Chemistry Chapter 18

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Chemistry Chapter 18

• Acids and Bases

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18.1 Introduction to Acids and Bases

• Objectives
• 1. Identify the physical and chemical properties
  of acids and bases
• 2. Classify solutions as acidic, basic or neutral
• 3. Compare Arrhenius, Bronsted-Lowry and Lewis
  models of acids and bases

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Physical and Chemical Properties

• Some physical properties of acids are that they
  taste sour, give carbonated beverages a sharp
  taste (carbonic and phosphoric acids), ability
  to conduct electricity
• Physical properties of bases include bitter taste
  (unsweetened chocolate) and slippery feel
• Chemical properties of acids include ability to
  turn blue litmus red
• Red litmus also stays red in the presence of an
  acid

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• Acids also react with metals chemically
• Chemical properties of bases include the ability
  to turn red litmus blue
• Blue litmus stays blue in the presence of a base
• Remember this by BB Bases are Blue
• Acids are hot, so they are red
• All water solutions contain varying amounts of
  hydrogen ions (H) and hydroxide ions (OH-)
• The relative amounts of each determine whether
  you have an acid or a base

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Acidic, Basic and Neutral

• An acidic solution contains more H than OH-
• A basic solution contains more OH- than H
• Water produces H and OH- by a process called
  self-ionization
• It looks like this H2O(l)H2O(l)?H3O(aq)OH-(aq)
  
• The hydronium ion is H3O which has a water
  molecule attached to it by a covalent bond
• Since H and H3O are used interchangeably,
  self-ionization also looks like this
  H2O(l) ?H(aq)OH-(aq)
• The equilibrium constant for water
  self-ionization of water is called Kw and it is
  1.0 x 10-14M

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Arrhenius

• According to Arrhenius an acid is a substance
  that contains hydrogen and produces hydrogen ions
  in aqueous solution
• A base is a substance that contains a hydroxide
  group and dissociates to produce hydroxide ions
  in aqueous solution
• This definition has some shortcomings
• Ammonia (NH3) and sodium carbonate (Na2CO3) are
  well-known bases that do not qualify as bases
  under this definition

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Bronsted-Lowry

• According to Bronsted-Lowry, an acid is a
  hydrogen ion (proton) donor
• A base is a hydrogen ion acceptor
• Under this definition conjugate acids and
  conjugate bases are produced
• This definition includes all the Arrhenius acids
  and bases, but it still excludes ammonia
• With Bronsted-Lowry, water is amphoteric it can
  act as both an acid and a base

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Lewis

• In the Lewis model, an acid is an electron-pair
  acceptor while a base is an electron-pair donor
• The Lewis models includes all of Arrhenius, all
  of Bronsted-Lowry and many more including NH3
• This models allows compounds that completely lack
  H and OH- to be classified
• Boron trifluoride and gaseous ammonia are an
  acid-base reaction

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Naming Practice (Remember These??)

• Name each acid or base
• A) HF
• B) HNO3
• C) KOH
• D) H2SO4
• E) H3PO4
• F) NH3
• G) H2CO3

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18.2 Strengths of Acids and Bases

• Objectives
• 1. Relate the strength of an acid or base to its
  degree of ionization
• 2. Compare the strength of a weak acid with the
  strength of its conjugate base
• 3. Explain the relationship between the strengths
  of acids and bases and the values of the their
  ionization constants

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Strength Degree of Ionization

• Def a monoprotic acid is an acid that can donate
  one hydrogen ion
• Def a polyprotic acid is an acid that can donate
  more than one hydrogen ion (such as diprotic,
  triprotic, etc.)
• If the bonds between the hydrogen and other
  element are polar, the hydrogen ion can be
  donated (ex HF, H2SO4, HC2H3O2, HCl)
• If the bonds are nonpolar, the substance will not
  be an acid, hydrogen ion cannot be donated
  (nonpolarC6H6-benzene)

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• Some acids have the ability to ionize completely,
  while others only partially ionize depending on
  bond strength
• Def a strong acid is an acid that ionizes
  completely in aqueous solution
• Notice that the word strong does not define the
  concentration, merely the ability to ionize
• Def a weak acid is an acid that partially
  ionizes in aqueous solution
• Weak acids do not conduct electricity as well as
  the strong acids
• The more ions in aqueous solution, the better the
  conductor will be

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Strength of Weak Acids with Their Conjugate Bases

