Acids and Bases

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Acids and Bases

• What is an Acid?

What is a Base?
Before we answer these questions, lets look at
the properties of acids and bases.
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Properties of Acids

• Tart or sour taste
• Electrolytic
• Reacts with many metals to produce Hydrogen gas
• Acid Metal ? Salt H2
• Example HCl Zn ? ZnCl H2
• Reacts with bases to form water and salt
• Acid Base ? Salt Water
• Example HNO3 NaOH ? NaNO3 H2O
• Reacts with Carbonates to form salt, water, and
  carbon dioxide.
• Acid Carbonate ? Salt Water CO2
• Example H2SO4 Na2CO3 ? Na2SO4 H2O
  CO2
• Reacts with ammonia to form ammonium salts.
• Acid Ammonia ? Ammonium salt
• Example HI NH3 ? NH4I

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Properties of Bases

• Taste Bitter
• Feel Slippery
• Reacts with acids to form salt and water

Picture of Quick Lime Ca(OH)2, used for making
cement, mortar, and soil stabilization.
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Recognizing and naming bases

• All bases that are ionic compounds will have the
  hydroxide ion present and are named just like
  ionic compounds.
• Examples
• NaOH
• Mg(OH)2
• Fe(OH)3
• Ammonia (NH3) is also a base
• NH3 H2O ? NH4OH
• The NH4OH then dissociates into NH4 OH-

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Definitions of acids and bases

• There are a three different definitions
• Arrhenius Acids and Bases
• According to the Arrhenius definition, an acid is
  a compound that ionize to yield H ions (often
  expressed as H3O) in aqueous solution. Bases
  compounds that ionize to produce OH- ions in
  solution.
• Example
• HCl H2O ? Cl- H3O
• Notice that in reality hydrogen ions are not
  produced but hydronium ions. For our purposes,
  they are synonymous.

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Bronsted-Lowry Acids (and bases)

• An acid is a hydrogen ion donor
• A B-L base is a hydrogen ion acceptor
• This is what allows ammonia to be a base it
  does not donate hydroxide ions (as an Arrhenius
  definition would imply) but it accepts hydrogen
  ions.
• Water in this instance is also an acid or a base!
  AMPHOTERIC

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The Amphoteric Water Molecule

• H2O H2O ? H3O OH-
• Notice that when the water molecule dissociates
  it donates or accepts both hydrogen ions
  (hydronium) and hydroxide ions.
• In pure water, this self ionization of water
  occurs such that the concentration of each of the
  respective ions is approximately 1 x 10-7 mol/L.
  We will discuss this later when we address the
  pH concept.
• In any aqueous solution H and OH- are
  interdependent. If H increases then OH- must
  decrease. The mathematical relationship is
• Kw H x OH- 1 x 10-14mol/L.
• Do you see why pure water has a H of 1 x 10-7
  mol/L?
• Kw ion product constant

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Lewis acids and bases

• The most general definition of acids and bases,
  which encompasses the Arrhenius and
  Bronsted-Lowry definitions is due to our old
  friend, Lewis and his dot structures.
• A Lewis acid is defined to be any species that
  accepts lone pair electrons.
• A Lewis base is any species that donates lone
  pair electrons.
• Thus, H is a Lewis acid, since it can accept a
  lone pair, while OH- and NH3 are Lewis bases,
  both of which donate a lone pair

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Odd Lewis Acids

• Interestingly, however, is that species which
  have no hydrogen to donate (a la the
  Bronsted-Lowry scheme) can still be acids
  according to the Lewis scheme. As an example,
  consider the molecule BF3 . If we determine Lewis
  structure of BF3 , we find that B is octet
  deficient and can accept a lone pair. Thus it can
  act as a Lewis acid. Thus, when reacting with
  ammonia, the reaction would look like

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Recognizing and Naming Acids

• Acids usually come in the form HX
• X is a mono or polyatomic anion
• Some Lewis and organic acids do not come in this
  form (we will not concern ourselves with naming
  these in General Chemistry).
• Some compounds like HCl(g) are named hydrogen
  chloride they are covalent molecules. However,
  when dissolved in water, become electrolytic (or
  ionic) and are considered an acid and are named
  differently.

