Acids Bases and Salts Chapters 14 and 15

Acids, Bases, and Salts Chapters 14 and 15

Some Properties of Acids 1. The word acid comes from the Latin word acere, which means "sour. " All acids taste sour. 2. In 1663, Robert Boyle wrote that acids would make a blue vegetable dye called "litmus" turn red. 3. Acids react with bases (they destroy the chemical properties of bases). 4. Acids conduct an electric current. 5. Upon chemically reacting with an active metal, acids will evolve hydrogen gas (H 2).

Some Properties of Bases 1. The word "base" has a more complex history and its name is not related to taste. All bases taste bitter. 2. Bases are substances which will restore the original blue color of litmus after having been reddened by an Thisacid. is because they dissolve the fatty 3. Bases destroyacids theand chemical oils fromproperties your skin andofthis acids (will react cutswith downacids) on the friction between your fingers you rub them together. 4. Bases will conduct anaselectric current. 5. Bases feel “slippery” (soap, bleach) on your skin.

Some Properties of Salts 1. A salt is the combination of an anion (- ion) and a cation (+ ion). 2. Salts are products of the reaction between acids and bases. 3. Solid salts are usually crystalline. 4. If a salt dissolves in water solution, it usually dissociates into the anions and cations that make up the salt (amount depends on Ksp)

The Acid Base Theory The three main theories regarding acids and bases are… (from least to most inclusive): 1. Arrhenius 3. Lewis 2. Brønsted-Lowry

Arrhenius Theory – late 1890 s DEFINITIONS: Acid - any substance which donates hydrogen ions (H+) to water (produces hydronium ions, H 3 O+): HA → H+ + A¯ Base - any substance which produces hydroxide ions (OH¯) in water. XOH → X+ + OH¯ When acids and bases react, they neutralize each other, forming water and a salt: HA + XOH → H 2 O + XA

Problems with Arrhenius Theory The theory did not explain why ammonia (NH 3) was a base (… OH- was the only base so far) n The theory only considers water as a solvent. We know that an acid added to benzene will not dissociate. Solvents are crucial to acid definition. n The end result of mixing certain acids and bases can be a slightly acidic or basic solution. Arrhenius had no explanation for this phenomenon (degrees of acidity). n

Brønsted – Lowry Theory – Early 1920 s Two chemists, independent of one another, proposed a new definition of an acid and a base: n An acid is a substance from which a proton can be removed (or donates protons). n A base is a substance that can remove a proton from an acid (or a proton acceptor). *This definition does not require acids and bases to be in aqueous solutions. n

Reactions Based on Bronsted - Lowry Which are the acids and bases? : HCl + H 2 O → H 3 O+ + Cl¯ n HCl - this is an acid, because it has a proton available to be transferred (it can give a proton). n H 2 O - this is a base, since it gets the proton that the acid lost (it has the capacity to accept a proton). n

Conjugate acid-base pairs Example: HCl + H 2 O → H 3 O+ + Cl¯ Notice that each pair (HCl and Cl¯ as well as H 2 O and H 3 O+ differ by one proton (H+). These pairs are called conjugate pairs. Example: HNO 3 + H 2 O → H 3 O+ + NO 3¯ n The acids are HNO 3 and H 3 O+ and the bases are H 2 O and NO 3¯. What are the pairs? n

Bases and Conjugate Acid Base Name C 2 H 3 O 2 - Acetate ion Conjugate acid Name CH 3 COOH Acetic acid NH 3 H 2 PO 4 - Ammonia Dihydrogen phosphate ion NH 4+ H 3 PO 4 HSO 4 - Hydrogen sulfate ion H 2 SO 4 Ammonium Phosphoric acid Sulfuric acid OHNO 3 H 2 O Hydroxide ion Nitrate ion water H 2 0 HNO 3 H 3 0 + water Nitric acid Hydronium ion