• Strong acids will have weak conjugate bases
• This is because the forward reaction toward
  complete ionization is stronger than the reverse
  reaction of conjugate base producing an acid
  again
• The equilibrium for a weak acid lies to the left
  (on the reactant side) because the conjugate base
  (on the product side) has a greater attraction
  for H from the conjugate acid on the product
  side than does the base on the reactant side

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Notice that these expressions have water as a
reactant
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Ionization Constants Acid/Base Strength

• Def Ka is the acid ionization constant which is
  the value of the equilibrium constant expression
  for the ionization of a weak acid
• Polyprotic acids has a unique Ka value for each
  ionization step, and the values decrease
  successively
• The weakest acids have the smallest Ka values
• Ka is calculated using the expression
  products/reactants, but excluding water

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• Def a strong base is a base that dissociates
  entirely into metal ions and hydroxide ions in
  aqueous solutions
• Def a weak base ionizes partially in aqueous
  solutions
• Kb is the base ionization constant
• Kb is the value of the equilibrium constant for
  the ionization a base
• Calculation is products/reactants it
  includes OH- (in the numerator)

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18.3 Hydrogen Ions and pH

• Objectives
• 1. Explain pH and pOH
• 2. Relate pH and pOH to the ion product constant
  for water
• 3. Calculate pH and pOH of aqueous solutions

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pH and pOH

• pH and pOH are logarithmic scales that express
  the concentration of hydrogen ions and hydroxide
  ions in aqueous solutions
• Recall the self-ionization of water
  H2O(l) ?H(aq)OH-(aq)
• The concentration of pure water is constant for
  the equilibrium constant will not have H2O in
  the denominator
• It looks like this KwHOH-

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Ion Product Constant for Water

• At 298K, H and OH- are equal to each other
  and have been scientifically measured in pure
  water to be 1.0 x 10-7 M each
• Therefore, Kw 1.0 x 10-7 M X 1.0 x 10-7 M or 1.0
  x 10-14 M2
• pH and pOH exist along a continuum of a scale
  that ranges from 0 14
• The lower the pH, the more acidic the substance
  is
• 7 is neutral pH values above 7 are basic (also
  called alkaline)

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Calculations

• Conversely, low pOH values are alkaline while
  high pOH values are acidic
• Three formulas to know and use
• 1. pH -logH (ALSO pOH -logOH-)
• 2. pH pOH 14
• 3. H x OH- 1.0 X 10-14 Kw
• A couple of hints as we begin to do calculations
  log x antilog1
• Neutral pH -log(1 x 10-7) 7

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Hints on Calculations

• To get H from the pH, take the negative
  antilog of the pH
• Look at book example problems on pages 653-657
• Q If the hydroxide ion concentration of an
  aqueous solution is 1.0 x 10-3 M, what is H
  in the solution? Is the solution acidic, basic or
  neutral?
• A 1.0 x 10-14 M, basic

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18. 4 Neutralization

• Objectives
• 1. Write chemical equations for neutralization
  reactions
• 2. Explain how neutralization reactions are used
  in acid-base titrations

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Write Equations

• In a neutralization reaction, an acid reacts with
  a base to produce a salt and water
• Def a salt is an ionic compound made up a cation
  from a base and an anion from an acid
• Most salts are composed a combination of a metal
  plus a nonmetal
• When writing neutralization reactions, you must
  know whether all of the reactant and products in
  solution exist as molecules (covalent bonding) or
  formula units (ions)
• For strong acids and strong bases, the net ionic
  equation is acceptable

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Titrations

• Def a titration is a method for determining the
  concentration of a solution by reacting a known
  volume of that solution with a solution of know
  concentration
• Def the titrant is the solution of known
  concentration, also called the standard solution
• Def the equivalence point is the point at which
  moles of H equal OH-

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• Def an acid-base indicator is a chemical dye
  whose colors are affected by the presence of an
  acid or a base
• Litmus is one example of a solid paper indicator
• Many indicators are liquid such as
  phenolphthalein which is pink in the presence of
  a base and is colorless in acids
• The equivalence point must be distinguished from
  the end point
• Def the end point is the point at which the
  indicator used in a titration changes color

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• To titrate
• 1. a measured volume of acid of unknown
  concentration is added to flask
• 2. several drops of indicator (often
  phenolphthalein) is added
• 3. a base of known concentration is slowly added
  through a device known as a buret until the
  indicator changes color
• HINT you must get the exact point or one drop
  where the color change occurs in order to get the
  correct volume of base added
• 4. use m1v1m2v2 to calculate the concentration
  of the acid