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Aqueous Reactions and Solution Stoichiometry

Aqueous Reactions and Solution Stoichiometry. Water based chemistry. Water, water everywhere and not a drop to drink. A solution in which water is the dissolving medium is called an aqueous solution . Water is the medium in which most of the chemical reactions in the body take place.

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Aqueous Reactions and Solution Stoichiometry

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  1. Aqueous Reactions and Solution Stoichiometry Water based chemistry

  2. Water, water everywhere and not a drop to drink. • A solution in which water is the dissolving medium is called an aqueous solution. • Water is the medium in which most of the chemical reactions in the body take place.

  3. General properties of aqueous solutions. • A solution is a homogenous mixture of two or more substances. • The solvent is the substance present in the greatest quantity. • The solutes are anything else in the solution.

  4. Electrolytic Properties • If two solutions are prepared one containing sugar the other containing salt, how do they differ. • Both are clear, colorless and odorless • They taste different • The one containing salt will conduct electricity. • Ions in the salt solution will carry electric charge from one electrode to the other. • The conductivity of a NaCl solution indicates the presence of ions.

  5. Ionic Compounds in Water • NaCl consists of an orderly arrangement of Na+ and Cl- ions. • When NaCl dissolves in water each ion separates from the solid structure and disperses throughout the solution. This process is called dissociation.

  6. Ionic Compounds in Water (cont) • Water is an excellent solvent for ionic compounds. • This is due to the dipole moment of the water molecule.

  7. Molecular Compounds in Water • When a molecular compound dissolves in water, the solution consists of intact molecules dispersed throughout the solution.

  8. Strong and weak electrolytes • Strong electrolytes are those that exist in solution completely or nearly completely as ions. • Weak electrolytes are those that exist in solution mostly as whole molecules. Such as acetic acid (CH3COOH) where only about 1% of the ions dissociate at any given moment.

  9. Chemical Equilibrium • The balance between two opposing processes produces a state called equilibrium. • At chemical equilibrium the relative numbers of each type of ion or molecule in the reaction are constant over time. • p 123

  10. Precipitation reactions • Reactions that result in the formation of an insoluble product are called precipitation reactions. • A precipitate is an insoluble solid formed by a reaction in a solution.

  11. Solubility Guidelines for Ionic Compounds • The solubility of a substance at a given temperature is the amount of the substance that can be dissolved in a given quantity of solvent at that temperature. • Table 4.1 on page 125

  12. Exchange Reactions • Mg(NO3)2 (aq) + 2 NaOH(aq)  Mg(OH)2 (s)↓ + 2 NaNO3 (aq) • The reaction appears to exchange Cations or Anions are called exchange reactions or metathesis reactions. • Practice exercise page 127

  13. Ionic Equations • It may be useful to indicate explicitly whether the dissolved substances are present as ions or molecules • Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq) • Molecular Equation • Pb2+ + 2NO3- + 2K+ +2I-  PbI2(s) + 2K+ + 2NO3- • Ionic Equation

  14. Acid-Base Reactions • Acids are substances that ionize in aqueous solutions to form hydrogen ions, thus increasing the concentration of H+ in water. • Bases are substances that ionize in aqueous solutions to form hydroxide ions, thus increasing the concentration of OH- in water or react to accept a H+ ion. (Ammoia) In both instances the concentration of OH- in water increases.

  15. Strong and Weak Acids and Bases • Acids and bases that are strong electrolytes are called strong acids and strong bases.

  16. Neutralizing Reactions and Salts. • When a solution of an acid and a solution of a base are mixed, a neutralization reaction occurs. • HCl(aq) + NaOH(aq)  H2O(aq) + NaCl(aq) • acid base water salt • A neutralization reaction produces water and a salt.

