Reactions in Solution and Titration

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Coral
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46465
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Reactions in Solution and Titration
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2010-11-01 12:20:35
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che130 combustion oxidation redox chemistry
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For CHE130
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  1. A homogeneous mixture of two or more substances.
    Solution
  2. A solution in which water acts as the solvent.
    Aqueous solution (aq)
  3. A solution that contains a small amount of solute relative to the solvent
    Dilute solution
  4. A solution that contains a large amount of solute relative to the solvent.
    Concentrated solution
  5. Molarity (M)
    Unit to express concentration. It is the amount of solute per liter of solution: Solute (in mol)/Solvet (in L).
  6. Solution Dilution
    M1V1 = M2V2, so V1=M2V2/M1 where M1 and V1 are the initial molarity and volume of the concentrated solution and M2 and V2 are the final diluted solution's molarity and volume.
  7. Solute-Solute and Solvent-Solute Interactions
    When a solid is put into a liquid solvent, the attractive forces that hold the solid together come into competition with the attractive forces between the solvent molecules and the particles that ompose the solid.
  8. Electrolytes
    Substances that dissolve in water to form solutions that conduct electricity.
  9. Strong electrolytes
    • Substances that completely dissociate into ions when they dissolve in water. The solutions which result conduct electricity. They tend to be composed of ionic compounds.
    • ex: Sodium chloride
  10. Nonelectrolytes
    Compounds that do not dissociate into ions when dissolved in water. The solutions which result do not conduct electricity. They tend to be composed of molecular compounds.`
  11. Acids
    Molecular compounds that ionize (form ions) when they dissolve in water. They produce H+ ions in aqueous solutions.
  12. Strong Acid
    • A molecular compound that completely ionizes in solution. They are equally strong electrolytes.
    • ex: Hydrochloric acid (HCl)
  13. Weak Acids
    • Acids that do not completely ionize in water. They are weak in conducting electricity and thus make poor electrolytes.
    • ex: Acetic acid (HC2H3O2)
  14. Soluable compound
    A compound which dissolves in water.
  15. Insoluable compound
    A compound which does not dissolve in water.
  16. Compounds Containing the Following Are Generally Soluable in Water
    • Li+, Na+, K+, NH4+, NO3-, C2H3O2-
    • Cl, Br-, and I - (except when paired with Ag+ Hg22+, or Pb2+)
    • SO42- (except when paired with Sr2+, Ba2+, Pb2+, Ag+, or Ca2+)
  17. Compounds Containing The Following Ions Are Generally Insoluable
    • OH- and S2-(except when paired with Li+, Na+, K+, NH4+, Ca2+, Sr2+, or Ba 2+, wherein OH- is soluable and S2- is slightly so)
    • CO32- and PO43- (except when paired with Li+, Na+, K+, NH4+)
  18. Soluability Rules
    • 1) Compounds containing the sodium ion are soluable.
    • 2) Compounds containing the NO3- ion are soluable
    • 3) When compounds containing polyatomic ions dissolve, the polyatomic ions remain as intact units as they dissolve
    • 4) With some exceptions, compounds containing the CO32- ion are insoluable.
  19. Precipitate from
    Come out from
  20. Precipitation Reactions
    Reactions in which a solid or precipitate forms upon mixing to solutions: Soluable + Soluable -> Soluable + Insoluable. Only insoluable compounds form precipitates.
  21. Procedure For Writing Equations For Precipitation Reactions
    • 1) Write the formulas of the two ompounds being mixed as reactants in a chemical equation.
    • 2) Below them,w rite the formula of the products that could form from the rectants. Combine the cation frm each reactant with the anion from the other.
    • 3) Use the solubility rules to determine whether any of the possible products are insoluable.
    • 4) If all the possible products are soluable, there will be no precipitate. Write NO REACTION.
    • 5) If any possible products are insoluable, write their formula as the products of the reaction.
    • 6) Balance the equation.
  22. Molecular Equation
    An equation showing the complete neutral formula for each compound in the reaction as if they existed as molecules.
  23. Spectator Ions
    Ions that do not participate in the reaction occuring in an ionic equation.
  24. Net ionic equation
    Equations which show on ly the species that actually change during the reaction.
  25. Complete ionic equation
    Equations which list individually all of the ions present as either reactants or products in a chemical reaction. These show all of the speies as they are actually present in solution.
  26. Acid-Base Reaction (aka Neutralization Reaction)
    An acid reacts with a base and the two neutralize each other, producing water or, rarely, a weak electrolyte. The H+(aq) from the acid combines with the OH- (aq) from the base to form H2O (l). They generally form a salt that remains dissolved in the solution.
  27. Gas-Evolution Reaction
    • Aqueous reactionns that form a gas when two solutions are mixed, resulting in bubbling. Some form a gaseous product directly when the cation of one reactant combines with the anion of another. Other form an intermediate product that then decomposes into a gas.
    • Many acid-base reactions are gas-evolution reactions.
  28. Base
    Substances that produces OH ions in aqueous solution.
  29. Hydronium ions
    H+ ion is a bare poton. Protons associate with water molecules in solution to form hydronium ions. H+ (aq) and H3O+ (aq) are used interchangeably.
  30. Polyprotic acids
    Acids that contain more than one ionized proton and release them sequentially
  31. Diprotic acid
    • A polyprotic acid which is strong in its first ionizable proton but weak in its second.
    • ex) Sulfuric acid
    • H2SO4 (aq) -> H+ (aq) + HSO4-
  32. Common Acids
    • Hydrochloric acid (HCl)
    • Hydrobromic acid (HBr)
    • Hydroiodic acid (HI)
    • Nitric acid (HNO3)
    • Sulfuric acid (H2SO4)
    • Perchloric acid (HClO4)
    • Acetic acid (HC2H3O2) *
    • Hydrofluoric acid (HF) *

