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Unit 3

Unit 3. Stoichiometry and Reactions . Chemical Equations. All chemical reactions can be represented using chemical equations. When hydrogen (H 2 ) burns it reacts with oxygen (O 2 ) in the air and produces water. H 2 + O 2  H 2 O All chemical equations must be balanced.

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Unit 3

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  1. Unit 3 Stoichiometry and Reactions

  2. Chemical Equations • All chemical reactions can be represented using chemical equations. • When hydrogen (H2) burns it reacts with oxygen (O2) in the air and produces water. H2 + O2 H2O • All chemical equations must be balanced. • The Law of Conservation of Mass requires that all chemical equations be balanced.

  3. Indicating the States of Reactants and Products • Not only is it important to balance all chemical reactions, but we must also indicate what physical state all of the reactants and products are in. • Na(s) + H2O(l) NaOH(aq) + H2(g) • This will become very important soon.

  4. Stoichiometry • Stoichiometry is defined as the relationship between the relative quantities of substances taking part in a reaction. • Stoichiometry is all about ratios. • A balanced chemical equation tells us in what ratio reactants react with one another and in what ratios products are formed. • 2 H2(g) + O2(g)  2H2O(l) • This equation tells us that two molecules of H2 react with one molecule of O2 to produce two molecules of water.

  5. Using Stoichiometry • We can use stoichiometry to perform a variety of calculations when given the amount of one component in a reaction. • Example: • 2 H2(g) + O2(g)  2 H2O(l) • How many moles of water can be produced from 1.57 moles of O2?

  6. Example: • How many grams of water are produced in the combustion of 1.00 g of glucose (C6H12O6)? • Solid lithium hydroxide is used in space vehicles to remove the carbon dioxide exhaled by astronauts. The lithium hydroxide reacts with gaseous carbon dioxide to produce solid lithium carbonate and liquid water. How many grams of carbon dioxide can be absorbed by 1.00 g of lithium hydroxide?

  7. Limiting Reactants • Suppose you want to make several sandwiches using one slice of cheese and two slices of bread for each sandwich. • Using bread = Bd and Cheese = Ch to make a Bd2Ch sandwich. • The recipe for this sandwich can be represented like a chemical equation. • 2 Bd + Ch  Bd2Ch • If you have 10 slices of bread and 7 slices of cheese you will only be able to make five sandwiches before you run out of bread, leaving two slices of cheese left over. • Bread is the limiting reactant.

  8. In the reaction between hydrogen and oxygen… • 2 H2(g) + O2(g)  2 H2O(l) • Suppose we have 10 moles of hydrogen and 7 moles of oxygen. • Because we need two moles of hydrogen for every one mole of oxygen we would need 5 moles of oxygen to react with all 10 moles of hydrogen. • This would mean that after the reaction we will be left with 2 moles of oxygen.

  9. Examples • The most important commercial process for converting N2 from the air into nitrogen-containing compounds is based on the reaction of N2 and H2 to from ammonia (NH3): N2(g) + 3 H2(g)  2NH3(g) • How many moles of ammonia can be produced from 3.0 moles of nitrogen and 6.0 moles of hydrogen?

  10. Consider the following reaction that occurs in a fuel cell: 2 H2(g) + O2(g)  2 H2O(l) • The reaction produces energy in the form of electricity and water. Suppose a fuel cell is set up with 150 g of hydrogen gas and 1500 g of oxygen gas. How many grams of water can be formed?

  11. Theoretical Yield • The quantity of product that is calculated to form when all of the limiting reactant reacts is called the theoretical yield. • The amount of product actually obtained in a reactions is called the actual yield. • Percent yield is calculated by dividing the actual yield by the theoretical yield and multiplying by 100.

  12. Example • Adipic acid, H2C6H8O4 is used to produce nylon. The acid is made commercially by a controlled reaction between cyclohexane (C6H12) and O2: 2 C6H12(l) + 5 O2 2H2C6H8O4(l) + 2 H2O(g) • Assume that you carry out this reaction starting with 25.0 g of cyclohexane and that cyclohexane is the limiting reactant. What is the theoretical yield of adipic acid? • If you obtain 33.5 g of adipic acid from your reaction, what is the percent yield?

  13. Solutions and Solubility

  14. The Solution Process • A solution is created when one substance disperses uniformly throughout another. • The ability of substances to form a solutions depends on two general factors… • The types of IMF’s involved in the solution process • The natural tendency for substances to diffuse. • All of the intermolecular forces we studied in the last unit affect two substances ability to form a solution.

  15. Examples • In solutions created by dissolving ionic compounds in water Ion-dipole interactions are the forces that allow this to happen. • NaCl in H2O • In a situation where we are dissolving a nonpolar substance in a non polar solvent which IMF’s are responsible for this? • In general “Like Dissolves Like”

  16. Energy Changes and Solution Formation • When two substances form a solution a variety of interactions must be over come. • These interactions are… • Solute-solute: The intermolecular forces holding the solute particles together. • Solvent-Solvent: The intermolecular forces holding the solvent particles together. • These two interactions must be over come by solvent-solute interactions.

  17. Saturated Solutions and Solubility • As a solute begins to dissolve in a solvent the concentration of the solute particles in the solution increases. • But frequently an already dissolved solute particle can collide with a still crystallized piece of solute and recrystallize. • When the rate of how often the particles dissolve from the main crystal and how often they recrystallize are equal we have a saturated solution.

  18. The solubility of a compound is the maximum amount of that compound that will dissolve in a given amount of solvent at a specified temperature.

