid and a strong base strong acid and a weak base strong base and weak acid Weak acid and weak base In cases 1, 2, and 3 it is easy to predict the outcome of hydrolysis of the resulting salt. In case 4 you would have to know the values for the strengths of the acid and base (information you do not have).

Hydrolysis Case #1 NaCl(s) + HOH(l) = Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) NaCl is the salt of a strong acid and a strong base so when it dissolves in water there will be an equal number of H+ and OH- ions in solution. The aqueous solution would have a pH of 7.

Hydrolysis Case #2 NH4NO3(s) + HOH = NH4+(aq) + OH-(aq) + H+(aq) + NO3-(aq) Since NH4NO3 is the salt of a weak base, NH4OH and a strong acid, HNO3 the salt will yield a higher concentration of H+ in the aqueous solution than the OH- ions. The solution of NH4NO3 with have an acidic pH.

Hydrolysis Case #3 K2CO3(s) + 2HOH(l) = 2K+(aq) + 2OH-(aq) + 2H+(aq) + CO3-2(aq) Since K2CO3 is the salt of a strong base, KOH and a weak acid, H2CO3 the salt will yield a higher concentration of OH- in the aqueous solution than the H+ ions. The solution of K2CO3 with have an basic pH.