Acids, Bases, and Salts

Acids, Bases, and Salts

Background A(g) + 2 B(g) 3 C(g) + D(g) Equilibrium constant (Keq) Keq = = [Products] [Reactants] [C]3 [D] [A][B] 2 A florence flask was getting dressed for the opera. All of a sudden she screamed: "Erlenmeyer, my joules! Somebody has stolen my joules!". The husband replied: "Take it easy honey, do not overreact. We'll find a solution". LeChateliers Principle

(lu-SHAT-el-YAYs) Acids, Bases, and Salts Ch. 19 Acids Properties Taste sour or tart Change the color of an acid-base indicator Can be strong or weak electrolytes in aqueous solution Bases Properties Q: Why do chemistry professors like to teach about ammonia? A: Because it's basic stuff. Taste bitter Feel slippery Will change the color of an acid-base indicator Can be strong or weak electrolytes in aqueous solution During class, the chemistry professor was

demonstrating the properties of various acids. Now Im going to drop this silver coin into this glass of acid. Will it dissolve? No sir, one student called out. No? queried the professor. Perhaps you can explain why the silver wont dissolve in this particular acid. Because if it would, you wouldnt have dropped it in! Acid vs. Base Alike Different Affects pH and litmus paper pH < 7 Topic sour taste react with metals

Different Acid pH > 7 Topic Related to H+ (proton) concentration pH + pOH = 14 Base bitter taste does not react with metals Properties ACIDS electrolyte s

BASES electrolytes sour taste bitter taste turn litmus red turn litmus blue react with metals to form H2 gas slippery feel vinegar, milk, soda, apples, citrus fruits ammonia, lye, antacid, baking soda ChemASAP Common Acids and Bases Strong Acids (strong electrolytes)

HCl HNO3 HClO4 H2SO4 hydrochloric acid nitric acid perchloric acid sulfuric acid Weak Acids (weak electrolytes) CH3COOH H2CO3 acetic acid carbonic Strong Bases (strong electrolytes) NaOH KOH Ca(OH)2 sodium hydroxide potassium hydroxide calcium hydroxide

Weak Base (weak electrolyte) NH3 ammonia NH3 + H2O NH4OH Common Acids Formula HF HBr HI HCl HClO HClO2 HClO3 HClO4 H2S H2SO3 H2SO4 HNO2 HNO3 H2CO3 H3PO3 H3PO4

Name of Acid Ion of Salt hydrofluoric hydrobromic hydroiodic hydrochloric hypochlorous chlorous chloric perchloric hydrosulfuric sulfurous sulfuric nitrous nitric carbonic phosphorous phosphoric Name of Negative fluoride bromide iodide chloride hypochlorite chlorite

chlorate perchlorate sulfide sulfite sulfate nitrite nitrate carbonate phosphite phosphate Common Bases Sodium hydroxide NaOH lye or caustic soda Potassium hydroxide KOH lye or caustic potash Magnesium hydroxide

Mg(OH)2 milk of magnesia Calcium hydroxide Ca(OH) 2 slaked lime Ammonia water NH3 H2O household ammonia Have you heard the one about a chemist who was reading a book about helium and just couldn't put it down? Arrhenius Acids and Bases Arrhenius said that acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution Bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solution Monoprotic acids = acids that contain 1 ionizable hydrogen like nitric acid

(HNO3) Diprotic acids = acids that contain 2 ionizable hydrogens like sulfuric acid (H2SO4) Triprotic acids = acids that contain 3 ionizable hydrogens like phosphoric acid (H3PO4) Hydroxides of group I metals are very soluble in water and caustic to skin, hydroxide of group II metals are not very soluble in water and very dilute Q: if(can both abe bear in Yosemite and one in Alaska even when saturated taken internally) fall into the water which one disolves faster? A: The one in Alaska because it is Polar. Bronsted-Lowry Acids and Bases But what about bases like sodium carbonate (Na2CO3) and ammonia (NH3)??? The Bronsted-Lowry theory defines an acid as a hydrogen-ion donor, and a base as a hydrogenion acceptor More complete definition

