Acids bases
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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

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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 + H2OH3O+ + F-

    • NH4+ + OH- NH3 + H2O

    • CO32- + H2OHCO3- + 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- + NaOHH2O(l)+ Na+ + SO42-+ OH-

      • The resulting solution is still basic.


    Acids bases

    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


    Practical examples

    Practical examples


    Acids bases

    More…


    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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