Lewis Theory –Early 1920 s Remember drawing Lewis Dot Structures for ionic and covalent compounds? n Lewis Theory focuses on the nature of electrons rather than proton transfer. DEFINITIONS: n An acid is an electron pair acceptor and a base as an electron pair donor. n Lewis Theory is much more general and will include also reactions that do not involve hydrogen or hydrogen ions… this will include formation of coordinate complex ions. n

Lewis acid-base reaction: BF 3 accepts an electron pair from ammonia to complete octet: A Lewis acid must have an empty orbital to accept an electron pair –all cations are L. A. s, and ions/atoms with incomplete octets. A Lewis base must have a pair of unshared electrons that can be donated. Typical Lewis bases are OH-, H 2 O, NH 3, Cl-, CN-… due to lone pair electrons.

Lewis AB reactions and Formation of Coordinate Complexes The metal ion is a Lewis acid and the ligands coordinated to the ion are Lewis bases.

Autoionization of Water

What are Strong Acids and Bases? n Strong acids are those that ionized completely in water. . . the H+ ions are not held strongly to the conjugate base. • The dissociation of a strong base also looks like the diagram at the right in that it dissociates into positive and negative ions.

7 Strong Acids HNO 3 - nitric acid HCl - hydrochloric acid HBr - hydrobromic acid HI - hydroiodic acid H 2 SO 4 - sulfuric acid HCl. O 4 - perchloric acid HCl. O 3 - chloric acid (wanna be) • Strong acids are assumed to ionize completely (100%) in water. They exist as H 3 O+ ions in water. This is known as “the leveling effect”. Water has a greater affinity for H+ than the conjugate bases do.

ANIMATION LINKS n Acid ionization equilibrium demo

What are Weak Acids and Bases? Some acids and bases ionize only slightly in water and are considered weak. n A weak acid has a high affinity for its H+. n The most important weak base is ammonia. n

Balance of ions in solutions Acidic Solution Neutral Solution

Strong Bases Li. OH - lithium hydroxide Na. OH - sodium hydroxide KOH - potassium hydroxide Rb. OH - rubidium hydroxide Cs. OH - cesium hydroxide Ba(OH)2 - barium hydroxide Sr(OH)2 - strontium hydroxide Ca(OH)2 - calcium hydroxide Uncommon in labs because too expensive GROUP 1 hydroxides Some GROUP 2 hydroxides • Strong bases also ionize completely in water, except for Sr(OH)2 and Ca(OH)2 which are only slightly soluble (remember Mg(OH) is insoluble, K =2 x 10 -11).

Polyprotic Acids Polyprotic acids dissociate in a stepwise fashion with different Ka values for each step… In the second and subsequent ionizations the acids are always weak, whether or not the original is a strong or weak acid. n For most of these acids (ex. H 3 PO 4), the first dissociation contributes the significant amount of H+ for p. H calculations, and the rest are negligible (except for H 2 SO 4 where second ionization is significant). n

Naming Acids -REVIEW -ide ending (elements): “hydro____ic acid” ex. chloride (HCl): hydrochloric acid -ate ending (polyatomics): “______ic acid” ex. chlorate (HCl. O 3): chloric acid -ite ending(polyatomics): “______ous acid” ex. chlorite (HCl. O 2): chlorous acid

Net Ionic Equations -REVIEW For aqueous acid-base reactions, it is common to write equations in the net ionic form. Standard form: Na. OH(aq) + HCl(aq) Na. Cl(aq) + H 2 O(l) Ionic form: Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) Na+(aq) + Cl-(aq) + H 2 O(l) Net ionic form: OH-(aq) + H+ (aq) H 2 O(l) (No spectator ions are included) n

