CHM1046 Exam 3 Study Guide Chapter 15: Acids and Bases 15.2: The Nature of Acids and Bases Acid: Characterized by: Sour taste. Ability to dissolve many metals. Ability to turn blue litmus paper red. Ability to neutralize bases. Dissolve into ions when dissolved in water. common acids: HCl, H2SO4, HNO3, HC3H3O2 Base: Characterized by: Bitter tate. Ability to turn red litmus paper blue. Ability to neutralize acids. common bases: NaOH, KOH, NaHCO3, Na2CO3, NH3 15.3: Definitions of Acids and Bases Arrhenius Definitions: acid: a substance that produces H+ ions in aqueous solutions HCl(aq)H+(aq)+Cl-(aq) base: a substance that produces OH- ions in aqueous solutions NaOH(aq)Na+(aq)+OH-(aq) Bronstedt-Lowry Definitions: **more likely to be asked about on the exam** acid: a proton donor base: a proton acceptor Amphoteric substance: can act as an acid or a base depending on how it acts in the reaction. Ex) water is an amphoteric substance. Conjugate Pairs: The product that is formed when a proton is donated or accepted from the reactants. Acids have conjugate Bases. Bases have Conjugate Acids. *the stronger the acid, the weaker the conjugate base and vice versa*15.4: Acid Strength and Acid Ionization Constant (Ka): A strong acid: completely ionizes in solution HA(aq)+H2OH3O+(aq)+A-(aq) A weak acid: partially ionizes in solution HA(aq)+H2OH3O+(aq)+A-(aq) *note the difference in arrows* Acids can be classified according to their number of protons: Monoproctic acid: HA Diproctic acid: H2A Triproctic acid: H3A Ka: a measure of the strength of the acid. The larger the Ka value, the stronger the acid. 15.5: Autoionization of Water and pH: Autoionization: When water acts as an acid and a base with itself H2O(l)+H2O(l)H3O+(aq)+OH-(aq) Kw: Equilibrium Constant for H2O and the Disassociation constant for water *the [H3O+]*[OH-] will always = 1.0*10^-14* An acidic solution: contains a base that creates additional H3O+ ions, causing [H3O+] to increase and [OH-] to decrease. A basic solution: contains a base that creates additional OH- ions, causing [OH-] to increase and [H3O+] to decreaseThe pH Scale: Ranges from 014 06.9999 is acidic 7.0 is neutral 7.114 is basic pH=-log[H3O+] The pOH and other p scales: pOH=-log[OH-] pOH is the opposite of pH, thus: pH+pOH=14 pKa=-log(Ka) pKb=-log)Kb) 15.6: Finding the [H3O+] & pH of Strong and Weak Acid SolutionsStrong Acid complete ionization (disassociation) Weak Acid Only partially ionized. must use ICE table to determine []’s **to determine if you can disregard the –x in the denominator of Ka equations: step 1: multiply given Ka value by 100. step 2: compare this new number to the original concentration given to you step 3: if the concentration is greater than the Ka*100, you can disregard the x Note: Pretty much every practice problem she’s given us you can throw out the X, so if you’re not sure if you can or not, odds are you can!** Percent Ionization of a Weak Acid: % ionization: ( [H3O+]aq/[HA]initial ) *100A Mixture of Two Weak Acids: *look for the strongest relative to the rest of the acids: the acid with the larger Ka, then use THAT Ka value to solve for the pH b/c it is the one that will have the most effect*15.7: Base Solutions Strong Base: completely disassociates in solution Weak Base: partially disassociates in solution 15.8: The Acid and Base Proportion of Ions and Salts: Anions as Weak Bases: Depends on the strength of the acid it comes from. Cl-(conjugate base of a weak acid) HC2H3O2 (weak acid) * an anion that is the conjugate base of a weak acid is itself a weak acid* C2H3OH2- is a weak base and will cause an increase in pH of solution NO3- HNO3(strong acid) :: NO3- is pH neutral because it’s the conjugate base of a strong acid NO2- HNO3 (weak acid) Ka=4.6*10^-4 :: NO2- is a weak base because it’s the conjugate of a weak acid.Ka*Kb=Kw Cations as Weak Acids: Cations that are the counterions of strong bases are themselves pH neutral. Cations that are conjugate acids of weak bases are weak acids Na+NaOH(strong base):: Na+ is the counter ion of a strong base so it’s pH neutral H4+NH3 (Kb=1.76*10^-5):: H4+ is the conjugate acid of a weak base so it is a weak acid Classifying Salt Solutions as Acidic, Basic, or Neutral: 1) The cation is pH neutral and the anion is pH neutral salt forms neutral solution 2) The cation is pH neutral and the anion is a weak base salt forms a basic solution 3) The cation is a weak acid and the anion is pH neutral salt forms an acidic solution 4) The cation is a weak acid and the anion is a weak base must compare the Ka and Kb values. The stronger species determines the pH15.9: Polyprotic Acids: Ionize in successive steps, each with its own Ka value. Use the step with the strongest Ka value to determine the pH 15.10: Acid Strength and Molecular Structure Binary Acid: H—Y : formed of hydrogen bonded to some other atom 2 factors that determine its strength: bond polarity and bond strength 1) Bond Polarity Bond between H and atom should be polar for the hydrogen atom to be acidic 2) Bond Strength: The weaker the H—Y bond, the stronger the acid is. H—F: Bond energy: 565 kJ/mol (weak acid) H—Cl: 431 kJ/mol (strong acid)* as you go leftright across the periodic table, the acidity increases. *as you go down the columns of the periodic table, the acidity increases Oxyacid: Contains a hydrogen atom bonded to an oxygen atom which in turn is bonded to another atom Acidity of oxyacids depend on electronegativity of Y: Y more electronegativestronger oxyacid The number of oxygen atoms bonded to Y more oxygen atomsstronger oxyacid 15.11: Lewis Acids and Bases Lewis Acid: electron pair acceptor Lewis Base: electron pair donor Chapter 16: Aqueous Ionic Equilibrium 16.2: Buffers Buffer: solutions that resist pH change A buffer solution: prepared from a weak acid and its conjugate base. Acid/Base pair is chosen with pKa value close to the required pH Calculate the pH of a buffer: Henderson-Hasselbach equation: (we had to use this for
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