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Acid-Base Theories

Acid-Base Theories. And Much More. Arrhenius Theory. An Acid is a substance that produces H+ ions in solution. A base is a substance that produces OH- ions in solution. Neutralization occurs when H+ ions combine with OH- ions to form water. Acid or Base (Arrhenius Style). H 2 SO 4

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Acid-Base Theories

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  1. Acid-Base Theories And Much More

  2. Arrhenius Theory An Acid is a substance that produces H+ ions in solution. A base is a substance that produces OH- ions in solution. Neutralization occurs when H+ ions combine with OH- ions to form water.

  3. Acid or Base (Arrhenius Style) • H2SO4 • Ca(OH)2 • Pb(NO3)2 • NaHCO3 • NaOH • HBr

  4. Limitations of Arrhenius Theory • Acids (like HCl) are able to be neutralized by NaOH and NH3. • Not all bases have OH- ions, NH3 in H2O does, but, if NH3 reacts as a gas with HCl gas, NH4Cl is still formed, but is not a solution, so this theory won’t always work.

  5. Bronsted-Lowry Theory • An acid is a proton donor. • A base is a proton acceptor. H3O+ is a hydronium ion (aka. THE PROTON) H2O is the base. HCl is the acid

  6. Writing Conjugates • See Board notes

  7. Conjugate Acid-Base Pairs A conjugate acid is the substance that has accepted the proton (gained an H+) A conjugate base is the substance that has lost the proton. Each acid has a conjugate base and each base has a conjugate acid.

  8. Lewis Theory In the Lewis theory of acid-base reactions, bases donate pairs of electrons and acids accept pairs of electrons. A Lewis acid is therefore any substance, such as the H+ ion, that can accept a pair of nonbonding electrons. In other words, a Lewis acid is an electron-pair acceptor. A Lewis base is any substance, such as the OH- ion, that can donate a pair of nonbonding electrons. A Lewis base is therefore an electron-pair donor.

  9. Ionization Constants Acid ionization constant (Ka) Base ionization constant (Kb) Magnitude dictates the strength of the base Strong bases have large Kb values – and dissociate completely (100%)  Weak bases have small Kb values - and do not dissociate completely (5% or so)  • Magnitude dictates the strength of the acid • Strong acids have large Ka values – this means that they dissociate completely (100%)  • Weak acids have small Ka values – and do not dissociate completely (usually 5% or so) 

  10. Relative Strengths of Acids and Bases

  11. Self-ionization of Water

  12. Ion Product of Water (Kw) Kw = 1.0 x 10-14=[H+][OH-] – changes with temperature Kw = Ka x Kb • Uses: Can find concentration of OH- or H+, the Kb of the conjugate base (and vice versa), or used with pH or pOH calculations when necessary

  13. pH and pOH pH = -log10[H+] or pH = -log10[H3O+] (same thing) [H+] = 10-pH pOH = -log10[OH-] [OH-] = 10-pOH

  14. pH Scales

  15. Strong Acids & Bases • Strong acids – hydrogen halides, oxyacids of halogens, sulfuric acid & nitric acid • Strong bases – NaOH, KOH, Cs(OH)2, Ca(OH)2 • Both of these ionize completely in H2O.

  16. Weak Acids & Bases • Weak acids and bases do not ionize completely in water • A weak acid/base will be shown with a double arrow reaction, and also have small Ka and Kb values

  17. Ionization Constants of Weak Acids & Bases

  18. Equilibrium Problems Using Acids Type 1: Solving for Ka You must have the following: - the equation for the reaction - the equilibrium expression - the concentrations at equilibrium (ICE chart) - The concentrations at equilibrium are usually found with pH or a % ionization.

  19. Equilibrium Problems Using Acids Type 2. Solving for EQ amounts You must have the following: - the reaction - the equilibrium expression and Ka value - the initial concentration of the acid (in mol/L) - With this type, it may ask for the pH at equilibrium

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