Acids, Bases, and Salts By: Chris Squires

Effects of Acid Rain on Marble (marble is calcium carbonate CaCO 3 ) George Washington: BEFORE acid rain AFTER acid rain

Acids have a pH less than 7

OPERATIONAL ACIDS Reacts with metals to form H 2 gas. HCl (aq) + Mg (s) → MgCl 2(aq) + H 2(g) HCl (aq) + Mg (s) → MgCl 2(aq) + H 2(g)

Taste Sour

Conducts electricity.

Neutralizes BASES

Acids Turn Litmus Red Blue litmus paper turns red in contact with an acid

Acids React with Carbonates and Bicarbonates HCl + NaHCO 3 NaCl + H 2 O + CO 2 Hydrochloric acid + sodium bicarbonate salt + water + carbon dioxide An old-time home remedy for relieving an upset stomach

OPERATIONAL BASES

BASES ARE SLIPPERY

BASES are Corrosive

BASES

BASES OPERATE

Bases have a pH greater than 7

Bases Turn Litmus Blue Red litmus paper turns blue in contact with a base (and blue paper stays blue).

Properties of Bases (metallic hydroxides) n React with acids forms water and salt. n Taste bitter & feel slippery n Electrolytes in aqueous n Change the color of indicators (bases go Blue, with litmus).

Examples of Bases  Sodium hydroxide, NaOH (lye for drain cleaner; soap)  Potassium hydroxide, KOH (alkaline batteries)  Magnesium hydroxide, Mg(OH) 2 (Milk of Magnesia)

Theories Acid Base H+ H+H+

Acid-Base Theories n a) Arrhenius, n b) Brønsted-Lowry n c) Lewis (not covered).

Svante Arrhenius ( )

1. Arrhenius Definition n Acids produce (H + ) in solution (HCl → H + + Cl - ) n Bases produce OH - in water. (NaOH → Na + + OH - ) n Limited to aqueous solutions. n Only one kind of base (hydroxides) n NH 3 (ammonia) could not be an Arrhenius base: no OH - produced.

3 flaws with Arrhenius 1. A proton (H + ) does not exist ALONE in water, it makes H 3 O + (aq) 2. NH 3 (g) is a base when dissolved in water; but does not contain the OH - (aq) 3. Amphiprotics like HCO 3 - (aq) appear to contain the H + ion BUT ARE bases

Johannes Brønsted Thomas Lowry ( ) ( ) Denmark England

Brønsted-Lowry n A broader definition than Arrhenius n Acid is a H +,base is H + acceptor. n HCl is an acid. –When it dissolves in water, it gives it’s proton to water. HCl (g) + H 2 O (l) ↔ H 3 O + (aq) + Cl - (aq) HCl (g) + H 2 O (l) ↔ H 3 O + (aq) + Cl - (aq) n Water is a base; makes hydronium ion.

A Brønsted–Lowry acid… …must have removable proton. HCl, H 2 O, H 2 SO 4 A Brønsted–Lowry base must accept this H+ NH 3, H 2 O, CO 3 2- and usually has negative charge

Acids and bases come in pairs n A “conjugate base” is the remainder of the original acid, after it donates it’s hydrogen ion n A “conjugate acid” is the particle formed when the original base gains a hydrogen ion

Conjugate Acids and Bases: n From the Latin word conjugare, meaning “to join together.” n Reactions between acids and bases always yield their conjugate bases and acids.

If can also be BOTH Acid and Base…...it is amphiprotic. HCO 3 – HSO 4 – H2OH2OH2OH2O

Who Theory: Acid When Arrheniusincreases H ’s Brønstedproton donor1923 Lewis electron pair acceptor same1923 Three THEORETICAL definitions

Hydrogen Ions from Water n Water ionizes, or falls apart into ions: H 2 O ↔ H + + OH - n Called the “self ionization” of water n Occurs to a very small extent: [ H + ] = [OH - ] = 1 x M n Since they are equal a neutral solution n K w = [H + ] x [OH - ] = 1 x n 14 = pH + pOH

pH and Significant Figures n [H + ] = M = 1.0 x M has 2 sig figs has 2 sig figs n the pH = 3.00, with the two numbers to the right of the decimal corresponding to the 2 sig figs

STRONG V WEAK

STRONG VS WEAK

Strength In any acid-base reaction, the equilibrium favors the reaction that moves the proton to the stronger base. HCl ( aq ) + H 2 O ( l )  H 3 O + ( aq ) + Cl – ( aq) Cl – equilibrium lies so far to the right K is not measured ( K >>1), so is not a base, it is neutral

Strength n Strong acids completely dissociate in water. –Their conjugate bases are actually so weak as not to be able to act as a base at all.