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• 1. When the name of the anion ends in ide, the
  acid name begins with the prefix hydro- and the
  anion takes on the suffix ic. It is followed by
  the word acid.
• For example H2S - Hydrogen Sulfide
• As an acid it is Hydrosulfuric Acid
• 2. When an anion ends in ite, the suffix of the
  anion is changed to ous, followed by the word
  acid.
• For example H2SO3 Hydrogen Sulfite
• As an acid it is Sulfurous Acid
• 3. When the anion ends in ate, the suffix ends
  in ic, followed by the word acid.
• For example H2SO4 Hydrogen Sulfate
• As an acid it is Sulfuric Acid

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Practice Naming Acids

• Nitrous Acid
• Hydroiodic Acid
• Hydrofluoric Acid
• Phosphoric Acid
• Phosphorous Acid
• Chlorous Acid
• Hypochlorous Acid
• Carbonic Acid
• Hydrocyanic Acid
• Acetic Acid
• Dichromic Acid
• Perchloric Acid

• HNO2
• HI
• HF
• H3PO4
• H3PO3
• HClO2
• HClO
• H2CO3
• HCN
• HC2H3O2
• H2Cr2O7
• HClO4

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Strong Acids and Bases

• As discussed earlier, acids and bases ionize in
  water to produce either hydrogen or hydroxide
  ions. Some acids and bases completely ionize
  some do not.
• Those that completely ionize are considered
  STRONG acids or bases. You will need to memorize
  these

Strong Bases Sr(OH)2 Ca(OH)2 Ba(OH)2
LiOH NaOH KOH RbOH
CsOH
Strong Acids HI HCl HBr HNO3 HClO4 H2SO4
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The ionization of a strong acid or base

• HI H2O ? H3O I-
• There is 100 ionization in this reaction.
• If one started with a 1.0M solution of HI, what
  would be the concentration of Hydrogen ions
  (Hydronium)?
• Answer 1.0M
• NaOH ? Na OH-
• Same is true for strong bases. In other words,
  there is no NaOH still in solution, it is all
  ionized.

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Weak Acids and Bases

• Weak acids and bases do NOT completely ionize in
  solution.
• HC2H3O2 H2O ?? H3O C2H3O2-
• Notice the double arrow, this means that the
  reaction goes both ways. Therefore all four
  species in this equation will exist in the
  solution at different concentrations.
• If one had a 1.0M solution of acetic acid, there
  would not be enough information to determine the
  concentration of H3O . This becomes important
  when determining pHs of solutions.

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The strengths of weak acids and bases Ka Kb
values.

• The more an acid or base dissociates in solution,
  the more ions it puts in solution, therefore it
  is stronger.
• Mathematically this can be expressed as a
  dissociation constant.
• Ka , the acid dissociation constant
  HA-/HA. This is the product of the
  concentrations of the dissociated form divided by
  the undissociated form.
• The larger the number the stronger the acid.
• Kb base dissociation constant, same principle
  as the acid dissociation constant

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Example of a problem using Ka

• What is the H of a 0.01M solution of carbonic
  acid?
• Look at the dissociation equation
• H2CO3 ?? H HCO3-
• Look up the Ka value for carbonic acid 4.3 x
  10-7
• Understand that the cation and anion dissociate
  equally and therefore will have equal
  concentrations HHCO3-
• Rearrange dissociation equation and solve for
  H
• H H2CO3 Ka

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Example continued

• H (0.01M)(4.3 x 10-7) 6.55
  x 10-5M
• The Kb for ammonia is 1.8 x 10-5. The equation
  looks like this
• NH3 H2O ?? NH4 OH-
• The concentration of water does not need to be
  considered.
• What is the hydroxide ion concentration when the
  ammonia concentration is 0.5M?
• Did you get 3.0 x 10-3M?

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Calculating pH or pOH

• pH is the unit we use to measure the
  concentration of acids and bases. (pOH can also
  be used)
• pH -logH pOH -logOH-
• Example what is the pH of a solution in which
  the H is 1 x 10-7 mol/L
• pH -log(1 x 10-7 mol/L)
• -(log 1 log 10-7)
• -(0.0 (-7)) 7.0

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Determining pH, pOH, OH-, H3O Remember H
H3O
Use this chart to determine unknowns given one
value
Kw/ H3O
H3O
OH-
Kw/OH-
-log H3O
10-pH
10-pOH
-logOH-
14.0 - pH
pH
pOH
14.0 - pOH
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Practice Problems

• What is the H of stomach acid with a pH of
  0.5?
• What is the pH of a 0.02M solution of HCl?
• What is the pOH of a solution that has a pH of
  6.2?
• What is the pH of a 0.05M solution of acetic
  acid? Ka of acetic acid is 1.8 x 10-5
• What is the OH- of a solution that has a H
  of 2.3 x 10-4?

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pH scale and some common household products