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Acids & Bases. …all you need to “get” for the test … In 20 minutes!. Definitions. Produces hydronium in aqueous (water) solutions (Arrhenius) Donates hydrogen ions to another species ( Bronsted -Lowry) Taste sour pH < 7 Turns litmus (and many other indicators red).

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acids bases

Acids & Bases

…all you need to “get” for the test…

In 20 minutes!

definitions
Definitions
  • Produces hydronium in aqueous (water) solutions (Arrhenius)
  • Donates hydrogen ions to another species (Bronsted-Lowry)
  • Taste sour
  • pH < 7
  • Turns litmus (and many other indicators red)
  • Produces hydroxide in aqueous (water) solutions (Arrhenius)
  • Receives hydrogen ions from acid (Bronsted-Lowry)
  • Taste bitter; feel slippery
  • pH > 7
  • Turns litmus (and many other indicators blue)

Acid

Base

the ionization process
The ionization process…..

A compound’s ability to behave as an acid is that’s compound’s ability to “donate” hydrogen ions (protons).

    • “Strong” acids release those ions VERY readily and completely
    • For example CH4 is NOT an acid—at all!
  • That donation is represented thusly:
    • H2SO4 + H2O HSO41- + H3O1+(1st ionization)
    • HSO41- + H2O SO42- + H3O1+ (2nd ionization)
ions in aqueous solutions exist in equilibrium
Ions in Aqueous solutions exist in equilibrium…
  • HSO41- + H2O SO42- + H3O1+
  • What you should notice:
    • HSO41- becomes SO42-; therefore, (donates H1+)
      • in the reverse, SO42- becomes HSO41- (receives H+)
    • H2O becomes H3O1+; therefore, (receives H+)
      • In the reverse, H3O1+ becomes H2O (donates H+)
  • Translation: for weak ionizations and/or dilute solutions, that are reversible (in equilibrium), acids become conjugate bases, and, conversely, bases become conjugate acids.
try these for examples
Try these for examples:
  • HF + H2O H3O+ + F-
  • NH4+ + OH- NH3 + H2O
  • CO32- + H2O HCO3- + OH-
consider
Consider:
  • Hydronium ions in the presence of hydroxide ions can form water!
  • Of course, the leftovers ions form a “salt”.
  • For example:
    • HCl(aq)+ NaOH(aq) H2O(l)+ Na+(aq) Cl-(aq)
    • Because both the acid and the base are “strong”, the resulting hydronium and hydroxide concentration are equal.
    • The resulting pH is neutral. The “salt” is sodium chloride.
  • Another example:
    • HSO4- + NaOH H2O(l)+ Na+ + SO42-+ OH-
    • The resulting solution is still basic.
slide7
pH
  • The actual measurements of concentration result in the calculation of pH.
  • Pure water is defined by equal concentrations of hydrogen ions and hydroxide ions.
    • [H3O+] = [OH-] = 1 x 10-7M
    • [H3O+] x [OH-] = 1 x 10-14(memorize these numbers)
the scale
The scale
  • Using the logarithmic function of those concentrations, we get the pH scale:
  • Water has a pH of 7
    • pH = -log [H3O+]
  • Higher concentrations of hydronium means a smaller log!
    • 2.34 x 10-4 [H3O+] = 3.63
  • Smaller concentrations mean higher logs!
    • 2.34 x 10-10 [H3O+] = 9.63
relating hydronium hydroxide
Relating [hydronium] & [hydroxide]
  • Because a species is only an acid or a base in water, the concentrations of these ions are related:
    • [H3O+] [OH-] = 1 x 10-14
    • Which means that as one concentration increases, the other decreases…. (don’t forget the constant.)
    • One can also take the pOH of the hydroxide concentration.
    • Interestingly, pH + pOH = 14
buffers a little extra
Buffers (a little extra!)
  • Buffer- a solution that resists changes in pH when limited amounts of acid OR base are added.
  • Ions of “weak” acids and bases, by definition, mean ions that are available to receive or to donate hydrogen ions &/or hydroxide ions.
  • CO2(g) + H2O(l) H2CO3 (aq) H+ (aq) + HCO3-(aq)
your turn
Your turn…
  • Compile 3 questions to ask/clarify/review:
  • 1.
  • 2.
  • 3.
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