Things to remember when writing Net Ionic Equations n n Binary Acids: HCl, HBr, and HI are strong: all other binary acids and HCN are weak. Strong acids are written in ionic form; weak acids are written in molecular form. Ternary Acids: If the number of oxygen atoms in an inorganic acid molecule exceeds the number of hydrogen atoms by two or more, the acid is strong (complete dissociation). We will consider all organic acids as weak. Strong: HCl. O 3, HCl. O 4, H 2 SO 4, HNO 3 Weak: HCl. O, H 3 As. O 4, H 2 CO 3, H 4 Si. O 4, HNO 2

Polyprotic Acids: (acids that contain more than one ionizable hydrogen atom. Ex: H 2 SO 4, H 3 PO 4, H 2 CO 3). n Bases: Hydroxides of Group 1 and 2 elements (except Be(OH)2 and Mg(OH)2) are strong bases. All others including ammonia, hydroxlamine, and organic bases are weak. n Salts: Salts are written in ionic form if soluble, and in undissociated form if insoluble. *Know the solubility rules. n Oxides: Oxides are always written in molecular or undissociated form (ex: Mg. O). n Gases: Gases are always written in molecular form (ex: SO 2). n

Practice Net Ionic Equations 1. Ag. NO 3 (aq) + H 2 SO 4 (aq) 2. H 4 Si. O 4 (aq) + Na. OH (aq) 3. HBr (aq) + KOH (aq) 1. Ag+ + HSO 4 - → Ag. HSO 4(s) 2. H 4 Si. O 4 + OH- → H 3 Si. O 4 - + H 2 O 3. H+ + OH- → H 2 O

How do we determine the strength (or p. H) of a weak acid?

Relationship between Ka and Kb Ka x K b = K w n For any acid and it’s conjugate base, this relationship can be used to determine Ka or Kb. n Ex: NH 3 + H 2 O ↔ NH 4+ + OHNH 4+ + H 2 O ↔ NH 3 + H 3 O+ Kb(NH 3)=[NH 4+][OH-] Ka(NH 4+)=[NH 3][H 3 O+] [NH 3] [NH 4+] n Therefore, Ka x Kb = [OH-][H 3 O+]= Kw

Weak acids and bases will have Ka or Kb values less than one, but greater than water dissociation, Kw (essentially Ka and Kb of H 2 O)

Sorenson (1868 -1939) invented the p. H scale while creating a wayp. H to test the (potential hydrogen scale) n The scale acidity of beer. Beer +) measure of hydronium ion (H O 3 has a p. H of about concentration. 4. 5. p. H Scale is a Hydronium ion concentration indicates acidity. Each increase in p. H # means a 10 fold decrease in [H+]. n The higher the [H 3 O+], the higher the acidity. n

p. H scale and [H+]

Calculating p. H The concentration (M or mol/L) of H 3 O+ is expressed in powers of 10, from 10 -14 to 100. n Scientists use p. H, which is the negative log of [H 3 O+]. n p. H = -log[H 3 O+] n Note: The significant figures for logarithmic numbers are given after the decimal, and the numbers preceding the decimal give the exponent.

Understanding p. H and [H+] Try filling in this table in your notes: x vs. -log x [H+] p. H -1 Estimate the p. H 10 of a 3 M HCl 1 0 solution. 0. 1 1 Estimate the p. H 0. 001 3 of a 3 x 10 -5 M HCl solution. 5 10 -5 7 10 -7 14 10 -14 n

Calculating p. H of a strong acid: Ex: Given a solution of 0. 50 M HCl, what is the p. H? Step 1: Find [H 3 O+] in mol/L 0. 50 mol/L = 5. 0 x 10 -1 mol/L Step 2: Place value in equation and solve. p. H = -log[5. 0 x 10 -1] = 0. 30

Practice p. H Calculations n 1. 2. 3. n 1. 2. Find p. H of the following solutions if [H 3 O+] is (try estimating first!): 1. 00 x 10 -3 6. 59 x 10 -6 9. 47 x 10 -10 Find [H 3 O+] if the p. H is: 6. 678 3. 10. 0 2. 533 4. 2. 56

p. OH You can calculate the p. H of a solution if you know the concentration of hydronium ion. [OH-] n If we use the ion product constant of water we can derive this equation: [p. H][p. OH] = 1. 00 x 10 -14 n Working with this equation leads to: p. H + p. OH = 14 n