Strength n Acids & Bases are classified according to the degree they ionize in water: –Strong are completely ionized, 100% –Weak ionize only slightly, like 1% Strength & Concentration Night V Day

Strong Acid Dissociation (makes 100 % ions)

Weak Acid Dissociation (only partially ionizes)

STRONG V WEAK ACIDS

LOL

Measuring strength n Ionization is reversible for WEAK acids: HA + H 2 O ↔ H + + A - n This makes an equilibrium n K a = [H + ][A - ] [HA] n Strong acid = more products & large K a (Note that water is NOT shown, (Note that the arrow goes both directions.)

Calculating pH from K a Calculate the pH of a 0.30 M solution of acetic acid, C 2 H 3 O 2 H, at 25°C. K a for acetic acid at 25°C is 1.8  K a for acetic acid at 25°C is 1.8 

Calculating pH from K a Use the ICE table: [C 2 H 3 O 2 ], M[H 3 O + ], M[C 2 H 3 O 2 − ], M Initial Change–x–x+x+x+x+x Equilibrium0.30 – xxx

Calculating pH from K a Use the ICE table: [C 2 H 3 O 2 ], M[H 3 O + ], M[C 2 H 3 O 2 − ], M Initial Change–x–x+x+x+x+x Equilibrium0.30 – x ≈ 0.30 xx Simplify: how big is x relative to 0.30?

Calculating pH from K a Now, (1.8  ) (0.30) = x  = x  = x pH = = 2.64

Calculating K a from the pH n The pH is 2.38 for a 0.10 M solution of formic acid, HCOOH,. Calculate K a K a, requires equilibrium concentrations. We can find [H 3 O + ], which is the same as [HCOO − ], from pH, so pH = X

Calculating K a from the pH pH = –log [H 3 O + ] – 2.38 = log [H 3 O + ] = 10 log [H 3 O + ] = [H 3 O + ] 4.2  = [H 3 O + ] = [HCOO – ] = X

Calculating K a from pH All values In Mol/L [HCOOH][H 3 O + ][HCOO − ] Initially Change –4.2   Equilibrium 0.10 – 4.2  = =  10 -3

Calculating K a from pH [4.2  ] [0.10] K a = = 1.8  10 -4

Calculating Percent Ionization [A - ] eq = [H 3 O + ] eq = 4.2  M [A - ] eq = [H 3 O + ] eq = 4.2  M [A - ] eq + [HCOOH] eq = [HCOOH] initial = 0.10 M [A - ] eq + [HCOOH] eq = [HCOOH] initial = 0.10 M %rx = [4.2  / 0.10] (x 100) %rx = [4.2  / 0.10] (x 100) = 4.2%

What about bases? n Strong bases dissociate completely. n MOH + H 2 O ↔ M + + OH - (M = a metal) n Base dissociation constant = K b n K b = [M + ][OH - ] [MOH] n Stronger base = more dissociated ions are produced, thus a n larger K b and a small Ka.

Weak Bases The Kb for this reaction is The Kb for this reaction is

pH of Basic Solutions What is the pH of a 0.15M solution of NH 3 ? [NH 4 + ] [OH − ] [NH 3 ] K b = = 1.8  10 -5

pH of Basic Solutions [NH 3 ][NH 4 + ][OH − ] Initial Equil x  0.15 xx

pH of Basic Solutions Ignore X (1.8  ) (0.15) = x  = x  = x 2 ( x ) 2 (0.15-X) 1.8  = pOH = –log (1.6  ) pOH = 2.80 pH = 11.20

Strength vs. Concentration n The words concentrated and dilute tell how much of an acid or base is dissolved in solution - refers to the # of moles of acid or base n Is a concentrated, weak acid possible?

