by Steven S. Zumdahl & Donald J. DeCoste University of Illinois Introductory Chemistry: A Foundation, 6th Ed. Introductory Chemistry, 6th Ed. Basic Chemistry, 6th Ed. by Steven S. Zumdahl & Donald J. DeCoste University of Illinois

Chapter 16 Acids and Bases

CaCO3 + 2 HCl CaCl2 + CO2 + H2O Properties of Acids Sour taste Turn blue litmus paper red Change color of vegetable dyes (red cabbage juice) React with “active” metals Like Al, Zn, Fe, but not Cu, Ag or Au Zn + 2 HCl ZnCl2 + H2 Corrosive React with carbonates, producing CO2 Marble, baking soda, chalk CaCO3 + 2 HCl CaCl2 + CO2 + H2O React with bases to form ionic salts, and often water 2

Properties of Bases Also known as alkalis Bitter Taste Feel slippery Change color of vegetable dyes Different color than acid Turn red litmus blue React with acids to form ionic salts, and often water Neutralization 3

Arrhenius Theory Acids ionize in water to H+ ions and anions Bases ionize in water to OH- ions and cations Neutralization reaction involves H+ combining with OH- to make water H+ ions are protons 4

Arrhenius Theory (cont.) Definition only good in water solution Definition does not explain why ammonia solutions turn litmus blue Basic without OH- ions

Brønsted-Lowry Theory H+ transfer reaction Since H+ is a proton, also known as proton transfer reactions Acids are proton donors, bases are proton acceptors In the reaction, a proton from the acid molecule is transferred to the base molecule Products are called the conjugate acid and conjugate base 5

Brønsted-Lowry Theory (cont.)

Brønsted-Lowry Theory (cont.) H-A + :B  A- + H-B+ A- is the conjugate base, H-B+ is the conjugate acid Conjugate acid-base pair is either the original acid and its conjugate base or the original base and its conjugate acid H-A and A- are a conjugate acid-base pair :B and H-B+ are a conjugate acid-base pair 6

Example #1: Write the conjugate base for the acid H3PO4 Determine what species you will get if you remove 1 H+1 from the acid. Conjugate base will have one more negative charge than the original acid H3PO4  H+ + H2PO4- 7

Self- check p 490 Which of the following represent conjugate acid base pairs? A. H2O, H3O+ B. OH-, HNO3 C. H2SO4, SO42- D. HC2H3O2, C2H3O2-

Brønsted-Lowery Theory (cont.) In this theory, instead of the acid, HA, dissociating into H+(aq) and A- (aq), the acid donates its H to a water molecule HA + H2O  A- + H3O+ A-1 is the conjugate base, H3O+ is the conjugate acid 8

Brønsted-Lowry Theory (cont.) H3O+ is called the hydronium ion In this theory, substances that do not have OH- ions can act as a base if they can accept a H+1 from water. H2O + :B  OH- + H-B+ :B is acting here as a base.

Strength of Acids & Bases The stronger the acid, the more willing it is to donate H+ 9

Strength of Acids & Bases (cont.) Strong bases will react completely with water to form hydroxide: CO3-2 + H2O HCO3- + OH- Only small fraction of weak base molecules pull H+ off water: HCO3- + H2O H2CO3 + OH-

Multiprotic Acids Monoprotic acids have 1 acid H, diprotic 2, etc. In oxyacids only the H on the O is acidic In strong multiprotic acids, like H2SO4, only the first H is strong; transferring the second H is usually weak H2SO4 + H2O  H3O+ + HSO4- HSO4- + H2O  H3O+ + SO4-2 10

Water As an Acid and a Base Amphoteric substances can act as either an acid or a base. Water as an acid, NH3 + H2O  NH4+ + OH- Water as a base, HCl + H2O  H3O+ + Cl- Water can even react with itself: H2O + H2O  H3O + + OH- 11

Autoionization of Water Water is an extremely weak electrolyte. Therefore there must be a few ions present H2O + H2O  H3O+ + OH- 12

Acid Nomenclature Acids Examples: Compounds that form H+ in water. Formulas usually begin with ‘H’. In order to be an acid instead of a gas, binary acids must be aqueous (dissolved in water) Ternary acids are ALL aqueous Examples: HCl (aq) – hydrochloric acid HNO3 – nitric acid H2SO4 – sulfuric acid

Acid nomenclature If anion ending is –ide (Binary compound), the acid name is hydro(stem)ic acid If ternary compounds- -Ate ending: (stem)ic acid -ite ending: (stem)ous acid

