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Acids & Bases

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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  1. Acids & Bases …all you need to “get” for the test… In 20 minutes!

  2. 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

  3. 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)

  4. 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.

  5. Try these for examples: • HF + H2O H3O+ + F- • NH4+ + OH- NH3 + H2O • CO32- + H2O HCO3- + OH-

  6. 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.

  7. 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)

  8. 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

  9. 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

  10. Practical examples

  11. More…

  12. 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)

  13. Your turn… • Compile 3 questions to ask/clarify/review: • 1. • 2. • 3.

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