University:
Eastern Kentucky University
Course:
CHE 111 | General Chemistry
Academic year:

2020
Views:

129

Pages:

1
Author:

Ean Mason

General Chemistry 10: Acids and Bases Definitions Polyvalence & Normality Arrhenius Acid: Produces H+ (same definition as Brønsted acid) Arrhenius Base: Produces OH- Equivalent: 1 mole of the species of interest. Normality: Concentration of equivalents in solution. Brønsted-Lowry Acid: Donates H+ (same definition as Arrhenius acid) Polyvalent: Can donate or accept multiple equivalents. Brønsted-Lowry Base: Accepts H+ Example: 1 mol H3PO4 yields 3 mol H+. So, 2 M H3PO4 = 6 N. Lewis Acid: Accepts e- pair Lewis Base: Donates e- pair Note: All Arrhenius acids/bases are Brønsted-Lowry acids/bases, and all Brønsted-Lowry acid/bases are Lewis acids/bases; however, the converse of these statements is not necessarily true. Amphoteric Species: Species that can behave as an acid or a base. Amphiprotic = amphoteric species that specifically can behave as a BrønstedLowry acid/base. Polyprotic Acid: An acid with multiple ionizable H atoms. Titrations Half-Equivalence Point: The midpoint of the buffering region, in which half the (midpoint) titrant has been protonated or deprotonated. [HA] = [A& ] and pH=pK) +and a buffer is formed. Equivalence Point: The point at which equivalent amounts of acid and base have reacted. 𝑁' +𝑉 ' =𝑁P +𝑉 P pH at Equivalence Point: Strong acid + strong base, pH = 7 Weak acid + strong base, pH > 7 Weak base + strong acid, pH < 7 Weak acid + weak base, pH > or < 7 depending on the relative strength of the acid and base Properties Indicators: Weak acids or bases that display different colors in the protonated and deprotonated forms. The indicator’s pKa should be close the pH of the equivalence point. &'( Water Dissociation Constant: 𝐾" = 10 at 298 K 𝐾" =𝐾) ×+𝐾 pH and pOH: pH=+−log+[H4 ] pOH=+−log+[OH& ] pH+ pOH= 14 Tests: Litmus: Acid = red; Base = blue; Neutral = purple Phenolphthalein: pH < 8.2 = colorless; pH > 8.2 = purple Methyl Orange: pH < 3.1 = red; pH > 4.4 = yellow Bromophenol Blue: pH < 6 = yellow; pH > 8 = blue [H4 ] = 10&67 p scale value approximation: −log+(𝐴+×10&=) p+value+ ≈ −(𝐵 + 0.𝐴) Endpoint: When indicator reaches full color. Polyvalent Acid/Base Multiple buffering regions and equivalence points. Titrations: Strong Acids/Bases: Dissociate completely Weak Acids/Bases: Do not completely dissociate Acid Dissociation Constant: 𝐾) = [7FG Base Dissociation Constant: 𝐾 = , H ][IJ ] [7I] [KH ][G7J ] [KG7] p𝐾) =+ −log+(𝐾) ) Titration Setup p𝐾, =+ −log+(𝐾, ) Burette p𝐾) +p𝐾, =p𝐾" = 14 Conjugate Acid/Base Pairs: Strong acids & bases / weak conjugate Weak acids & bases / weak conjugate Neutralization Reactions: Form salts and (sometimes) H2O Titrant (strong acid in this example) Analyte / Titrand (weak base in this example) Conical flask Buffers Buffer: Weak acid + conjugate salt Weak base + conjugate salt Buffering Capacity: The ability of a buffer to resist changes in pH. Maximum buffering capacity is within 1 pH point of the pKa. Henderson-Hasselbalch pH=pK) + log Equation: pOH=pK, + log [IJ ] [7I ] [ KH] [7G7 ] When [A-] = [HA] at the half equivalence point, log(1) = 0, so pH = pKa Titration Curve When titrating a weak base with a strong acid Midpoint pOH =pK, Equivalence Point 𝑁' +𝑉' =𝑁P +𝑉 P