    * = weak
  33. Common Bases
    • Sodium hydroxide (NaOH)
    • Lithium hydroxide (LiOH)
    • Potassium hydroxide (KOH)
    • Calcium hydroxide (Ca(OH)2)
    • Barium hydroxide (Ba(OH)2)
    • Ammonia (NH3) *

    * weak base. Ammonia does not contain OH-, but produces it in a reaction with water.
  34. Net Ionic Equation for Acid-Base Reaction
    • H+(aq) + OH-(aq) -> H2O(l)
    • Acid + Base -> Water + Salt (acid-base reactions)
  35. Salt
    Acid-base reactions generally form water and an ionic compound
  36. Decompose
    Break down into component elements
  37. Types Of Compounds That Undergo Gas-Evolution Reactions
    • Sulfides
    • No intermediate product
    • Gas evolved: H2S

    • Carbonates and bicarbonates
    • Intermediate product: H2CO3
    • Gas evolved: CO2

    • Sulfites and bisulfites
    • Intermediate product: H2SO3
    • Gas evolved: SO2

    • Ammonium
    • Intermediate product: NH4OH
    • Gas evolved: NH3
  38. Oxidation-Reduction Reactions (aka Redox Reactions)
    • Reactions in which electrons are transferred from one reactant to the other. Many involve the reaction of a substance with oxygen, but do not necessarily need it.
    • ex) A metal (which has a tendency to lose electrons) reacts wit a nonmetal (which has a tendency to gain). Metal atoms lose electrons to nonmetal atoms.
  39. Oxidation
    The loss of electrons
  40. Reduction
    The gain of electrons (and reduction of charge).
  41. Oxidation number (aka Oxidation state)
    A number given to each atom based on the electron assignments. It is the charge it would have if all shared electrons were assigned to the atom with a greater attraction for those electrons.
  42. Rules for Assigning Oxidation States
    • 1. The oxidation state of an atom in a free element is 0.
    • ex) Cu (0 ox state); Cl2 (0 ox state)

    • 2. The oxidation state of a monoatomic ion is equal to its charge.
    • ex) Ca2+ (+2 ox state); Cl- (-1 ox state)

    • 3. The sum of the oxidation states of all atoms in a) a neutral molecule or formula unit is 0 b) an ion is equal to the charge of the ion.
    • ex) H2O 2(H ox state) + 1(O ox state) = 0
    • ex) NO3- 1(N ox state) + 3(O ox state) = -1

    • 4. In their compounds metals have positive oxidation states.
    • ex) Group 1A metals always have an oxidation state of +1.
    • ex) Group 2A metals always have an oxidation state of +2.

    • 5. In their compounds, nonmetals are assigned oxidation states as follows:
    • Fluorine +1, Hydrogen +1, Oxygen -2, Group 7A -1, Group 6A -2, Group 5A -3
    • NOTE: Rules are hierarchical. If two rules conflict, consult the higher rule.
  43. Oxidation
    An increase in oxidation state
  44. Reduction
    A decrease in oxidation state
  45. Oxidizing agent
    A substance that causes the oxidation of another substance. In a redox reaction, the oxidizing agent is always reduced.
  46. Reducing Agent
    A substance that causes he reduction of another substance. In a redox reaction, the reducing agent is always oxidized.
  47. Combustion Reaction
    A type of redox reaction characterized by the reaction of a substance with O2 to form one or more oygen-containing compounds, often including water. They aways emit heat.

    ex) Compounds containing carbon and hydrogen or carbon, hydrogen, and oxygen always form arbon dioxide adn water upon omplete combustion.

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