1. Unit VI Acids, Bases & Salts
Textbook Chapter 15
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Unit Outline
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2. Electrolytes
Dissolve in water, solution formed will conduct electricity. Why conduct?
Conductivity due to presence of free ions. How free ions?
Free ions can be created from processes of dissociation or ionization.
Acids, Bases & Salts are examples of electrolytes.
Electrolytes Poem
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3. Cations are positive.
Anions are negative.
Dissociation
Occurs when ions of an ionic solid are separated from the crystal lattice structure.
Polar water molecules pull ions from lattice, surround ions and form hydrated ions.
Examples: all ionic compounds (salts) such as NaCl(aq), CuSO4(aq), MgBr2(aq)
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4. Occurs when a molecular solid is dissolved and creates ions in solution.
Unlike dissociation, ions were not present originally.
Common examples: Acids
HC2H3O2 H2O> H+(aq) + C2H3O2-(aq)
Ionization
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5. Weak or Strong Electrolyte?
Depends on amount of ionization.
Reference Tables K & L list strong electrolytes closer to the top (stronger acids and bases)
Keq values such as Ka or Kb can indicate amount of ionization. (Remember Keq’s Δ only w/ temp.)
Large Ka  strong electrolyte (acid); many free ions
Small Ka  weak electrolyte (acid); few free ions
Complete ionization of HCl animation
Incomplete ionization of weak acid animation
Comparing HCl ionization to HF ionization animation
Conductivity of HCl vs. Acetic Acid video
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6. Nonelectrolytes Do not conduct electricity when in solution.
Examples: All organic compounds except organic acids.
Examples: C12H22O11(aq), C2H5OH(aq)
AB(s)  AB(aq) {no free ions}
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7. How do you define something?
Operational Definition-list properties or behaviors. Include observations from experiments.
Conceptual Definition-tries to answer “Why?” and “How?”. Based on interpretations or conclusions from observed facts.
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8. What is an Acid? 6 Operational Definitions of an Acid
Bases

9. Acids: Op Def #1 Aqueous solutions of acids conduct electricity
Weak acidsweak electrolytefew ionsslight ionization
Strong acidsstrong electrolyteMany ionsAlmost complete ionization
Chinese for strong electrolyte (strong acid) ???
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10. Acids: Op Def #2
Acids will react with certain active metals to liberate hydrogen gas.
Think Rocket Lab!
2HCl(aq) + Zn(s)  ZnCl2(s) + H2(g)
Use Ref Table J. Where are metals found on table that will undergo this reaction?
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11. Acids: Op Def #3
Acids cause color changes for acid-base indicators. See Ref Table M.
Note from Table M that all indicators have their own pH range. Depends on the concentration of H+ in solution. Indicators do not change right at pH of 7.
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12. Acids: Op Def #4
Acids react with hydroxides (bases) to form water and a salt. Called neutralization.
Example: HNO3(aq) + NaOH(aq)  H2O + NaNO3(aq)
Remember from Lab activity to make salt!!!
Neutralization simplified: H+(aq) + OH-(aq)  H2O
Try another example on next slide!
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What is a salt?

13. Acids react with carbonates to release carbon dioxide.
HCl(aq) + Li2CO3(s)  complete this rxn
2HCl(aq) + Li2CO3(s)  LiCl(aq) + H2CO3
Unstable and decomposes
2HCl(aq) + Li2CO3(s)  LiCl(aq) + H2O(l) + CO2(g)

14. Acids: Op Def #5 Dilute aqueous solutions of acids have a sour taste.
Examples: vinegar (acetic acid) in pickles, citric acid in lemonheads/sour candy or citrus fruits (lemons, limes, oranges, etc.)
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15. Acids: Op Def #6
Acids react with metallic oxides to form salts and water.
Example: HCl + K2O  2KCl +H2O
Potassium Oxide
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16. What is a Base? 5 Operational Definitions of a Base
Conceptual Definitions

17. Bases: Op Def #1 Aqueous solutions of bases conduct electricity
Weak basesweak electrolytesfew ionsslight dissociation
Strong basesstrong electrolytesMany ionsAlmost complete dissociation
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18. Bases: Op Def #2
Bases cause color changes in acid-base indicators. See Ref Table M.
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19. Bases: Op Def #3
Bases react with acids to form water and a salt. Called neutralization.
Example: HC2H3O2(aq) + NaOH(aq)  H2O + NaC2H3O2(aq)
Neutralization simplified: H+(aq) + OH-(aq)  H2O
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What is a salt?

