chem unit 2

arhennius acid

produces H ions when dissolved in aqueous solution; [HA]

Using ICE vs SRF tables

ICE: determining eq concentrations, will give Ka or Kb, use molarity; SRF: used mostly for buffer reactions, SA +SB reactions

Arrhenius base

produces hydroxide ions (OH-) when dissolved in aqueous solution; [A]

Equivalence point

Moles of titrant are EQUAL

bronsted lowry acid

proton (H+) donor (HCl)

bronstead Lowry base

proton (H+) acceptor (NH3)

conjugate acid

the species formed when a proton is transferred to the base; NH3→ NH4+

conjugate base

what remains of an acid molecule after a proton is lost; HCl → Cl-

hydronium ion

H3O+; the predominating form of the proton in an aqueous solution; can use H+ and H3O+ interchangeably

strong acid

complete dissociation of that ion to produce H+ and the conjugate base; lose H+ much easier; HA + H2O → A +H3O ; single arrow forward means complete dissociation

weak acid

partial dissociation; HA + H2O ←→ A+H3O ; arrow forward and back to show it has to move forward and back to reach equilibrium

strong base

a metal hydroxide salt that completely dissociates into its ions in water; holds onto H+ strongly; H+H2O → HB + OH; single arrow shows complete dissociation

weak base

a base that reacts with water to produce hydroxide ions to only a slight extent in aqueous solutions; BH2 + H2O ←→ BH3 +OH ; arrow forward and back to show it has to move forward and back to reach equilibrium

Ka

acid dissociation constant; numerical measure of an acids strength in aqueous solution; higher Ka means stronger acid; [A][H3O]/[HA]

autoinization

the transfer of a proton from one molecule to another of the same substance; OH to H2O or vice versa; water is amphoteric- can be acid or base

Kw

ion product constant; equilibrium constant for the autoionization of water; = 1.0×10^-14; Kw=[H3O][OH]; pKw=pH + pOH=14

significant figures and pH: characteristic and mantissa

pH=7.00; characteristic: related to the exponent of the power of 10 (7); mantissa: indicates the number of sig figs (00)

calculating the pH of weak acids

weak acids will dissociate and form an equilibrium with their conjugate bases; set up Ka expression; set up ice table; solve for x and determine concentrations and then convert to pH

calculating pH of weak bases

weak bases will dissociate and form an equilibrium with their CA; set up Kb expression and use ice table

salt

ionic compound

predicting whether a salt solution will be acidic, basic, or neutral

1) strong base and strong acid: salt has no hydrolyzable ions and form a neutral salt solution; 2)strong base and weak acid: produces basic salt; 3)weak base and strong acid: produces acidic salt

common ion effect

the shift in equilibrium position caused by the addition or presence of an ion involved in the equilibrium reaction

buffer solution

resists change in pH when OH ions or H protons are added; consists of a weak acid and its salt (conjugate base) or consists of a weak base and its salt (conjugate acid)

Henderson hasselbach equation (provided)

used to calculate buffer pH or design buffers; relate the pH of a buffer solution to the pKa of a weak acid and its conjugate base

buffering capacity

[A]/[HA]; the ability of a buffered solution to absorb protons or hydroxide ions without a significant change in pH ; determined by magnitudes of [HA] and [A] in solution; the most effective buffer is one where the ratio is close or equal to 1, meaning when pH is close to pKa

titration

the process of reaction a solution of unknown concentration with one of known concentration (standard solution)

titration curve

graph of the pH as a function of the volume of added titrant

strong acid-strong base titration

equivalence point

weak acid strong base titration

initial pH, buffer region: between initial and before equivalence; equivalence point; half equivalence point right in the middle of the buffer region: where pH=pKa; pH at equivalence point is above 7

strong acid weak base titration

2: half equivalence point and buffer region, pH=pKa; 3= equivalence point; equivalence point is below 7

determining equivalence point experimentally

use an acid-base indicator which marks the endpoint of titration by changing color

how to choose which indicator to use; you will have the color chart

select one whose pH color change range overlaps the steep vertical section of your titration curve, usually centered near the equivalence point

percent dissociation

concentration dissoicated/original concentration