Properties of acids

• Taste Sour (kids, don’t try this at home).
• Conduct electricity.
• Some are strong, some are weak electrolytes.
• React with metals to form hydrogen gas.
• Change indicators (litmus red).
• React with hydroxides to form water and a salt.

Properties of bases

• React with acids to form water and a salt.
• Taste bitter.
• Feel slippery (Don’t try this either).
• Can be strong or weak electrolytes.
• Change indicators (litmus blue).

Types of Acids and Bases

Several Definitions

Arrhenius Definition

• Acids produce hydrogen ions in aqueous solution.
• Bases produce hydroxide ions when dissolved in water.
• Limits to aqueous solutions.
• Only one kind of base.
• NH₃ ammonia could not be an Arrhenius base.

Bronsted-Lowry Definitions

• An acid is a proton (H⁺) donor and a base is a proton acceptor.
• Acids and bases always come in pairs.
• HCl is an acid.
• When it dissolves in water it gives its proton to water.
• HCl(g) + H₂O(l) H₃O⁺ + Cl^-
• Water is a base -makes hydronium ion.

Polyprotic Acids

• Some compounds have more than 1 ionizable hydrogen.
• HNO₃ nitric acid - monoprotic
• H₂SO₄ sulfuric acid - diprotic - 2 H⁺
• H₃PO₄ phosphoric acid - triprotic - 3 H⁺

Come in Pairs

• General equation
• HA(aq) + H₂O(l) H₃O⁺(aq) + A^-(aq)
• Acid + Base Conjugate acid + Conjugate base
• This is an equilibrium.
• B(aq) + H₂O(l) BH⁺(aq) + OH^-(aq)
• Base + Acid Conjugate acid + Conjugate base
• NH₃(aq)+H₂O(l) NH₄⁺(aq)+OH^-(aq)

Water

• Water ionizes- falls apart into ions.
• H₂O → H⁺ + OH^-
• Called the self ionization of water.
• Only a small amount.
• [H⁺ ] = [OH^-] = 1 x 10^-7M
• A neutral solution.
• In water Kw = [H⁺ ] x [OH^-] = 1 x 10^-14
• Kw is called the ion product constant.

Ion Product Constant

• H₂O H⁺ + OH^-
• Kw is constant in every aqueous. solution [H⁺] x [OH^-] = 1 x 10^-14M²
• If [H⁺] > 10^-7 then [OH^-] < 10^-7
• If [H⁺] < 10^-7 then [OH^-] > 10^-7
• If we know one, we can determine the other.
• If [H⁺] > 10^-7 acidic [OH^-] < 10^-7
• If [H⁺] < 10^-7 basic [OH^-] > 10^-7

Logarithms

• Powers of ten.
• A shorthand for big, or small numbers.
• pH = -log[H⁺]
• in neutral pH = - log(1 x 10^-7) = 7
• in acidic solution [H⁺] > 10^-7
• pH < -log(10^-7)
• pH < 7
• in base pH > 7

pH and pOH

• pOH = - log [OH^-]
• [H⁺] x [OH^-] = 1 x 10^-14M²
• pH+pOH = 14

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Practice problems

• Calculate the pH, pOH, [H+ ], [OH- ] of the following solutions and tell whether its acidic or basic.
• a) [H+ ] = 1.36 x 10 ‐9 M
• b) [OH ] = 3.86 × 10 ‐8 M
• c) pOH = 7.50

• Give me all five if you start with 5.0 grams of hydrogen sulfate and mix it in 5000 mL of water.

0

1

3

5

7

9

11

13

14

Basic

Acidic

Neutral

[OH^-]

pH

[H⁺]

pOH

How Strong

Strength

• Strong acids and bases are strong electrolytes
• They fall apart completely.
• Weak acids don’t completely ionize.
• Concentrated- much dissolved.
• Strong forms many ions when dissolved.
• Mg(OH)₂ is a strong base- it falls completely apart when dissolved.
• Not much dissolves.

Measuring strength

• Ionization is reversible.
• HA H⁺ + A^-
• makes an equilibrium.
• Equilibrium constant for an acid(acid dissociation constant.)
• K_a = [H⁺ ][A^- ] [HA]
• Stronger acid- more products.
• larger K_a (pg 450)

What about bases?

• Strong bases dissociate completely.
• B + H₂O BH⁺ + OH^-
• Base dissociation constant.
• K_b = [BH⁺ ][OH^-] [B]
• We can ignore the water because it’s concentration doesn’t change.
• Stronger base more dissociated.
• Larger K_b.

pH/pOH prac problems

• For all the following give me all five!!
• A) 50 grams of Barium Hydroxide in .5 L of water.
• B) 200 grams of hydrocyanadic acid(K_a = 4.9 x 10^-10) in 500 mL of water.
• C) .002 mol of Nitrous acid in 100 L of water. (Ka is 4.5 x 10^-4 )
• D) .002 mole of Nitric Acid in .5 L of water.