Calculating p. H of a strong base: Ex: Find the p. H of a solution with an [Na. OH] of 1. 0 x 10 -8. Step 1: Solve for [H 3 O+] in equation: [H 3 O+] = 10 -14 [OH-] Step 2: Place values in: [H 3 O+] = 10 -14 = 10 -6 M [1. 0 x 10 -8]

n Step 3: Solve for p. H by placing [H 3 O+] in p. H = –log[H 3 O+] p. H = -log(1. 0 x 10 -6) p. H = 6. 0

Practice p. H Calculations Using p. OH n 1. 2. 3. 4. Find the p. H of the following solutions with [OH] of: 1. 00 x 10 -4 2. 64 x 10 -13 5. 67 x 10 -2 3. 45 x 10 -11

Calculating p. H, p. OH, [H+] and [OH-] n If one of these values is known, all others can be found using the following relationships: p. H + p. OH = 14 p. H=-log[H+]=10 -p. H [H+] p. OH=-log[OH-] [H+] [OH-] = Kw [OH-]=10 -p. OH [OH-]

Calculating p. H of a weak acid: Ex: Find the p. H of a 1. 00 M HF solution (Ka=7. 2 x 10 -4): HF(aq) ↔ H+(aq) + F-(aq) Ka= [H+][F-] [HF] Use ICE box method: HF(aq) ↔ H+(aq) + F-(aq) I 1. 00 M 0 0 C -x +x +x E 1. 00 -x x x

Ka= [H+][F-] = (x)(x) ≈ x 2 [HF] (1. 00 -x) (1. 00) x 2 = Ka(1. 00) = (7. 2 x 10 -4)(1. 00) x ≈ 2. 7 x 10 -2 CHECK: 5% rule is valid since 2. 7 x 10 -2/1. 00 = 2. 7% p. H = -log[H+] = 1. 57 Try AP 2005 practice problem!!

Acid-base properties of salts n n Salt hydrolysis is the reaction in which a salt dissolved in water produces an acid or basic solution (opposite of a neutralization reaction): Ex 1: Al. Cl 3 + H 2 O → Al(OH)3 + 3 HCl weak base q strong acid A salt that contains the conjugate base of a strong acid will produce a slightly acidic solution when dissolved in water.

Strength of Acid-base pairs Strong acids yield WEAK conjugate bases in water… they have a low affinity for H+ n Weak acids yield STRONG conjugate bases n Strong bases yield WEAK conjugate acids n Weak bases yield STRONG conjugate acids n n This is important when we determine the acidic/basic properties of salts….

Acid-base properties of salts n Ex 2: Na. C 2 H 3 O 2 + H 2 O → Na. OH + HC 2 H 3 O 2 strong base weak acid q. A salt that contains the conjugate base of a weak acid will produce a slightly basic solution when dissolved in water. q What will happen when Na. Cl dissolves in water?

Relative strength of conjugate A/B pairs. Reactions that are most likely to occur will be between top left SA and bottom right SB. Which acids are stronger than water? Which bases are stronger than water?

Percent dissociation = amount dissociated x 100% initial concentration

Practice Problem: find Ka given % dissociation of a weak acid: Find [H+] first using % dissociation n Use formula Ka=x 2/[acid]0 - x to find Ka n n Make up your own problem from a % dissociation value…

The effect of structure on acid-base properties n n Relative acidity of oxyacids (hydrogen is attached to oxygen, and acid contains one other element): Acid strength increases as the O-H bond is weakened (or with increasing # of oxygen atoms on central atom). HCl. O 4>HCl. O 3>HCl. O 2>HCl. O Bond strength of H-F is too strong (F is so small) to make it a strong acid (will not dissociate easily). HI>HBr>HCl>>HF