  19. Ways Of Expressing Concentration • Mass percent: • Mass % = • ppm = • Mole fraction =

  20. Molarity = • Molality =

  21. Aqueous Reactions and Solution Stoichiometry

  22. General Properties of Aqueous Solutions • Solute • Solvent • Electrolytic Properties • A substance whose aqueous solution contains ions is called an electrolyte (NaCl) • A substance whose aqueous solution does not contain ions is called a nonelectrolyte. (C12H22O11) • Many ionic compounds are electrolytes while many covalent molecules are not.

  23. Ionic Compounds in Water • When highly soluble ionic compounds are dissolved in water they dissociate into ions. • The solvation process helps stabilize ions in solution and prevents ions from recombining.

  24. Strong and Weak Electrolytes • Strong electrolytes are those that completely ionize in a solution. • Weak electrolytes are those that do not completely ionize in solution. • HCl Vs. CH3OOH

  25. Relative Numbers of Ions in Solution • This diagram represents an aqueous solution of one of the following compounds: MgCl2, KCl, or K2SO4. Which compound does the diagram best represent?

  26. 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 solution. Pb(NO3)2(aq) + 2 KI(aq)  PbI2(s) + 2 KNO3(aq) • This reaction forms solid Lead (II) Iodide (the yellow solid). • The other product is KNO3 and remains in solution. • PbI2 becomes a solid because the lead and the iodine attract each other very strongly making the compound insoluble.

  27. Solubility Guidelines • 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 the given temperature. • PbI2 only has a solubility of 1.2 x 10-3 mol/L at 25oC. • For now any compound that has a solubility of less than 0.01 mol/L will be considered insoluble.

  28. Solubility Rules • 1. All salts of Group IA, and ammonium are soluble.  • 2. All salts of nitrates, chlorates and acetates are soluble. • 3. All salts of halides are soluble except those of silver(I),  copper(I), lead(II), and mercury(I).  • 4. All salts of sulfate are soluble except for barium sulfate, lead(II) sulfate, and strontium sulfate.  • 5. All salts of carbonate, phosphate and sulfite are insoluble, except for those of group IA and ammonium.  • 6. All oxides and hydroxides are insoluble except for those of group IA, calcium, strontium and  barium.  • 7. All salts of sulfides and insoluble except for those of Group IA and IIA elements and of ammonium.

  29. Ionic Equations • Molecular equation: • Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq) • Ionic Equation: • Pb2+(aq) +2NO3-(aq) + 2K+(aq) + 2I-(aq)  PbI2(s) + 2K+(aq) + 2NO3-(aq)

  30. Writing Net Ionic Equations • Write a balanced molecular equation for the reaction. • Rewrite the equation to show the ions that form in solution when each soluble strong electrolyte dissociates into ions. • Identify and cancel any spectator ions.

  31. Predicting Precipitates • Note the ions present in the reactants. • Consider the possible combinations of the cations and anions • Use the solubility rules to determine if any of these combinations are insoluble. • Mg(NO3)2(aq) + 2NaOH(aq) 

  32. Example • Predict the identity of the precipitate that forms when solutions of BaCl2(aq) and K2SO4(aq) are mixed. Write a balanced net ionic equation for this reaction.

  33. Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride and sodium carbonate are mixed.

  34. Acid-Base Reactions • Acids are substances that ionize in aqueous solutions to form hydrogen ions, thereby increasing the concentration of H+(aq) ions. • Bases are substances that accept H+ ions. Bases produce hydroxide ions (OH-) when they dissolve in water. • In general ionic compounds are strong acids and bases and molecular compounds are weak acids and bases.

  35. Neutralization Reactions and Salts • When a solution of an acid and a solution of a base are mixed a neutralization occurs. • The products of an acid-base neutralization have the properties of neither. • HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)

  36. Example • Write a balanced molecular and net ionic equation for the reaction between acetic acid (CH3COOH) and barium hydroxide

  37. Oxidation Reduction Reactions • In oxidation-reduction reactions electrons are transferred. • Oxidation refers to the loss of electrons. • Reduction refers to the gain of electrons. • We can assign an oxidation number to any atom in a molecule or compound.

  38. Oxidation Number Rules • For any atom in it’s elemental form the oxidation number is always zero. • For any monatomic ion the oxidation number equals the charge on the ion. • Nonmetals usually have negative oxidation numbers. • The oxidation number of oxygen is usually -2. The major exception is in compounds called peroxides, which contain the O2-2 ion, giving each oxygen an oxidation number of -1. • The oxidation number of H is usually +1 when it is bonded with nonmetals and -1 when bonded to metals.

  39. The oxidation number of fluorine is -1 in all compounds. The other halogens have an oxidation number of -1 in most binary compounds. • The sum of the oxidation numbers of all atoms in a neutral compound is zero. The sum of the oxidation numbers in a polyatomic ion equals the charge of the ion.

  40. Example • Determine the oxidation number of sulfur in each of the following: (a) H2S (b) S8 (c) SCl2 (d) Na2SO3 (e) SO42-

  41. Oxidation of Metals by Acids and Salts • A + BX AX + B • Zn(s) + 2HBr(aq) ZnBr2(aq) + H2(g) • A list of metals in order of decreasing ease of oxidation is called an activity series.

  42. Example • Write the balanced molecular and net ionic equations for the reaction of aluminum with hydrobromic acid.

  43. Solution Stoichiometry And Chemical Analysis • How many grams of Ca(OH)2 are needed to neutralize 25.0 mL of 0.100 M HNO3? • If 45.7 mL of 0.500 M H2SO4 is required to neutralize a 2.0 mL of NaOH solution, what is the concentration of the NaOH solution.

  44. A 70.5 mg sample of potassium phosphate is added to 15.0 mL of 0.050 M silver nitrate, resulting in the formation of a precipitate. Write the molecular equation for the reaction. What is the limiting reactant in the reaction? Calculate the theoretical yield, in grams, of the precipitate that forms?

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