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This concise overview of acids and bases condenses essential definitions and concepts you'll need for your test. Understand how acids donate hydrogen ions and how bases accept them. Learn about ionization processes, pH calculations, and the relationship between hydronium and hydroxide ions. Explore the significance of strong and weak acids, as well as practical examples of neutralization and buffers. Prepare effectively and quickly with this guide that ensures you grasp key ideas in just 20 minutes.
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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) • 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….. 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… • 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: • HF + H2O H3O+ + F- • NH4+ + OH- NH3 + H2O • CO32- + H2O HCO3- + OH-
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.
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 • 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] • 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!) • 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… • Compile 3 questions to ask/clarify/review: • 1. • 2. • 3.