Conjugate Acids and Bases Conjugate acid = the particle formed when a base gains a hydrogen ion Conjugate base = the particle that remains when an acid has donated a hydrogen ion Conjugate acids and bases are always paired w/ a base or an acid, respectively Conjugate acid-base pair = consists of 2 substances related by the loss or gain of a single hydrogen ion. Cont A water molecule that gains a hydrogen ion becomes a positively charged Hydronium ion (H3O+) Amphoteric = a substance that can act as both an acid and a base EX: Water Lewis Acids and Bases Lewis proposed that an acid accepts a pair of electrons during a reaction while a base donates a pair of electrons More general than either of the other 2 theories

Lewis acid = a substance that can accept a pair of electrons to form a covalent bond Lewis base = a substance that can donate a pair of electrons to form a covalent bond Acid Base Systems Type Acid Base Arrhenius H+ or H3O + producer OH - producer BrnstedLowry Proton (H +) donor Proton (H +) acceptor

Lewis Electron-pair acceptor Electron-pair donor Many Lewis acids are also Bronsted-Lowry acids and vice versa but not all *PP 1-2, 19.1 sect. assessment #8 pg. 593 Copper leaves Copper Sulfate and says see you: he answers CuSO4!!!!! Ion Product Constant for Water Self-ionization of water = the reaction in which water molecules produce ions For aqueous solutions, the product of the hydrogen-ion concentration and the hydroxide-ion concentration equals 1.0 x 10-14 [H+] x [OH-] = 1.0 x 10-14 ion-product constant for water (KW) = the product of the concentrations of the hydrogen ions and hydroxide ions in water

Acidic solution = one in which [H+] is greater than [OH-] *The [H+] is greater than 1 x 10-7 M* Basic solution = one in which [H+] is less than [OH-] *The [H+] is less than 1 x 10-7 M* *SP 19.1, PP 9-10 pg. 596 Basic 7 Acid 14 Neutral pH Scale Acidic 0 Base [H+]

pH 10-14 14 10-13 13 10-12 12 10-11 11 10-10 10 10-9 9

10-8 8 10-7 7 10-6 6 10-5 5 10-4 4 10-3 3 10-2

2 10-1 1 100 0 1 M NaOH Ammonia (household cleaner) Blood Pure water Milk Vinegar Lemon juice Stomach acid 1 M HCl

pH of Common Substances vinegar 2.8 gastric juice 1.6 carbonated beverage 3.0 0 1 2 urine 6.0 acidic 4 5

6 bile 8.0 seawater 8.5 7 neutral [H+] = [OH-] 8 9 1.0 M NaOH (lye) 14.0 milk of magnesia 10.5

detergents 8.0 - 9.0 milk 6.4 tomato 4.2 coffee 5.0 3 blood 7.4 potato 5.8 apple juice 3.8 lemon juice 2.2 drinking water 7.2

bread 5.5 orange 3.5 1.0 M HCl 0 water (pure) 7.0 soil 5.5 ammonia 11.0 bleach 12.0 10 11 basic

12 13 14 pH of Common Substance More acidic More basic pH NaOH, 0.1 M Household bleach Household ammonia Lime water Milk of magnesia Borax Baking soda Egg white, seawater Human blood, tears Milk Saliva

Rain Black coffee Banana Tomatoes Wine Cola, vinegar Lemon juice Gastric juice 14 13 12 11 10 9 8 7 6 5 4 3 2 1 0 [H1+]

1 x 10-14 1 x 10-13 1 x 10-12 1 x 10-11 1 x 10-10 1 x 10-9 1 x 10-8 1 x 10-7 1 x 10-6 1 x 10-5 1 x 10-4 1 x 10-3 1 x 10-2 1 x 10-1 1 x 100 [OH1-] 1 x 10-0 1 x 10-1 1 x 10-2 1 x 10-3 1 x 10-4 1 x 10-5 1 x 10-6

1 x 10-7 1 x 10-8 1 x 10-9 1 x 10-10 1 x 10-11 1 x 10-12 1 x 10-13 1 x 10-14 pOH 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

