Acids and Bases

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Acids and Bases

• Chp 16

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Old Definitions

• Classic
• Acids taste sour
• Bases taste bitter
• Arrhenius model
• Acids produce hydronium ions (H3O) in solution
• Bases produce hydroxide ions (OH-) in solution

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Current Definition

• Bronsted Lowry Model
• Acids donate protons
• Bases accept protons

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Conjugate Acid/Base Pairs

• When an acid donates a proton, it forms a
  conjugate base
• HCl donates and becomes Cl-
• When a base accepts a proton, it forms a
  conjugate acid
• OH- accepts and becomes H(OH)
• These differ by only ONE hydrogen atom
• If the acid is strong, its conjugate base is
  weak (and vice versa)

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Acid-Base Reactions

• Can occur in either direction, but both sides
  compete for the free H ions so 1 direction often
  dominates
• HCl H2O H3O Cl-
• Acid base conj. acid conj. base

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Some Terms

• Amphoteric
• can be acids or bases, depending on the
  circumstance
• Would be weak in both cases
• Water is the most common amphoteric substance
• Buffer
• A weak acid or base that can be added to a
  solution to help it resist changes in pH

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Indicators

• Substances that change color when exposed to an
  acid or a base
• Occurs because of reactions with the hydrogen
  ions
• Many substances are natural pH indicators (ex.
  Blueberries, red cabbage, ammonia)
• Some are commonly used in labs
• Litmus paper (bases turn red paper blue, acids
  turn blue paper red)
• Phenalthalien turns red when basic

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Concentrations

• Can be measured with molarity, but can be
  misleading because some acids have multiple H
  ions to donate
• Use Normality instead
• N mols (equivalents)
• Liters
• Equivalents how many H ions (or OH ions) that
  acid or base has
• Ex. 0.5 mol H3PO4 dissolved in 1.5 L has what
  normality?
• N 0.5(3) 1 N
• 1.5

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Neutralization

• Adding just enough acid to completely react with
  a base (or vice versa)
• Need to have every proton used
• Creates water and a salt (ionic compound)
• Can determine the amount of acid or base in a
  reaction using this in a process known as
  titration
• Adding a little of a known acid or base until the
  equivalence point is reached as shown by a pH
  indicator
• NacidVacid NbaseVbase

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An Example

• What volume of 0.10 M NaOH is required to
  completely neutralize 50.0 mL of 0.20 M H2SO4?
• 0.10 M NaOH 0.10 (1) 0.10 N NaOH
• 0.20 M H2SO4 0.20 (2) 0.40 N H2SO4
• SO
• 0.40 (50) 0.10 (V)
• V 200 ml

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H and OH- Ions

• Acids create H ions when mixed in water because
  they donate H to it
• Bases create OH- ions when mixed in water because
  they take a H away from it
• If H gt OH-, the solution is acidic
• If H lt OH-, the solution is basic
• If H OH-, the solution is neutral
• However, for any solution H OH- 1 x 10-14

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An example

• If a solution has a OH- 3.4 x 10-9 M, what is
  the concentration of H?
• H (3.4 x 10-9) 1 x 10-14
• H 2.9 x 10-6 M
• Is the solution acid or basic?
• since H gt OH-, it is acidic
• (remember that negative exponents mean more zeros
  after the decimal, so a smaller )

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pH Scale

• An easier way to measure acids and bases because
  the concentrations we work with are so small
• 0-7 acidic
• 7 neutral
• 7-14 basic
• The closer to 7 the weaker the substance is, the
  further away the stronger it is
• Each step is actually a factor of 10
• So pH 2 is 100 X stronger than pH 4

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pH continued

• pH measures the concentration of H ions (or
  hydronium ions) in solution
• pH -log H
• pOH -log OH-
• pH pOH 14
• To get H from pH
• H inverse log (-pH)

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An example

• What is the pH of a solution that has a H of
  1.3 x 10-11M?
• pH -log (1.3 x 10-11) 10.88
• Is it an acid or a base?
• basic since the pH is greater than 7

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Another Example

• What is the H if the pH of a solution is 2.3?
• H inverse log (-2.3) 5.01 x 10-3 M
• What would the pOH of this solution be?
• 2.3 pOH 14
• pOH 11.7