Chapter 14: Acids and Bases

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1 Ch 14 Page 1 Chapter 14: Acids and Bases Properties of Acids Sour taste React with some metals Turns blue litmus paper red React with bases Some Common Acids HCl, hydrochloric acid H 2 SO 4, sulfuric acid HNO 3, nitric acid HC 2 H 3 O 2, acetic acid (common name) What is the IUPAC name of this carboxylic acid? Properties of Bases Taste bitter Solutions feel slippery Turns red litmus paper blue React with acids Arrhenius Acid-a hydrogen containing compound that ionizes to produce hydrogen ions (H + ) when dissolved in water Arrhenius Acid-Base Theory HCl (aq) H + (aq)+ Cl (aq) Some Common Bases NaOH, sodium hydroxide KOH, potassium hydroxide NaHCO 3, sodium bicarbonate Because molecular acids are not made of ions, they cannot dissociate. They must be pulled apart, or ionized, by the water.

2 Ch 14 Page 2 Arrhenius Acids and Bases Hydrogen ions released from acids do not float about freely in solution. Instead, they attach to water molecules to form HYDRONIUM IONS = H 3 O + In the formula for an acid, the hydrogen atom that ionizesis many times written first- ex. HC 2 H 3 O 2 This hydrogen atom is the one that ionizes These hydrogen atoms do not ionize in water Arrhenius Base-a hydroxide containing compound that dissociates to produce OH - when dissolved in water NaOH(aq) Na + (aq) + OH (aq) Bases that contain OH - are ionic compounds. Arrhenius Acid Base Reactions Ionic substances dissociate in water. The H + from the acidcombines with the OH - from the baseto make a molecule of H 2 O HCl(aq) + NaOH(aq) NaCl(aq) + H 2 O(l) (acid + base salt + water) The cation from the base combines with the anion from the acid to make a salt (an exchange reaction)

3 Ch 14 Page 3 Bronsted-Lowry Acid-Base Theory Although widely used, the Arrhenius acid-base theory has limitations. (For example, some compounds act as bases even though they do not contain OH - ). Bronsted-Lowry theory broader definition than Arrhenius theory Bronsted-Lowry Acid- any substance that can donate a proton (H+) Bronsted-Lowry Base-any substance that can accept a H + A Brønsted-Lowry acid base reactionis any reaction in which an H + is transferred. The acid molecule donates an H + to the base molecule: H A + :B :A + H B + Base structure must contain an atom with an unshared pair of electrons to bond to H +. Examples: HCl (aq) + H 2 O (l) Bronsted-Lowry Acid-Base reactions don't have to take place in aqueous solution. Cl (aq) + H 3 O + (aq) Proton donor Proton acceptor HCl (aq) + NH 3(aq) NH 4 + (aq) + Cl - (aq) Proton donor Proton acceptor

4 Ch 14 Page 4 Conjugate Acid-Base Pairs Identify the Bronsted-Lowry acid and base in the following reactions: NH 3(aq) + H 2 O (l) NH 4 + (aq) + OH - (aq) HCO 3 - (aq) + H 2 O (l) CO 3 2- (aq) + H 3 O + (aq) Conjugate Acids and Bases A conjugate acid-base pair is two species that differ from each other by one proton HCl (aq) + H 2 O (l) A = acid; CB = conj. base B = base; CA = conj. acid Cl (aq) + H 3 O + (aq) A B CB CA NH 2 - (aq) + H 2 O (l) NH 3(aq) + OH - (aq) B A CA CB Identify the conjugate acid-base pairs in the following reactions: HF (aq) + H 2 O (l) F (aq) + H 3 O + (aq) NH 3(aq) + H 2 O (l) NH 4 + (aq) + OH - (aq) We will not cover Section 14.5 "Reactions of Acids and Bases" (we already discussed part of this in Ch.7) or Section 14.6 "Acid-Base Titration" Note: Water is an amphoteric molecule - it can function as an acid or a base

5 Ch 14 Page 5 Strong and Weak Acids and Bases A strong acidis an acid that transfers 100%(or nearly 100%) of its acidic hydrogens to water HCl + H 2 O H 3 O + + Cl - 100% transfer Hydrochloric acid: a strong acid Common Strong Acids HCl, HBr, HI, HNO 3, H 2 SO 4, HClO 4 A weak acidis an acid that transfers only a small amount of its acidic hydrogens to water HC 2 H 3 O 2 + H 2 O H 3 O + + C 2 H 3 O 2 - Acetic acid: a weak acid Common Weak Acids HC 2 H 3 O 2, HF, H 2 CO 3, H 3 PO 4, H 2 SO 3 When acetic acid is dissolved in water only 0.42%of its acidic protons are transferred to H 2 O % of the molecules remain un-ionized. Basescan also be strongor weak Common Strong Bases LiOH, NaOH, KOH, Ca(OH) 2 NaOH Na + + OH - Dissociates 100% 100% dissociation NH 3 + H 2 O NH 4+ + OH - Only small amount of dissociation Common Weak Bases NH 3, HCO 3-, CH 3 NH 2 (ammonia, bicarbonate, methyl amine) When NH 3 reacts with H 2 O only a small amount of OH - produced