Sren Sorensen (1868 - 1939) pH Concept pH = the negative logarithm of the hydrogen-ion concentration of a solution A solution in which [H+] if greater than 1 x 10-7 M has a pH less than 7.0 and is acidic. The pH of pure water or a neutral aqueous solution is 7.0. A solution with a pH greater than 7 is basic and has a [H+] of less than 1 x 10-7 M. Cont The pOH of a solution equals the negative logarithm of the hydroxide-ion concentration A solution w/ a pOH less than 7 is basic, greater than 7 is acidic For pH calculations, you should express the hydrogen-ion

concentration in scientific notation *SP 19.2, PP 11 pg. 599 *Given pH = 4.6 determine the hydronium ion *SP 19.3-19.4, PP 13-16 pg. 600601 pH Calculations pH pH = -log[H3O+] [H3O+] [H3O+] = 10-pH pH + pOH = 14 pOH [H3O+] [OH-] = 1 x10-14 pOH = -log[OH-] [OH-] [OH-] = 10-pOH

Strength of Acids and Bases Strong Acids = completely ionized in aqueous solution HCl and H2SO4 Weak Acids = ionize only slightly in aqueous solution Ethanoic acid (acetic acid) Comparison of Strong and Weak Acids Type of acid, HA Reversibility of reaction Ka value Ions existing when acid, HA, dissociates in H2O Strong Not reversible Ka value very large

H+ and A-, only. No HA present. Weak reversible Ka is small H+, A-, and HA HA(aq) + H2O(l) H3O+(aq) + A-(aq) The equilibrium expression for the reaction is Ka = [H3O+] [A-] [HA] Note: H3O+ = H+ Relative Strengths of Acids and Bases

Formula perchloric HClO4 hydrogen chloride HCl nitric HNO3 sulfuric H2SO4 hydronium ion H3O+ hydrogen sulfate ion HSO4phosphoric H3PO4 acetic HC2H3O2 carbonic H2CO3 hydrogen sulfide H2S ammonium ion NH4+ hydrogen carbonate ion HCO3water H2O ammonia

NH3 hydrogen H2 acid Conjugate base Formula perchlorate ion ClO4chloride ion Clnitrate ion NO3hydrogen sulfate ion HSO4water H2O sulfate ion SO42dihydrogen phosphate ion H2PO4acetate ion C2H3O2hydrogen carbonate ion HCO3hydro sulfide ion HSammonia NH3 carbonate ion CO32hydroxide ion OHamide ion NH2hydride ion Hconjugate base + H+

Decreasing Base Strength Decreasing Acid Strength Acid Acid Dissociation Constant For dilute solutions, the conc. of water is a constant. It can be combined w/ Keq to give the acid dissocation constant. Acid dissociation constant = the ratio of the concentration of the dissociated (or ionized) form of an acid to the concentration of the undissociated (nonionized) form. Weak acids have small Ka values. The stronger an acid is, the larger is its Ka value. Nitrous acid (HNO2) has a Ka of 4.4 x 10-4, acetic acid has a Ka of 1.8 x 10-5 so nitrous acid is more ionized and has a higher [H 3O+] or [H+] thus is a stronger acid Di and triprotic acids lose each H separately so they have multiple dissociation constants Equilibrium and pH Calculations Weak acid HA + H2O Strong acid

HA H + + AHA + H2O H3O+ + A- H3O+ + Aacid-dissociation constant calculations [A-] [H3O+] Ka = [H3O+] [HA] + antilog(-pH) 7 1 x 10-14 = [OH-] [OH-]

-log [H3O+] pH 0 [HA] = [H3O+] 14 1 x 10-14 = [H3O+][OH-] Kw = [H3O+][OH-] 1 x 10-14 [H3O+] = - Base Dissociation Constant Strong bases = dissociate completely into metal ions and hydroxide ions in aqueous solution Ex: Ca(OH)2 Weak bases = react w/ water to form the

hydroxide ion and the conjugate acid of the base Ex: ammonia NH3 Base dissociation constant (Kb) = the ratio of the concentration of the conjugate acid times the conc. of the hydroxide ion to the conc. of the base Calculating Dissociation Constants To find the Ka of a weak acid or the Kb of a weak base, substitute the measured concentration of all the substances present at equilibrium into the expression for Ka or Kb. *SP 19.5, PP 22-23 pg. 610 Weak Acids (pKa) Weak Acids dissociate incompletely (~20%) Strong Acids dissociate completely (~100%) A(g) + 2 B(g) 3 C(g) + D(g) Equilibrium constant (Keq) = Keq = [Products]

