AP Chapter 4 Notes1 by HC120808193043

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									           Chapter 4
Aqueous Reactions and Solution
        Stoichiometry
General Properties of Aqueous
          Solutions
Ionic Compounds in Water
Molecular Compounds in Water
  Strong and Weak Electrolytes

• Simulation
• Electrolyte activity
Precipitation Reactions
Solubility Guidelines
Double Displacement (Exchange)
           Reactions
Ionic Equations
           Acid-Base Reactions
• Acids and Bases are also common electrolytes, which
  dissociate in water (soluble/separate)
• Acids are substances that ionize in aqueous solutions to
  form hydrogen ions (H+) thereby increasing the
  concentration of H+(aq) ions.
• Because hydrogen atoms consist of only one proton and
  one electron, is H+ simply one proton. Thus acids are
  often called proton donors.
• Just like the ionic compounds we have been studying,
  water also causes the H+(aq) ion to separate from its other
  half.
• Molecules of different acids can ionize to form different
  numbers of H+(aq) ions. Both HCl and HNO3 are
  monoprotic acids yeilding only one H+(aq) ion. H2SO4,
  however, is a diprotic acid that yields two H+(aq) ions. This
  occurs to two steps however.
      Acid-Base Reactions cont.
• Bases are substances that
  accept (react with) H+(aq) ions.
  Bases produce hydroxide ions
  (OH-) when they dissolve in
  water. Ionic hydroxide
  compounds such as NaOH,
  KOH, and Ca(OH)2 are among
  the most common bases,
  separating just as the ions we
  have been studying separate.
• Compounds that do not contain
  OH- ions can also be bases. For
  example, ammonia is a common
  base that when added to water,
  it accepts an H+ from the water
  and thereby produces an OH-
  ion.
Strong and Weak Acids and Bases
• Acid and bases that completely dissociate in
  solution (strong electrolytes) are strong acids
  and bases. (pg 130 in AP Text)
• Those that only partially dissociate in solution
  (weak electrolytes) are weak acid and bases.
• Not only strong acids are highly reactive
  though…
• So strong electrolytes include all soluble ionic
  compounds and strong acids. Weak electrolytes
  include weak acids and bases.
Neutralization Reactions and Salts
• When a solution of a strong acid and strong
  base are mixed a neutralization reaction occurs
  because the H+ the acid combines with the OH-
  from the base to make water and a salt (ionic
  compound whose anion comes from an acid and
  whose cation comes from an base.
      HCl + NaOH H2O + NaCl
• Video simulation
• Also have net ionic equations…
• Sample Exercise 4.7 pg 133
   Acid-Base Reactions with Gas
            Formation
• Many bases beside OH- react with H+ to
  for molecular compounds such as the
  sulfide ion and the carbonate ion. Both of
  these ions reaction with acids to form
  gases. See page 134 for net ionic
  equations.


                            Net Ionic Equation
 Oxidation-Reduction Reactions
• Reactions in which electrons are transferred between
  reactants-redox reactions
• When an atom, ion or molecule has become more
  positively charged by losing electrons it is oxidized
  (oxidation). LEO
• When an atom, ion or molecule has become more
  negatively charged by gaining electrons, it is reduced
  (reduction). GER
• The oxidation of one substance is always accompanied
  by the reduction of another substance. As electrons are
  transferred between them.
• Calcium Water Demo on each table
                 Oxidation Numbers
• We need a system to keep track of whether electrons are lost
  or gained by a substance to determine if a redox reaction has
  taken place.
• Oxidation number is the hypothetical charges assigned to the
  atom before or after a reaction takes place.
• Oxidation numbers are assigned by the following rules:
   – For an atom in it elemental form, the oxidation number is always zero.
   – For any monatomic ion the oxidation number equals the charge on the
     ion.
   – The oxidation number of hydrogen is usually +1 when bonded to
     nonmetals and -1 when bonded to metals.
   – The oxidation number of fluorine is -1 in all compounds, but the other
     halogens have positive oxidation numbers when combined with oxygen.
   – The sum of the oxidation numbers of all atoms in a neutral compound is
     zero and the sum of the oxidation numbers in a polyatomic ion equals
     the charge of the ion.
Oxidation of Metals by Acids and Salts
 • In this chapter we will look at the oxidation of
   metals by acids and salts, but in Ch 20 we will
   look at more complex redox reactions.
 • The reaction of a metal with either an acid or a
   metal salt is a single displacement reaction.
   A + BX  AX + B
 • Also has net ionic reactions…
 • Sample Exercise 4.9 pg 140
           Activity Series of Metals
• Different metals vary in the ease with which they are oxidized
  also known as activity.
• Zn is oxidized by solutions of Cu+2 but Ag is not… meaning that
  Zn loses electrons more readily than Ag. Elemental metals
  become ions by losing electrons through oxidation.
• Therefore, extensive oxidation of metals can lead to the failure of
  metal machinery parts or structures if the parts are made of
  metals with high activity (ease of oxidation).
• A list of metals arranged in order of decreasing ease of oxidation
  is called an activity series (seen in our previous lab).
• Alkali and Alkaline Earth Metals are most easily oxidized,
  elemental form turns into a compound, and therefore are called
  the active metals.
• Transition metals are less active and form compounds less
  readily, called noble metals because of their low reactivity in
  elemental form, and are therefore used for jewelry and coins.
• Any metals on the list can be oxidized by the ions of elements
  below it. Which also means that each metal is reduced by the
  elements above it.
    Concentrations of Solutions
• The behavior of solutions often depends on how much of
  a solute is in it.
• Concentration is the term used to explain the amount of
  solute dissolved in a given quantity of solvent. More
  solute, greater concentration.
• Molarity (M) expresses the concentration of a solution as
  the number of moles of solute in a liter of solution
      Molarity= moles of solute/volume in liters of solution
• Sample Exercises 4.11 on pg 144 and practice exercise
  pg 145.
• Three quantities needed for the calculation so if we know
  two of the three and then we can figure out the third.
• Sample Exercises 4.13 on pg 146
                 Dilutions
• Solutions that are used routinely in the lab
  are often purchased in concentrated form.
  HCl for example is purchased as a 12M
  solution. Solutions of lower concentrations
  can them be obtained by adding water,
  process called dilution.
• Mconc X Vconc =Mdil X Vdil
• Sample Exercise 4.14 pg 148
         Solution Stoichiometry
• If you know the chemical equation and the amount of
  one reactant consumed in the reaction, you can
  calculate the quantities of the other reactants and
  products.
• Coefficients in a balanced equation give the relative
  number of moles of reactants and products. So all values
  given should be converted to moles so that the mole
  ratio in the equation can be used to do conversion from
  one reactant to another part of the equation.
• If you have mass convert into moles before you convert
  to mass of other substance. If you have molarity, convert
  into moles as well before you convert to molarity of other
  substance. See chart on pg 149.
                  Titrations
• To determine concentration of a particular solute in
  a solution, chemists often carry out a titration,
  which involves combining a sample of the solution
  with a reagent solution of known concentration
  called a standard solution.
• Can be conducted using acid-base, precipitation or
  oxidation-reduction reactions.
• The point at which stoichiometrically equivalent
  quantities are brought together is known as the
  equivalence point for the titration (ratios from
  balanced equation achieved).
• Easy to determine the equivalence point for acid
  and bases (solution becomes neutral and can be
  detected by indicators such as phenothalein)
• Practice Exercises on pgs 152-153.

								
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