Usually happen in solutions

• If it takes 87 mL of an HCl solution to neutralize 0.67 moles of Mg(OH)₂ what is the concentration of the HCl solution?
• If it takes 58 mL of an H₂SO₄ solution to neutralize 0.34 moles of NaOH what is the concentration of the H₂SO₄ solution?

Neutralization reactions

Neutralization Reactions

• Acid + Base → Salt + water
• Salt = an ionic compound
• Water = HOH
• HNO₃ + KOH →
• HCl + Mg(OH)₂ →
• H₂SO₄ + NaOH →
• Really just double replacement.

Net Ionic Equations

• Use the double replacement reaction method and see what partners are made.
• Get rid of spectator ions
• What you have left is the net ionic equation.
• What is the net ionic equation for an acid-base reaction between solutions of Ca(OH)2 and H2S?
• H2S (aq) + Ca (OH)2 (aq)→ CaS(aq) + H2O(l)
• H+ + S-2 + Ca +2 + OH -1→ Ca + 2 + S -2 + H2O
• S-2 and Ca+2 Don’t change on Either side of the equation, they are spectators.
• Net= H+ + OH- → H2O

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Practice!

a. HCl (aq) + Ba(OH)2 (aq) →

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b. H2SO4 (aq) + NaOH (aq) →

Dilution

Adding solvent to a solution

Dilution

• The number of moles of solute doesn’t change if you add more solvent.
• The moles before = the moles after
• M₁ x V₁ = M₂ x V₂
• M is concentration and V is volume.
• Stock solutions are pre-made to known M

Practice

• 2.0 L of a 0.88 M solution are diluted to 3.8 L. What is the new molarity?

Quiz Add-On

• You have 150 mL of 6.0 M HCl. What volume of 1.3 M HCl can you make? How much water did you have to add to make that solution?

Practice

• Write the ionization for HNO₃ . If you start with .10 moles of Nitric Acid in .067 L of water, what are the five?

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Acid- Base Reactions Happen in Moles

• How many moles of HNO₃ are need to neutralize 0.86 moles of KOH?
• How many moles of HCl are needed to neutralize 3.5 moles of Mg(OH)₂ ?

Titration

Determining an unknown

Titration

• When you add the same number of moles of acid and base, the solution is neutral.
• By measuring the amount of a base added you can determine the concentration of the acid.
• If you know the concentration of the base.
• This is a titration.

Titration equations

• M_a x V_a x # of H⁺ = M_b x V_b x # of OH^-
• really moles of H⁺= moles of OH^-

Normality: Used in Acids and bases so you can take the number of H (acids) or the number of OH (bases) into account.

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N = M x # H (acids) N = M x #OH (base)

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stoichiometry with solutions as well.

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Calculate the M, m and N for a solution that contains 100 g of phosphoric acid in 500 mL of water.

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100/ 98 = 1.02 mol 500/1000 = .5 L

M = 1.02/.5 = 2.04 M m = 1.02 mol/ .5 kg of water = 2.04 m

N = M x 3 = 2.04 x 3= 6.12 N

Give me all 5 for each of the following!

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1. .10 M HCl
2. 0.0044 M Ca(OH)2
3. .0057 M Acetic Acid (K = 1.8 x 10 -5)

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More practice titrations

More Practice

• A-If it takes 45 mL of a 1.0 M NaOH solution to neutralize 57 mL of HCl, what is the concentration of the HCl ?
• B-If it takes 67 mL of 0.500 M H₂SO₄ to neutralize 15mL of Al(OH)₃ what was the concentration of the Al(OH)₃ ?
• C-How much of a 0.275 M Hydrochloric Acid will be needed to neutralize 25mL of .154 M Strontium Hydroxide?

Last Practice problems!

• 1) If you have a of a .025 M solution of Sulfuric acid, what is the pOH? pH? [H+] ? and [OH-]?
• 2) Give me all 5
• 0.0044 M Ca(OH)2
• .0057 M Acetic Acid (K = 1.8 x 10 -5)
• 3) What is the Normality of a 25 mL solution of .34 g of phosphorous acid?
• 4) It takes 83 mL of a 0.45 M sodium hydroxide solution to neutralize 235 mL of an hydrochloric acid solution. What is the concentration of the acidic solution?
• 5) It takes 38 mL of 0.75 M Gallium Hydroxide solution to completely neutralize 155 mL of a sulfuric acid solution . What is the concentration of the acidic solution?

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Solubility

• Dissolving stuff is an equilibrium
• CaCl₂(s) Ca⁺² (aq) + 2 Cl^- (aq)
• You can write and equilibrium constant for dissolving
• K_eq = [Ca⁺²] [Cl^-]²� [CaCl₂]
• The concentration of a solid does not change, so we can combine it with K_eq
• K_sp = [Ca⁺²] [Cl^-]²

Solubility

• K_sp is called the solubility product constant
• The more soluble a solid is the greater K_sp
• Used for slightly soluble salts.
• Can tell if a precipitate will form.
• Do the math
• If the answer is bigger than the Ksp it will form
• If not, it will all stay dissolved