[Reactants] [C] 3[D] [A][B] 2 LeChateliers Principle (lu-SHAT-el-YAYs) Values of Ka for Some Common Monoprotic Acids HSO4HClO2 HC2H2ClO2 HF HNO2 HC2H3O2 HOCl HCN NH4+ HOC6H5 Name Value of Ka* hydrogen sulfate ion 1.2 x 10-2

chlorous acid 1.2 x 10-2 monochloracetic acid 1.35 x 10-3 hydrofluoric acid 7.2 x 10-4 nitrous acid 4.0 x 10-4 acetic acid 1.8 x 10-5 hypochlorous acid 3.5 x 10-8 hydrocyanic acid 6.2 x 10-10 ammonium ion 5.6 x 10-10 phenol 1.6 x 10-10 *The units of Ka are mol/L but are customarily omitted. Increasing acid strength Formula Sample 1) One gram of concentrated sulfuric acid (H2SO4) is diluted to a 1.0 dm3 volume with water. What is the molar concentration of the hydrogen ion in this solution?

What is the pH? Solution) First determine the number of moles of H2SO4 x mol H2SO4 = 1 g H2SO4 H2SO4 1 mol H2SO4 98 g H2SO4 H+ + HSO41- & = 0.010 mol H2SO4 HSO41- H+ + SO42- OVERALL: H2SO4 0.010 M

2 H+ + SO42- in dilute solutions...occurs ~100% 0.020 M pH = - log [H+] substitute into equation pH = - log [0.020 M] pH = 1.69 A volume of 5.71 cm3 of pure acetic acid, HC2H3O2, is diluted with water at 25 oC to form a solution with a volume of 1.0 dm3. What is the molar concentration of the hydrogen ion, H+, in this solution? (The density of pure acetic acid is 1.05 g/cm3.) Step 1) Find the mass of the acid Mass of acid = density of acid x volume of acid = 1.05 g/cm3 x 5.71 cm3 = 6.00 g Step 2) Find the number of moles of acid. (From the formula of acetic acid, you can calculate that the molar mass of acetic acid is 60 g / mol). x mol acetic acid = 6.00 g HC2H3O2 Molarity: M = mol / L Substitute into equation

1 mol HC2H3O2 = 0.10 mol acetic acid (in 1 L) 60 g HC2H3O2 M = 0.10 mol / 1 L M = 0.1 molar HC2H3O2 Step 3) Find the [H+] Ka = HC2H3O2 0.1 M Step 3) Find the [H ] + weak acid H+ + C2H3O21- ? 0.1 M Ka = 1.8 x 10-5 @ 25 oC for acetic acid 1

[H ][C2H3O2 ] Ka = [HC2H3O2 ] 1 1.8 x 10 -5 [H ][C2H3O2 ] = [HC H O ] 2 3 2 How do the concentrations of H+ and C2H3O21- compare? [x][x] [HC2H3O2 ] Substitute into equation: 1.8 x 10 -5 1.8 x 10 -5

x2 [0.10 M] pH = - log[H+] x2 = 1.8 x 10-6 M x = 1.3 x 10-3 molar pH = - log [1.3 x10-3 M] = [H+] pH = 2.9 Practice Problems: 1a) What is the molar hydrogen ion concentration in a 2.00 dm3 solution of hydrogen chloride in which 3.65 g of HCl is dissolved? 1b) pH 2a) What is the molar concentration of hydrogen ions in a solution containing 3.20 g of HNO3 in 250 cm3 of solution? 2b) pH 3a) An acetic acid solution is 0.25 M. What is its molar concentration of hydrogen ions?