Acid and Base Strength Acetate is a stronger base than H 2 O, LOWER on table, The stronger base “wins” the proton. HC 2 H 3 O 2 ( aq ) + H 2 OH 3 O + ( aq ) + C 2 H 3 O 2 – ( aq)

TITRATIONS

Titration n Titration is the process of adding a known amount of solution of known concentration to determine the concentration of another solution n The equivalence point is when the moles of hydrogen ions is equal to the moles of hydroxide ions (= neutralized!)

Titration n The concentration of acid (or base) in solution can be determined by performing a neutralization reaction –An indicator is used to show when neutralization has occurred –Often we use phenolphthalein- because it is colorless in neutral and acid; turns pink in base

Steps - Neutralization reaction #1. A measured volume of acid of unknown concentration is added to a flask #2. Several drops of indicator added #3. A base of known concentration is slowly added, until the indicator changes color; measure the volume

Neutralization n The solution of known concentration is called the standard solution –added by using a buret n Continue adding until the indicator changes color –called the “end point” of the titration

strong acid (HCl) v. strong base (NaOH)

Very little pH change during the initial 20cm 3 Very sharp change in pH over the addition of less than half a drop of NaOH pH 1 as 0.1M HCl (strong acid) strong acid (HCl) v. strong base (NaOH)

Steady pH change Sharp change in pH over the addition of less than half a drop of NaOH Curve levels off at pH 13 due to excess 0.1M NaOH (a strong alkali) pH 4 due to 0.1M CH 3 COOH (weak monoprotic acid) w. acid (CH 3 COOH) v s.base (NaOH)

strong acid (HCl) v. strong base (NaOH) Any of the indicators listed will be suitable - they all change in the ‘vertical’ portion PHENOLPHTHALEIN LITMUS METHYL ORANGE

weak (CH 3 COOH) v. weak (NH 3 ) Types Steady pH change pH 4 due to 0.1M CH 3 COOH (weak monoprotic acid) NO SHARP CHANGE IN pH Curve levels off at pH 10 due to excess 0.1M NH 3 (a weak alkali)

Polyprotic Acids? n Some compounds have more than one ionizable hydrogen to release n HNO 3 nitric acid - monoprotic n H 2 SO 4 sulfuric acid - diprotic - 2 H + n H 3 PO 4 - triprotic - 3 H + n IMPT for TITRATING BUT NOT FOR pH calculating…why?

Polyprotic Acids n Here are three successive ionizations of phosphoric acid: n Here are three successive ionizations of phosphoric acid: n The first dissociation constants for phosphoric acid is greater than the second, 100,000 times greater n This means nearly all the H+ ions in the solution comes from the first step of dissociation. H 3 PO 4(s) + H 2 O (l) H 3 O + (aq) + H 2 PO 4 − (aq) K a1 = 7.25×10 −3 H 2 PO 4 − (aq) + H 2 O (l) H 3 O + (aq) + HPO 4 2− (aq) K a2 = 6.31×10 −8 HPO 4 2− (aq) + H 2 O (l) H 3 O + (aq) + PO 4 3− (aq) K a3 = 3.98×10 −13

Other pH curves - Other pH curves - acid v. carbonate Sodium carbonate reacts with hydrochloric acid in two steps... Step 1 Na 2 CO 3 + HCl ——> NaHCO 3 + NaCl Step 2 NaHCO 3 + HCl ——> NaCl + H 2 O + CO 2 Overall Na 2 CO 3 + 2HCl ——> 2NaCl + H 2 O + CO 2 First rapid pH change around pH = 8.5 due to the formation of NaHCO 3. Can be detected using phenolphthalein Second rapid pH change around pH = 4 due to the formation of acidic CO 2. Can be detected using methyl orange. There are two sharp pH changes

Polyprotic acids (H 3 PO 4 ) There are three sharp pH changes Each successive addition of NaOH is the same as equal number of moles are involved. Phosphoric acid is triprotic; it reacts with sodium hydroxide in three steps... Step 1 H 3 PO 4 + NaOH ——> NaH 2 PO 4 + H 2 O Step 2 NaH 2 PO 4 + NaOH ——> Na 2 HPO 4 + H 2 O Step 3 Na 2 HPO 4 + NaOH ——> Na 3 PO 4 + H 2 O