20. Bases: Op Def #4
Aqueous solutions of bases feel slippery and taste bitter.
Examples: soaps, cleaners
Armor All is great on tires and the dusty dashboard. Why never on the steering wheel?
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21. Bases: Op Def #5
Strong bases have a caustic action on the skin, corrosive to skin and poisonous.
Examples: really strong cleaners, concentrated solutions of KOH or NaOH can cause blindness if in eyes
Burn from conc. NaOH
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22. 2 Conceptual Definitions of Acids and Bases
Remember conceptual definitions add reasons to observations. Answer Why?
Arrhenius Theory
Brönsted-Lowry Theory (AKA an alternate acid/base theory according to NYSED)
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23. Arrhenius’ Theory of an Acid (1887)
A substance that yields hydrogen ions (H+) in aqueous solutions.
Arrhenius said properties of acids in aqueous solutions due to an excess of H+ or H3O+ (hydrogen or hydronium ions)
Good definition but limited to only (aq).
Examples:
HNO3 H2O H+(aq) + NO3-(aq)
H2SO4 H2O 2H+(aq) + SO42- (aq)
Svante Arrhenius,
The Nobel Prize in Chemistry 1903
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24. Arrhenius’ Theory of a Base (1887)
A substance that yields hydroxide ions (OH-) as the only negative ions in aqueous solutions.
Examples:
KOH H2O K+(aq) + OH-(aq)
Ba(OH)2 H2O Ba+2(aq) + 2OH-(aq)
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25. Brönsted-Lowry Theory of an Acid (1923)
Any species (molecule or ion) that can donate a proton to another species.
AKA: “a proton donor”
Brönsted-Lowry does not replace Arrhenius’ theory, it just extends it. Includes Arrhenius examples of acids but adds other examples.
Johannes Nicolaus Brønsted
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Thomas Martin Lowry
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26. Brönsted-Lowry Theory of a Base (1923)
Any species(molecule or ion) that can accept a proton.
AKA: “a proton acceptor”
Again Brönsted-Lowry does not replace Arrhenius’ theory, it just extends it. More species can accept a proton than just an OH-.
Ex: H2O + HCl ⇆ H3O+ + Cl-
Water is base in forward rxn, accepts H+ to become H3O+
In reverse rxn, Cl- accepts proton and acts as base to form HCl
I luv accepting protons.
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27. B-L Conjugate Acid-Base Pairs
Bronsted-Lowry acid-base rxn,
Transfer of proton (H+) from acid to base
To accept proton, must have pair of unshared electrons.
An acid give a proton to form a conjugate base.
A base gains a proton to form a conjugate acid.
Ex: HSO4- + H2O  H3O+ + SO42-
Label reacting acid and base and conjugates formed.
Do Examples!
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28. Lewis A&B Theory defined in terms of electron pair
Lewis acids- electron-pair acceptor
Lewis bases- electron pair donor.
Remember coordinate covalent bonds?
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29. Amphoteric(Amphiprotic) Substances
Can act as either an acid or a base.
Pick out the amphoteric substances in the following:
NH3 + H2O ⇆ NH4+ + OH-
H2O + HCl ⇆ H3O+ + Cl-
NH3 + OH- ⇆ NH2- + H2O
Wafflers Unite!
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30. Naming Acids Binary acids (2 elements) use hydro……ic acid
ex: HCl hydrochloric acid
Ternary acids (3 elements) use polyatomic ion name and change ending
ate ic
ite  ous
Ex: H2SO4
(sulfate) Sulfuric acid
Ex: H2SO3
(sulfite) Sulfurous acid
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31. Naming Salts Binary salts Ex: CuCl2 Ternary salts Ex: Ca3(PO4)2
Change ending to ide
Use Roman numeral if necessary.
Ex: CuCl2
Copper (II) chloride
Ternary salts
Use polyatomic ion name
Ex: Ca3(PO4)2
Calcium phosphate
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What is a salt?