3b) pH 4) A solution of acetic acid contains 12.0 g of HC2H3O2 in 500 cm3 of solution. What is the molar concentration of hydrogen ions? 1a) 0.0500 M 1b) pH = 1.3 2a) 0.203 M 2b) pH = 0.7 3a) 2.1 x 10-3 M 3b) pH = 2.7 4) 2.7 x 10-3 M Weak Acids Cyanic acid is a weak monoprotic acid. If the pH of 0.150 M cyanic acid is 2.32. calculate Ka for cyanic acid. + H3O (aq) 4.8 x 10-3 M 0.150 M Ka =

[Products] Ka = [Reactants] [4.8 x 10-3 M] Ka = + CN1-(aq) H+(aq) HCN(aq) [4.8 x 10-3 M] [CN1-] [0.150 M] Ka = 1.54 x 10-4 4.8 x 10-3 M [H3O+] [CN1-] [HCN] pH = -log[H3O+]

10-pH = [H3O+] 10-2.32 = [H3O+] 4.8 x10-3 M = [H3O+] Titration Q: How did the chemist survive the famine? A: By subsisting on titrations. Neutralization Reaction = a reaction b/w an acid and a base in aqueous solution to produce salt and water Equivalence point = when the # of moles of hydrogen ions equals the number of moles of hydroxide ions Titration = the process of adding a known amt. of solution of known conc. to determine the conc. of another solution Standard solution = the solution of known conc. End point = the point at which the indicator changes color The point of neutralization is the end pt. of the titration *SP 19.7, PP 32-33 pg. 616 Q: How do you get lean molecules? A: Feed them titrations. Q: What did one titration say to the other? A: Let's meet at the endpoint!

Buffers Buffer = a solution in which the pH remains relatively constant when small amts. of acid or base are added A buffer is a solution of weak acids and one of its salts or a solution of a weak base and one of its salts A buffer solution is better able to resist drastic changes in pH than is pure water Buffer capacity = the amt. of acid or base that can be added to a buffer solution before a significant change in pH occurs Naming Acids Anion Acid _________ ide (chloride, Cl1-) add H+ _________ ate (chlorate, ClO3-) (perchlorate, ClO4-) add H+

_________ite (chlorite, ClO2-) (hypochlorite, ClO-) add H+ ions ions ions Hydro____ ic acid (hydrochloric acid, HCl) _________ic acid (chloric acid, HClO3) (perchloric acid, HClO4) ______ous acid (chlorous acid, HClO2) (hypochlorous acid, HClO) Review A physicist, biologist and a chemist were going

to the ocean for the first time. What does one do with a dead body? Barium in a krypt-on Maybe he was killed oxydentally. They should have seen the doctor first, he'd Curium. Ah, barium anyway, just to see how he reacts. better though to have helium. Perhaps with a houseplant, a Germanium. And if they stole it, the police would Cesium. Locked up for life, in Irons. They would go crazy in jail, a Silicon. The physicist saw the ocean and was fascinated by the waves. He said he wanted to do some research on the fluid dynamics of the waves and walked into the ocean. Obviously he was drowned and never returned. The biologist said he wanted to do research on the flora and fauna inside the ocean and walked inside the ocean. He too, never returned. The chemist waited for a long time and afterwards, wrote the observation, "The physicist and the biologist are soluble in ocean water".

Soren Sorenson developed pH scale 7 neutral pH scale 0 [H+] = [OH-] acid 14 base (alkalinity) Arnold Beckman invented the pH meter pH = -log [H+] pOH = -log [OH-] pH + pOH = 14 kW = [H ] [OH ] + -

kw = 1 x 10-14 H+ + H2O proton H 3O+ hydronium ion Strong / Weak Acid Strong HA H+ + A- (~100% dissociation) Weak HA H+

+ A- (~20% dissociation) H2A 2 H+ [Product] Ka = [Reactant] + Ka = A[H+]2 [A-] [H2A] acid dissociation constant Ka 0.8 0.0021

H3PO4 HF 3H+ + PO43H + + F- Acid + Base Salt + Water How would you make calcium sulfate in the lab? H2SO4 + Ca(OH)2 ?

? BASE ACID Sour taste, litmus CaSO4 + 2 H2O red bitter taste, litmus blue Arrhenius H+ as only ion in water Arrhenius OH- as only ion in water Brnsted-Lowry proton donor Brnsted-Lowry proton acceptor

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