Other pH curves - polyprotic acids (H 3 PO 4 ) pH of H 3 PO 4 = 1.5 Phosphoric acid is triprotic; it reacts with sodium hydroxide in three steps... Step 1 H 3 PO 4 + NaOH ——> NaH 2 PO 4 + H 2 O Step 2 NaH 2 PO 4 + NaOH ——> Na 2 HPO 4 + H 2 O Step 3 Na 2 HPO 4 + NaOH ——> Na 3 PO 4 + H 2 O

Other pH curves - polyprotic acids (H 3 PO 4 ) pH of NaH 2 PO 4 = 4.4 pH of H 3 PO 4 = 1.5 Phosphoric acid is triprotic; it reacts with sodium hydroxide in three steps... Step 1 H 3 PO 4 + NaOH ——> NaH 2 PO 4 + H 2 O Step 2 NaH 2 PO 4 + NaOH ——> Na 2 HPO 4 + H 2 O Step 3 Na 2 HPO 4 + NaOH ——> Na 3 PO 4 + H 2 O

Other pH curves - polyprotic acids (H 3 PO 4 ) pH of Na 2 HPO 4 = 9.6 pH of NaH 2 PO 4 = 4.4 pH of H 3 PO 4 = 1.5 Phosphoric acid is triprotic; it reacts with sodium hydroxide in three steps... Step 1 H 3 PO 4 + NaOH ——> NaH 2 PO 4 + H 2 O Step 2 NaH 2 PO 4 + NaOH ——> Na 2 HPO 4 + H 2 O Step 3 Na 2 HPO 4 + NaOH ——> Na 3 PO 4 + H 2 O

Other pH curves - polyprotic acids (H 3 PO 4 ) pH of Na 3 PO 4 = 12 pH of Na 2 HPO 4 = 9.6 pH of NaH 2 PO 4 = 4.4 pH of H 3 PO 4 = 1.5 Phosphoric acid is triprotic; it reacts with sodium hydroxide in three steps... Step 1 H 3 PO 4 + NaOH ——> NaH 2 PO 4 + H 2 O Step 2 NaH 2 PO 4 + NaOH ——> Na 2 HPO 4 + H 2 O Step 3 Na 2 HPO 4 + NaOH ——> Na 3 PO 4 + H 2 O

Salt Hydrolysis n A salt is an ionic compound that: –formed in a neutralization reaction –some neutral; others acidic or basic

Salt Hydrolysis n To see if the resulting salt is acidic or basic, check the by dissociating the salt NaCl, a neutral salt (NH 4 ) 2 SO 4, acidic salt CH 3 COOK, basic salt

Adding a “salt” to water n If the salt is from a strong acid and base then the pH is 7, for example KNO 3. n A salt formed between a strong acid and a conjugate of a weak base is an acid salt, example NH 4 Cl, pH =acidic. n A salt formed between a conjugate of a weak acid and a strong base is a basic salt, for example NaCH 3 COO, and the pH will be basic.

Salt solutions: hydrolysis NaCl ( aq ) NH 4 Cl ( aq ) NaClO ( aq )

Buffers n Buffers are solutions in which the pH remains relatively constant, even when small amounts of acid or base are added –made from a pair of chemicals: a weak acid and one of it’s salts; or a weak base and one of it’s salts

Buffers n A buffer system is better able to resist changes in pH than pure water n It is a pair of chemicals (conjugates) –one chemical neutralizes any acid added, while the other chemical would neutralize any base

Buffers n Example: Ethanoic (acetic) acid and sodium ethanoate (also called sodium acetate) n The buffer capacity is the amount of acid or base that can be added before a significant change in pH

Buffers n The two buffers that are crucial to maintain the pH of human blood are: 1. carbonic acid (H 2 CO 3 ) & hydrogen carbonate (HCO 3 - ) 2. dihydrogen phosphate (H 2 PO 4 1- ) & monohydrogen phoshate (HPO 4 2- )

Aspirin (which is a type of acid) sometimes causes stomach upset; thus by adding a “buffer”, it does not cause the acid irritation. Bufferin is one brand of a buffered aspirin that is sold in stores.

Indicators 1. Moisten the pH indicator paper strip with a few drops of solution, by using a stirring rod. 2.Compare the color to the chart on the vial – then read the pH value.

Some of the many pH Indicators and their pH range

How Do We Measure pH? –An indicator Compound that changes color in solution.