Acids and Bases

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Acids and Bases

• Chapter 16 and Chapter 17 (part)

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Arrhenius Definition

• Acid - produces H in solution
• Base - produces OH- in solution
• Best definition - any substance that raises the
  H concentration (or OH- concentration) of water
• Examples HCl ? H(aq) Cl-(aq)
• NaOH ? Na(aq) OH-(aq)

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Bronsted-Lowry Definition

• Arrhenius deffinition was too narrow - only
  solutions and some substances didnt fit.
• Water is more active in this definition
• H H2O ? H3O
• B-L Acid Substance that can donate a proton to
  another substance
• B-L Base Substance that can receive a proton
  from another substance
• Pairs! HCl H2O ? Cl- H3O
• BLA BLB CB CA

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Broader Definition

• Doesnt have to be in solution
• Vapor (from video)
• HCl(g) NH3(g) ? Cl- NH4
• To be a B-L acid, a molecule or ion must have a
  hydrogen that it can lose
• To be a B-L base, a molecule or ion must have an
  unshared pair of electrons that it can use to
  bind the hydrogen (look at NH3)
• Some substances can act as both - amphoteric (H2O)

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Relative Strength of Acids and Bases

• The stronger the acid, the weaker its conjugate
  base. (overhead)
• Three categories
• Strong acids completely donate their proton to
  water - name them
• Weak acids partly dissociate to form a few
  conjugate ions weak bases have slight ability to
  remove protons from water
• Negligible acidity like CH4, molecule has
  hydrogen but doesnt show any acidic behavior.

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Equilibrium

• In every acid-base reaction, the position of the
  equilibrium favors transfer of the proton to the
  stronger base.
• HCl H2O ? H3O Cl-
• Which are the two bases above? Which is
  stronger? Called Battle of the bases.

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Autoionization of Water

• H2O H2O ? H3O OH-
• Two in 10 billion molecules do this.
• Keq H3OOH-
• Called Kw
• 1 X 10-14 at 25? C
• Very important - this value stays constant for
  most aqueous solutions at 25? C

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More

• Concentration of each ion is 1.0 X 10-7 in a
  neutral solution.
• When you know one concentration, you can
  calculate the other.
• PE page 621

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The pH Scale

• Too difficult to use negative exponents
• It is the negative log in base 10 of the H
  concentration
• pH -logH
• Try your calculator, both directions
• pOH is the same but for OH- concentration

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pH of Common Substances
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Measuring pH

• pH meter uses small changes in voltage caused by
  H changes to determine pH.
• Indicators - substances which change to specific
  colors depending upon pH.
• It is actually an equilibrium shift.
• HA H2O ? H3O A-

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Strong Acids and Bases

• Name them!
• Complete ionization or dissociation.
• HNO3 ? H NO3- (ONE WAY!)
• Easy to calculate pH
• It is the molarity that you begin with
• Practice

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Weak Acids

• HA ? H A-
• Write the Keq expression for this
• This is called Ka
• Would a stronger weak acid have a larger or
  smaller Ka?
• Equilibrium is reached simultaneously, so pH will
  give H which will give Ka
• PE

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Using Ka to Calculate pH

• RICE with unknown x
• What is the pH of a solution of 0.30 M acetic
  acid? Try this.
• Use table to get Ka
• This would lead to quadratic
• Assume it is negligible to original amount for
  denominator
• Use the 5 rule! Ignore if x is less than 5
  of original amount.

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Polyprotic Acids

• Use Ka1 to estimate pH. It is always much
  greater than the others.
• H2SO4 ? H HSO4-
• HSO4- ? H SO42-
• Ka 1 (completely ionizes) Ka 2
• 1.2 X 10-2

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Weak Bases

• The same as weak acids, except you are getting
  pOH and then must get pH
• Called Kb
• Categories
• Neutral substance with a nonbonding pair of
  electrons that can serve as a proton acceptor,
  usually contain N ammonia and amines (NH2)
• Anions of weak acids F-, CH3COO-

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Relationship Between Ka and Kb

• Write each equation for the NH3/NH4 conjugate
  pair
• Add them algebraically
• This means that the constants are multiplied
• Calculate Ka X Kb
• This is Kw!
• Ka or Kb may be listed as pKa.
• pKa pKb 14

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Acid-Base Properties of Salts

• Assume 100 dissociation at low concentrations
• Ions may ionize (react with) water to yield H3O
  or OH-
• Lets look at the anion in a salt which forms as
  the conjugate to an acid. (Cl- or CH3COO-)
• Only one of these will cause water to lose an H
  and become OH-. Which one and why?

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Cations

• Result from bases
• Stong bases (Na) neutral
• Weak base (NH4) acidic
• Summary Cations or anions from strong acids and
  bases form neutral salts. Learn these and figure
  out the rest.
• Look at summary on page 642
• Metal cations cause solutions to become slightly
  acidic (16-11) READ!
• Practice NaCl, KCN, NH4Br

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Structure and Acid-Base Behavior

• Read 644-652.