6 Ch 14 Page 6 Self Ionization of Water A VERY SMALL PERCENTAGE(about 1 out of every 10 million) of water moleculescan dissociate into H 3 O + and OH - H 2 O H 2 O + H 2 O H + + OH - (Arrhenius) H 3 O + + OH - (Bronsted-Lowry) All aqueous solutions contain both H 3 O + and OH At 25 o C, there are equal numbers of hydronium and hydroxide ions: [H 3 O + ] = [OH ] = 1 x 10-7 M at 25 C in pure water Note: [brackets] mean molar concentration in H 2 O [H 3 O + ] x [OH ] = Ion Product for Water (K w ) [H 3 O + ] x [OH ] = 1 x M The product of the H 3 O + and OH concentrations is constant:1 x M If H 3 O + ions are addedto water, there will be adecrease in [OH - ] If OH - ions are addedto water, there will be a decrease in [H 3 O + ] Sufficient acid is added so that the [H 3 O + ] is now 6.52 x 10-4 M. What is the [OH - ] in this solution? Sufficient base is added so that the [OH - ] reaches 5.41 x 10-6 M. What is the new [H 3 O + ]?

7 Ch 14 Page 7 The ph Scale Acidic Solutions: [H 3 O + ] > [OH ] Basic Solutions: [H 3 O + ] < [OH ] Neutral Solutions: [H 3 O + ] = [OH ] A strongly acidicsolution has a relatively high concentration of. A strongly basicsolution has a relatively high concentration of. The ph scaleis a 0-14 scale which measures the acidity or basicity of solution. The ph scale was derived because very small concentrations(such as 4.22 x M and 3.12 x M) are difficult to compare. ph = -log [H 3 O + ] ph = 2.0 means [H 3 O + ] = 1 x 10-2 M ph = 3.0 means [H 3 O + ] = 1 x 10-3 M If [H 3 O + ] = 1 x 10-9 M, then ph = The ph of a solution is the negative logarithm of the hydronium ion concentration Which is more acidic, ph = 2.0 or ph = 3.0?

8 Ch 14 Page 8 ph Calculations ph = -log [H 3 O + ] ph < 7acidicsolution ph > 7basicsolution ph = 7neutral solution The lower the ph, the more acidicthe solution; the higher the ph, the more basic the solution. 1 ph unit corresponds to a factor of 10 difference in acidity. Normal rangeis 0 to 14 ph 0 then [H + ] = 1 M, ph 14 then [OH ] = 1 M; but ph can be negative (very acidic) or larger than 14 (very alkaline) If [H 3 O + ] = 1 x 10-6 M, then ph = -log 10-6 = If [H 3 O + ] = 1.0 x M, ph = acidic or basic? If [H 3 O + ] = 3.6 x M, ph = acidic or basic? Text calculator: (-) LOG 3.6 E (-) 12 = Numerical calculator: 3.6 E 12 + /- LOG + /- (in ph values, the sig figs are after the decimal point)

9 Ch 14 Page 9 ph and H 3 O + Calculations Calculate the ph of the following solutions: a) [H 3 O + ] = 5.9 x b) [H 3 O + ] = 2.5 x 10-3 c) [H 3 O + ] = 6.32 x 10-6 Conversions between ph and [H ] If ph = 8.0, [H 3 O + ] = 1 x 10-8 M [H 3 O + ] = 10 -ph If ph = 2.87, [H 3 O + ] = = text calculator: 10 x (-) 2.87 = numerical calculator: /- 10 x If ph = 6.43, [H 3 O + ] =10 = (Use MODE or SCI if your calculator gives you ) Calculate the [H ] of the following solutions: a) ph = 8.92 b) ph = 2.664

10 Ch 14 Page 10 [OH - ] and poh Calculations If [OH - ] = 1 x 10-4 M, poh = 4 poh = -log [OH - ] [OH - ] = 10 -poh poh and ph are related ph + poh = 14 Calculate the ph of the following solutions: a) [OH - ] = 1.0 x 10-5 b) [OH - ] = 6.7 x 10-8 c) [OH - ] = 6.32 x 10-6 These equations will be given on the final exam: ph = -log[h 3 O + ] [H 3 O + ] = 10 -ph ph + poh = 14 poh = -log[oh - ] [OH - ] = 10 -poh We will not cover Sections "Buffers" or "Acid Rain" (this material will not be on the final exam)

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