Acid Nomenclature Flowchart

Solved examples HBr – 2 elements-ide, hydrobromic acid H2CO3- 3 elements- ate, carbonic acid H2SO3- 3 elements- ite, sulfurous acid Hydrofluoric acid: 2 elements= HF Sulfuric acid: 3 elements, –ic= -ate, H2SO4 Nitrous acid: 3 elements, -ous= -ite, HNO2

Now your turn! HI (aq) HCl H2SO3 HNO3 HIO4

Hydrobromic acid Nitrous acid Carbonic acid Phosphoric acid

Acidic and Basic Solutions Acidic solutions have a larger [H+] than [OH-] Basic solutions have a larger [OH-] than [H+] Neutral solutions have [H+]=[OH-]= 1 x 10-7 M [H+] = 1 x 10-14 [OH-] [OH-] = [H+] 13

Ion product of water

Example #2 Determine the [H+] and [OH-] in a 10.0 M H+ solution 14

Example #2 (cont.) Determine the given information and the information you need to find Given [H+] = 10.0 M, find [OH-]

Example #2 (cont.) Given [H+] = 10.0 M = 1.00 x 101 M Kw = 1.0 x 10-14 15

Self check p 497 Calculate [H+] in a solution in which [OH-] = 2.0X 10-2 M. Is this solution acidic, neutral or basic?

pH & pOH The acidity/basicity of a solution is often expressed as pH or pOH. pH = -log[H3O+] pOH = -log[OH-] pHwater = -log[10-7] = 7 = pOHwater [H+] = 10-pH [OH-] = 10-pOH 16

pH scales

pH & pOH (cont.) pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral The lower the pH, the more acidic the solution; the higher the pH, the more basic the solution 1 pH unit corresponds to a factor of 10 difference in acidity pOH = 14 - pH

pH of Common Substances

Example #3 Calculate the pH of a solution with a [OH-] = 1.0 x 10-6 M 17

Example #3 (cont.) Find the concentration of [H+]

Example #3 (cont.) Enter the [H+] concentration into your calculator and press the log key log(1.0 x 10-8) = -8.0 Change the sign to get the pH pH = -(-8.0) = 8.0 18

Example #4 Calculate the pH and pOH of a solution with a [OH-] = 1.0 x 10-3 M Enter the [H+] or [OH-] concentration into your calculator and press the log key log(1.0 x 10-3) = -3.0 Change the sign to get the pOH pOH = -(-3) = 3.0 Subtract the calculated pH or pOH from 14.00 to get the other value pH = 14.00 – 3.0 = 11.0 19

Solving concentration from pH or pOH Calculate the [OH-] of a solution with a pH of 7.41 If you want to calculate [OH-] use pOH; if you want [H+] use pH. It may be necessary to convert one to the other using 14 = [H+] + [OH-] pOH = 14.00 – 7.41 = 6.59 20

Example #5 (cont.) Enter the pH or pOH concentration into your calculator Change the sign of the pH or pOH -pOH = -(6.59) Press the button(s) on you calculator to take the inverse log or 10x [OH-] = 10-6.59 = 2.6 x 10-7 M

Self check p 499 Calculate the pH value for each of the following solutions at 25°C. A. a solution in which [H+] =1.0 X 10-9M. B. a solution in which [OH-] =1.0 X 10-6M. P501- A sample of rain in an area with severe air pollution has a pH of 3.5. What is the pOH of this rainwater?

Calculating the pH of a Strong, Monoprotic Acid A strong acid will dissociate 100% HA  H+ + A- Therefore the molarity of H+ ions will be the same as the molarity of the acid Once the H+ molarity is determined, the pH can be determined pH = -log[H+] 21

Example #6 Calculate the pH of a 0.10 M HNO3 solution. 22

Example #6 (cont.) Determine the [H+] from the acid concentration HNO3  H+ + NO3- 0.10 M HNO3 = 0.10 M H+ Enter the [H+] concentration into your calculator and press the log key log(0.10) = -1.00 Change the sign to get the pH pH = -(-1.00) = 1.00

Self check p503 The pH of rainwater in a polluted area was measured to be 3.5. What is the [H+] in this rainwater? The pOH of a liquid drain cleaner was found to be 10.50. What is the [OH-] for this cleaner? P505- Calculate the pH of a solution of 5.0 X 10-3 M HCl.

Buffered Solutions Buffered solutions resist change in pH when an acid or base is added to it. Used when need to maintain a certain pH in the system Blood 23

Buffered Solutions (cont.) A buffer solution contains a weak acid and its conjugate base. Buffers work by reacting with added H+ or OH-ions so they do not accumulate and change the pH. Buffers will only work as long as there are sufficient weak acid and conjugate base molecules present.

Buffered Solutions (cont.)