Chapter 4: Chemical Reactions

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Chapter 4 Chemical Reactions
Chemistry 1061 Principles of Chemistry I Andy
Aspaas, Instructor
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Chemical reactions

• Ions in aqueous solution
• Molecular and ionic equations
• Types of reactions
• Precipitation reactions
• Acid-base reactions
• Oxidation-reduction reactions
• Solutions
• Concentration and dilutions
• Quantitative analysis
• Gravimetric and volumetric analyses

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Ions in aqueous solutions

• Ionic theory of solutions Arrhenius, 1884
• When dissolved in water, the individual ions of
  ionic substances completely separate and enable
  the solution to conduct electricity
• Pure water is a poor conductor of electricity
• Electrolyte substance that dissolves in water to
  give an electrically conducting solution
• Generally, ionic solids that dissolve in water
  are electrolytes
• A few molecular electrolytes, Ex. HCl (g)
• Nonelectrolytes dissolve in water, poorly
  conducting solution, usually neutral molecular
  substances

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Strong and weak electrolytes

• The extent to which a solution conducts
  electricity indicates the strength of the
  dissolved electrolyte
• Strong electrolytes exist in solution almost
  entirely as ions
• Ex. NaCl
• Weak electrolytes dissolve in water to give only
  a small percentage of dissociated ions
• Ex. NH3

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Solubility rules

• Solubility ability of a substance to dissolve
  completely in water
• Ex. Sugar, NaCl, ethyl alcohol are soluble
• Ex. Calcium carbonate, benzene are insoluble
• Soluble ionic compounds are strong electrolytes
• 8 solubility rules can determine whether an ionic
  compound is soluble or not

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Solubility rules
Li, Na, K, NH4
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Molecular and ionic equations

• Molecular equation chemical equation in which
  reactants and products are written as if they
  were molecular substances, even if they exist as
  ions in solution
• Explicit in the actual compounds added to a
  solution, and the products obtained
• Complete ionic equation all strong electrolytes
  are written as their dissociated ions (aq)
• Insoluble compounds are written as a solid
  compound, not ions

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Net ionic equations

• Spectator ion ion in an ionic equation that does
  not take part in the reaction
• Appears in ionic form on both sides of a reaction
• Net ionic equation equation in which all
  spectator ions have been canceled
• Several different reactions can have the same net
  ionic equation

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Precipitation reactions

• Precipitate insoluble compound formed during a
  chemical reaction in solution
• Predicting precipitation reactions
• Exchange reaction most common, each compound
  trades partners to form products
• Write molecular equation
• Use solubility rules to determine phase lables
  for each product and reactant (aq) if soluble,
  (s) if insoluble
• If all components of reaction are soluble, no
  reaction occurs
• If a product is insoluble, it forms as a
  precipitate
• A net ionic equation shows the reaction at the
  ionic level

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Acid-base reactions

• Acids vinegar (acetic acid), lemon juice (citric
  acid), Coca-Cola (phosphoric acid and carbonic
  acid), battery acid (sulfuric acid)
• Bases Drano (sodium hydroxide), ammonia, Milk of
  Magnesia (magnesium hydroxide)
• Brønsted-Lowry acid molecule or ion that donates
  a proton to another species in a proton-transfer
  reaction
• Brønsted-Lowry base molecule or ion that accepts
  a proton in a proton transfer reaction

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Strong acids and strong bases

• Strong acids and bases ionize completely in water
• Strong acids HClO4, H2SO4, HI, HBr, HCl, HNO3
• Strong bases LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2,
  Ba(OH)2
• Weak acids and bases only partly ionize in water

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Neutralization reactions

• Reaction between acid and base to produce a salt
  and possibly water
• Salt ionic compound formed in neutralization
  reaction
• Start by writing molecular equation
• Acid anion and base cation form the salt
• Water is usually a product
• Net ionic equation write any strong acid or base
  as its dissociated ions

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Acid-base reactions with gas formation

• Carbonates (CO32-) form H2O and CO2 when reacted
  with acids
• Sufites (SO32-) form H2O and SO2 when reacted
  with acids
• Sulfides (S2-) form H2S when reacted with acids

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Oxidation-reduction reactions

• Oxidation-reduction reactions (redox) involve
  transfer of electrons
• Oxidation number actual charge of an atom if it
  exists as a monatomic ion, or a hypothetical
  charge assigned by a few rules
• Elemental atoms always have ox. 0
• Oxygen is usually -2
• Hydrogen is usually 1
• Halogens usually -1 (unless bonded to another
  halogen or oxygen)
• Sum of ox. s of atoms in a compound is 0, sum
  of ox s in a polyatomic ion is the charge on
  the ion

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Describing oxidation-reduction reactions

• If a species loses electrons, it is oxidized
• If a species gains electrions, it is reduced
• LEO, GER
• Use oxidation numbers to determine this
• Oxidizing agent species that oxidizes another
  species, and is itself reduced
• Reducing agent species that reduces another
  species, and is itself oxidized

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Combustion reaction

• Reaction in which a substance reacts with oxygen,
  usually accompanied by release of heat and
  production of a flame
• Organic compounds combust to form CO2 and H2O
• Metals combust to form metal oxides

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Molar concentration

• Molarity measure of concentration
• (moles of solute / liters of solution)
• Unit mol/L
• Diluting solutions MiVi MfVf

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Gravimetric analysis

• Determination of amount of a species by
  precipitating that species out as an insoluble
  compound, and weighing the product
• Mass precipitated product ? moles product ? moles
  unknown species ? mass unknown species

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Volumetric analysis

• Titration method for determining amount of one
  substance by adding a precise volume of another
  substance until the two substances completely
  react
• Colored pH indicator often used to detect
  endpoint
• Volume added solution ? moles added solution ?
  moles unkn. solution ? molarity or grams unkn.
  solution