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Chapter 17 Additional Aspects

• Common Ion Effect
• CH3COOH ? CH3COO- H
• What does LeChatelier say will happen when
  NaCH3COO is added?
• Therefore, pH goes up!
• The extent of ionization of a weak electrolyte is
  decreased by adding to the solution a strong
  electrolyte that has a common ion.
• PE

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Buffers

• This is a solution that resists changes in pH.
• Very important in nature (blood pH 7.4)
• Prepared by adding a weak acid or base and the
  salt of its conjugate
• Example NH4Cl to NH3
• We can choose the right components to buffer at
  any pH.

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Equation Derivation

• HA ? H A-
• Ka H A-
• HA
• H Ka HA
• A-
• Adding acid or base causes little change in
  ratio because of equilibrium

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Half-equivalence point

• Half way to the stoichiometric equivalence point,
  HA A-
• What happens to equation
• pKa pH

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Choosing a Buffer

• Pick a buffer whose acid has a pKa near the pH
  that you want
• Must consider pH and capacity
• Capacity amount of acid or base a buffer can
  neutralize before changing appreciably.
• .1M/.1M has 1/10 the buffering capacity of 1M/1M

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Henderson-Hasslbalch Equation

• Put previous equation in terms of pH
• pH pKa log base
• acid
• Try PE with new equation

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Buffer Problem

• Look at ratio in pH pKa log base
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  acid
• How would adding HCl affect it?
• It would decrease the base and add to the acid.
• How would adding NaOH affect it.
• It would decrease the acid and add to the base.
• PE

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Excess OH- or H

• Subtract the ion from the acid or base as
  appropriate. When in reaches 0, the remainder
  will be in excess.
• Now calculate the pH based upon the excess that
  is in the solution as pure OH- or H
• For example, if there is only 0.30 moles of weak
  base and you add 0.32 moles of H, there will be
  0.02 moles of H in solution.

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Titration Intro

• Quantitative process which uses a solution of
  known concentration to precisely determine the
  concentration of a species in a second solution
  usually acid/base
• Equivalence point stoichiometry says that equal
  amounts of acid and base are present
• End point indicator changes color, should be
  very near the equivalence point

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Monitor pH throughout titration

• Handouts
• Know general characteristics of each
• Use common sense
• Dont mix moles and molarity
• Stoichiometry is based upon moles, pH is based
  upon molarity

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Strong Acid and Strong Base

• 1. Initial pH determined by initial
  concentration of the strong acid (base)
• 2. Between initial and equivalence point pH
  change is slow and then rapid as it approaches it
  and leaves it.
• 3. pH at equivalence point is 7. Why?
• 4. After equivalence point pH is determined by
  excess base (or acid).

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Titration Curve
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Major Plan

• Find moles of unreacted acid (base) by
  stoichiometry
• Find molarity using new volume
• Find pH
• Practice

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Weak Acid- Strong Base Titration

• 1. Initial pH is a RICE problem using Ka.
• 2. Calculate the total number of moles of weak
  acid, and amount of base necessary to neutralize
  this. (Road map).
• 2. Before equivalence point change is gradual.
  Determine pH by using stoichiometry to get acid
  and conjugate base ratio, then use
  Henerson-Hasselbalch.
• 3. pH at equivalence point RICE with conjugate
  base that has been formed
• 4. pH afterwards increases rapidly with excess
  OH-

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Weak Acid- Strong Base Curve
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Strong Acid-Weak Base

• Handout for curve
• Mirror image
• Same math
• Try several of each!

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Strong Acid-Weak Base Curve
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Polyprotic Acids

• Multiple Equivalence Points
• Calculations done with first Ka

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Solution Equilibrium

• Saturated solution dissolved solute is in
  equilibrium with undissolved solute
• Rules tell us what will dissolve, Ksp tells us
  how much.
• BaSO4(s) ? Ba2(aq) SO42-(aq)
• Ksp Ba2 SO42-
• Ksp is called the solubility constant

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Differences

• Solubility is expressed in g/l or mol/l
• Solubility changes, Ksp does not!
• Can get solubility from Ksp and vice versa
• PE

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Common Ion and Solubility

• Common ion the solubility of a slightly soluble
  salt is decreased by the presence of a second
  solute that furnishes a common ion.
• This becomes a RICE problem with an initial
  amount of an ion.

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pH and Solubility

• The solubility of almost any ionic compound is
  affected if the solution is made sufficiently
  acidic or basic.
• The effects are noticeable and can be calculated
  if one or more ions are acidic or basic.
• General Rule The solubility of slightly soluble
  salts containing basic anions increases as pH
  goes down.
• CaF2 ? Ca2 2F-
• H F- ? HF
• CaF2 2H ? Ca2 2HF

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Precipitation and Separation of Ions

• Q problem what is Q?
• Add two ions from two salts that might
  precipitate
• Example BaCl2 and Na2SO4
• Find Ksp for BaSO4
• Calculate Q
• Q gt Ksp what will happen?