32. Naming Bases Ex: Ba(OH)2 Ex: Fe(OH)2
What is the ion Arrhenius said all bases have?
Use it to name bases. Use Roman numeral if necessary.
Ex: Ba(OH)2
Barium hydroxide
Ex: Fe(OH)2
Iron (II) hydroxide
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33. Neutralization A + B  water + salt
Complete and balance the following neutralization rxn: HNO3 + Ca(OH)2 
2HNO3 + Ca(OH)2  2H2O + Ca(NO3)2
Name all compounds in this rxn.
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34. Hydrolysis The reverse of neutralization.
Some salts in aqueous solutions react with water and form solutions that are acidic or basic.
Salt + H2O  Acid + Base
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35. Examples of Hydrolysis
Salt made from strong acid and weak base give aq solutions that are acidic.
Ex: NH4Cl + H2O  HCl + NH4OH
Salts from weak acids and strong bases give aq solutions that are basic.
Ex: NaC2H3O2 + H2O  HC2H3O2 + NaOH
Salts from strong acid and strong base usually do not hydrolyze. If they do, form neutral pH of 7.
Ex: NaOH and HCl
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36. Acid-Base Titration:a very precise neutralization
The molarity (strength) of an acid (or base) of unknown concentration can be determined by slowly combining it with a base (or acid) of known molarity (standard solution).
A titration is finished when the equivalence or end point is reached, this indicates neutralization and is determined from color changes of indicators.
Buret
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37. Titration Formula Ref Table T
MAVA = MBVB
Show derivation of formula on overhead.
Before using titration formula, determine mole ratio of acid:base from balanced equation.
Practice Examples!
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38. Ionization Constants
Remember Keq from last unit?
In this unit we deal mainly with ionization of acids (Ka) and bases (Kb)
Ka & Kb provide method for comparing relative strength of acids or bases.
Example below:
CH3COOH H20> H+(aq) + CH3COO-(aq)
Size of Ka value tells strength of acid
NYS Chem Reference Table pre-2002

39. Ionization Constant for Water (Kw)
H2O H+ + OH- or 2H20  H3O+ + OH-
Since concentration of H2O is relatively constant, Keq expression can be simplified to ….
At STP only 1.0 X 10-7 moles/L of H2O will ionize into equal amts of H+ & OH- so therefore….
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40. Kw = 1.0X10-14 @ 25oC w/standard pressure.
Kw = [H+][OH-]
Kw = [1.0X10-7][1.0X10-7]**
Kw = 1.0X10-14**
** This is why water is neutral. Because acidity (H+) and basicity (OH-) are in equal concentrations.
Kw = 25oC w/standard pressure.
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41. Sample Problem
What happens to [H+] if [OH-] concentration jumps to STP?
Kw = [H+][OH-]
1.0X10-14 = [H+][1.0X10-5]
[1.0X10-14]/ [1.0X10-5] = [H+]
[H+] = [1.0X10-9]
[OH-] got bigger, [H+] got smaller

42. What do you know about the pH scale?
0-14
7 is neutral
Acids below 7, Bases above 7.
Stronger the acid or base the further away from neutral.
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43. pH pH is the logarithm of the reciprocal of the H+ concentration
pH of a solution tells the concentration of H+ (how acidic it is)
Ex: [H+]=1.0 X 10-9
pH = 9
pH is absolute value of exponent for [H+]
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44. pH scale What does a logarithmic scale mean?
Each Δ in pH # by 1 = Δ in [H+] or [H3O+] by 10
Powers of 10 video
Scale of Universe (powers of 10)
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45. [H+]= 1.0 X 10-7 same as [OH-], balanced pH= 7, neutral
Remember water at 25oC!
[H+]= 1.0 X 10-7 same as [OH-], balanced
pH= 7, neutral
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46. pH sample problems What is the pH in each of the following examples?
[H+]=1.0 X 10-3
pH = 3
[OH-]=1.0 X 10-11
Kw=[H+][OH-]
1.0X10-14 = [H+][1.0X10-11]
1.0X10-14/1.0X10-11 = [H+]
[H+] = 1.0X10-3
pH=3

47. pH sample problems What is the pH in each of the following examples?
0.001M HCl solution
0.001M  1.0X10-3 = [H+]
pH=3
0.01M NaOH solution
0.01M  1.0X10-2 = [OH-]
Kw = [H+][OH-]
1.0X10-14 = [H+][1.0X10-2]
1.0X10-14/1.0X10-2 = [H+]
[H+] = 1.0X10-12
pH=12  basic

48. pOH pOH = -log [OH-] pH + pOH = 14.0
Taken from sciencegeek.net on 7/29/11.

49. Buffers
Solutions used to minimize ( not prevent) a change in pH when an additional acid or base is introduced into solution.
Buffers consist of a conjugate pair of a weak acid and weak base in an equilibrium system.
Taken from boughen.com.au on 7/28/11.

50. Calculating the pH of Buffers
The Henderson-Hasselbalch equation:

51. Salts
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An ionic compound containing positive ions other than H+ and negative ions other than OH-
Ionic compound, bonding?
Ionic bonding
Many salts are strong electrolytes.
Completely dissociate in solution into free ions and resulting solution will conduct electricity. “Break crystal lattice.”

52. 5 Methods for Preparing a Salt
Neutralization
Single Replacement Reaction
Direct Combination of Elements
Double Replacement
Reaction of a metallic oxide with a nonmetallic oxide.