  17. Acid-Base Reactions with Gas Formation • 2HCl(aq) + Na2S(aq)  H2S(g) + 2NaCl(aq) • Carbonates and Bicarbonates react with acids to form CO2 gas. • HCl(aq) + NaHCO3(aq)  NaCl + H2O + CO2

  18. Oxidation-Reduction Reactions • These are reactions that involve the transfer of electrons between the reactants. • When a metal corrodes or rusts it loses electrons and forms cations. • Ca(s) + 2H+(aq)  Ca2+ + H2(g) • Loss of electrons by a substance is called oxidation. Thus Ca loses 2 electrons and is oxidized forming Ca2+

  19. Oxidation-Reduction Reactions • When an atom, ion, or molecule has become more negatively charged, that is, it gained electrons, it is said to be reduced. • The gain of electrons by a substance is called reduction. • When one reactant loses electrons (oxidized) another reactant must gain them.

  20. Oxidation Number • Oxidation numbers are a bookkeeping system to keep track of electrons gained or lost by a reactant. • Each atom in a neutral molecule or ion is assigned an oxidation number. • It is important to remember that this is really a bookkeeping system because electrons are not really kept solely by one atom or another.

  21. Rules for assigning Oxidation Numbers • For an atom in its elemental form, the oxidation number is always zero. • For any monatomic ion the oxidation number equals the charge on the ion. • K+ has an oxidation number of +1 • S2- has an oxidation number of -2 • Nonmetals usually have negative oxidation number • The sum of the oxidation number of all atoms in a neutral compound must be zero.

  22. Oxidation of Metals by Acids and Salts • The reaction of a metal with either an acid or a metal salt conforms to the following pattern. • A + BX  AX + B • As an example: • Zn(s) + 2HBr  ZnBr2 + H2 • Mn(s) + Pb(NO3)2(aq)  Mn(NO3)2 (aq)+ Pb(s) • These reactions are called displacement reactions.

  23. Metals and Acids • Many metals undergo displacement reactions with acids, producing salt and hydrogen gas. • P 139 example • Whenever one substance is oxidized, some other substance must be reduced.

  24. The Activity Series • A list of metals sorted by decreasing ease of oxidation is called an activity series. • P 141

  25. Concentration of Solutions • Concentration is a term describing the amount of a solute dissolved in a given quantity of solvent. The greater the amount of solute the more concentrated the resulting solution.

  26. Molarity • Molarity expresses the concentration of a solution as the number of moles of solute in a liter of solution. • moles of solute • Molarity = volume of solution in liters • A 1.00 molar solution contains 1.00 moles of solute in every liter of solution.

  27. Molarity Example • Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate (Na2SO4) in enough water to form 125 ml of solution. • 1st find number of moles of sodium sulfate • 0.165 mol • How many liters of solution • 0.125 L • Molarity = 0.165 mol/0.125 L = 1.32 M

  28. Expressing the Concentration of an Electrolyte • When an ionic compound dissolves, the relative concentration of the ions intoduced into the solution depend upon the chemical formula of the compound. • For example a 1.0 M solution of NaCl contains 1.0 M of Na+ and 1.0 M of Cl- • 1.0 M solution of Na2SO4 contains 2.0 M of Na+ and 1.0 M of SO42-

  29. Converting between Molarity, Moles and Volume • There are three quantities in the Molarity calculation, Molarity, moles and volume. If we know any two we can then calculate the third. • Example : how many moles of HNO3 in 2.0 L of 0.200 M solution. • moles HNO3 = (2.0 L)(0.200 mol HNO3/1.0 L soln) • = 0.40 mol HNO3

  30. Dilution • Adding water to a solution lowers the concentration of the solution a process called dilution. • The number of moles of a solute prior to dilution is equal to the number of moles of solute after dilution. • Example 4.14 p 148 • How many mL of 3.0 M sulfuric acid are needed to make 450 mL of 0.10 M sulfuric acid?

  31. Solution Stoichiometry • Example 4.15 p 150 • How many grams of Ca(OH)2 are needed to neutralize 25 mL of 0.100 M HNO3

  32. Titrations • Titration is a common laboratory method of quantitative chemical analysis that is used to determine the unknown concentration of a known reactant. • The point at which stoichiometrically equivalent quantities are brought together is known as the equivalence point of the titration. • Acid-Base indicators are dyes used to determine the equivalence point in neutralization reactions.

  33. Titration • Titration is the combining of a known concentration of a solution with an unknown to determine the concentration of the unknown. • Example 4.17 p 152

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