BUFFERS A buffer solution is one which resists changes in p. H when small quantities of an acid or a base are added to it. n How do buffer solutions work? A buffer solution has to contain ions or molecules which will remove any hydrogen ions or hydroxide ions that you might add to it -otherwise the p. H will change. n buffer demo n

ex: HF and Na. F ex: NH 3 and NH 4 Cl

Calculating p. H of a buffer n Taking the log of the Ka equation (rearranged to solve for H+)… …. yields the Henderson-Hasselbalch equation: This is used to find the p. H of a buffer solution

H-H Special case… n What happens with equimolar amounts of acid and conjugate base ion? n This situation simplifies to p. H = p. Ka, since log 1=0.

Preparing buffers: 1) A buffer can be prepared by adding a common ion (the conjugate base) to a weak acid. ex: CH 3 COOH + Na. CH 3 COO 2) Alternatively, a buffer can also be prepared by partially neutralizing a weak acid with a strong base to produce the conjugate base anion. ex: CH 3 COOH + OH- → CH 3 COO- + H 2 O

Practice Problem A buffer problem is simply a weak acid equilibrium problem that involves a common ion. . n Ex: Calculate the [H+], p. H and percent dissociation of HF in a buffer solution that contains 1. 0 M HF (Ka= 7. 2 x 10 -4) and 1. 0 M Na. F. n ANS: HF ↔ H+ + FKa= 7. 2 x 10 -4 = [H+][F-] = x(1. 0 + x) ≈ x(1. 0) [HF] 1. 0 – x 1. 0 n

Solving for x = 7. 2 x 10 -4 [H+] = 7. 2 x 10 -4 and p. H = 3. 14 % dissociation is 7. 2 x 10 -4 x 100% = 0. 072% 1. 0 M

Buffer Capacity n n n The buffering capacity of a solution represents the amount of H+ or OH- the buffer can absorb without significant changes in p. H. The p. H of a buffered solution is determined by the ratio [A-]/[HA]. The capacity of a buffered solution is determined by the magnitudes of [HA] and [A-]. Therefore, a good buffer solution will have relatively high concentrations of BOTH components.

Practice Problem: Buffer Preparation n A chemist needs a solution buffered at p. H 4. 30 and can choose from the following acids and their sodium salts: a) CH 2 Cl. COOH, cloroacetic acid (Ka = 1. 35 x 10 -3) b) CH 3 CH 2 COOH, propanoic acid (Ka = 1. 3 x 10 -5) c) C 6 H 5 COOH, benzoic acid (Ka = 6. 4 x 10 -5) d) HOCl, hypochlorous acid (Ka = 3. 5 x 10 -8) Calculate the ratio [HA]/[A-] required for each system to yield a p. H of 4. 30. Which will work best? A p. H of 4. 30 corresponds to [H+] = 5. 0 x 10 -5 M, and [H+] = Ka[HA]/[A-] Using Ka values above, the [HA]/[A-] ratio for each of the salts are: a) cloroacetic acid: 3. 7 x 10 -2 b) propanoic acid: 3. 8 best since [HA]/[A-] c) benzoic acid: 0. 78 ratio is closest to 1 d) hypochlorous acid: 1. 4 x 103

Indicators

p. H and titration curves –SA and SB Adding SA to SB Adding SB to SA In what way are the curves different? Similar?

Comparing WA and SA added to SB…. What are the differences? Similarities?

Adding Acid to Base Adding WA to SB Adding SA to WB In what way are the curves different? Similar?

Adding Base to Acid Adding WB to SA Adding SB to WA

Finding Ka from titration curve

More complicated titration curves… Adding a weak acid to weak base… Adding H 2 CO 3 to strong base… usually too difficult to measure two equivalence points seen in the endpoint polyprotic acid titration

Neutralization Reaction

Acid-Base Titration

Make up a weak acid name, formula and Ka…. What will the titration curve look like?