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MCAT Gen. Chem Ch.10
| Term | Definition |
|---|---|
| Arrhenius Acids | Dissociate to produce an excess of hydrogen ions in solution |
| Arrhenius Bases | Dissociate to produce an excess of hydroxide ions in solution |
| Bronsted-Lowry Acids | Species that can donate hydrogen ions. |
| Bronsted Lowry Bases | Species that can accept hydrogen ions |
| Lewis Acid | Electron-pair acceptors |
| Lewis Bases | Electron-pair donators |
| Note About Arrhenius And Lewis Acid / Bases | All Arrhenius acids and bases are Bronsted-Lowry acids and bases. All Bronsted-Lowry acids and bases are Lewis acids and bases. |
| Amphoteric Species | Species that can behave as an acid or base |
| Amphiprotic Species | Amphoteric species that are amphoteric species that specifically can behave as a Bronsted-Lowry acid or Bronsted-Lowry base |
| Water Is An Example Of A: | Amphoteric, amphiprotic species since it can accept a hydrogen ion to become a hydronium ion, or donate a hydrogen ion to become a hydroxide ion. |
| Conjugate Species Of Polyvalent Acids And Bases Can Also: | Behave as amphoteric and amphiprotic species. |
| Water Dissociation Constant, Kw | 10^-14 at 298 K. This is only affected by changes in temp. |
| Strong Acids And Bases: | Completely dissociate in solution |
| Weak Acids And bases: | Do not completely dissociate in solution and have corresponding dissociation constants (Ka and Kb) |
| Neutralization Reactions | Form salts and sometimes water |
| Equivalent | One mole of the species of interest |
| Normality | Conc. of acid or base equivalents in solution. |
| Polyvalent Acids And Bases | Acids and bases that can donate or accept multiple electrons. |
| Titrations | Used to determine the conc. of a known reactant in a solution. |
| Titrant | Has a known conc. and is added slowly to the titrand to reach the equivalence point |
| Titrand | Has an unknown conc. but a known volume |
| Half-equivalence Point | Midpoint of the buffering region in which half of the titrant has been protonated (or deprotonated) thus [HA] = [A-] and a buffer is formed. |
| Equivalence Point | Indicated by the steepest slope in a titration curve. This is reached when the number of acid equivalents in the original solution equals the number of base equivalents added, or vice-versa. |
| pH = 7 For Equiv Points | This is for Strong Acid and Strong Base Titrations |
| pH > 7 For Equiv Points | This is for Weak Acid and Strong Base Titrations |
| pH < 7 For Equiv Points | This is for Weak Base and Strong Acid Titrations |
| pH Above Or Below 7 For Equiv Points | This can be for weak acid and weak base titrations, depending on the relative strength of the acid and base |
| Indicators | Weak acids or bases that display different colors in their protonated and deprotonated forms |
| Indicator Chosen For A Titration Should Have A pKa Close To: | The pH of the expected equivalence point |
| Endpoint Of A Titration | When the indicator reaches its final color. |
| Multiple Buffering Regions And Equivalence Points Are Observed In: | Polyvalent acid and base titrations |
| Buffer Solutions | Consist of a mix. of a weak acid and its conjugate salt or a weak base and its conjugate salt. They resist large fluctuations in pH. |
| Buffering Capacity | Ability of a buffer to resist changes in pH. Maximal buffering capacity is seen within 1 pH point of the pKa of the acid in the buffer solution. |
| Henderson-Hasselbach Equation | Quantifies the relationship between pH and pKa for weak acids and between pOH and pKb for weak bases. |
| When A Solution Is Optimally Buffered: | pH = pKa and pOH = pKb |
| Eq. 10.1 Autoionization Constant For Water | Kw = [H30+][OH-] = 10^-14 at 25C (298K) |
| Eq. 10.2: Definitions of pH and pOH | pH = -log[H+] = log 1/[H+]. pOH = -log[OH-] = log 1/[OH-] |
| Eq. 10.3: Relationship of pH and pOH at 298 K | pH + pOH = 14 |
| Eq. 10.4: P Scale Value Approximation | p value = m - 0.n |
| Eq. 10.5: Acid Dissociation Constant | Ka = [H30+][A-] / [HA] |
| Eq. 10.6: Base Dissociation Constant | Kb = [B+][OH-] / [BOH] |
| Eq. 10.7: Relationship of Ka and Kb at 298 K | Ka,acid * Kb,conjugate base = KW. Kb,base * Ka,conjugate acid = Kw |
| Eq. 10.8: Equivalence Point | NaVa = NbVb. Na and Nb = acid and base normalities. Va and Vb = acid and base volumes. |
| Eq. 10.9: Henderson-Hasselbalch (Acid Buffer) | pH = pKa + log ([A-]/[HA]) |
| Eq. 10.10: Henderson-Hasselbalch (Base Buffer) | pOH = pKb + log([B+]/[BOH]) |
Created by:
SamB91
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