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Chapter 4 Aqueous Reactions and Solution Stoichiometry

CHEMISTRY The Central Science 9th Edition. Chapter 4 Aqueous Reactions and Solution Stoichiometry. David P. White. General Properties of Aqueous Solutions. Electrolytic Properties Aqueous solutions, solutions in water, have the potential to conduct electricity.

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Chapter 4 Aqueous Reactions and Solution Stoichiometry

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  1. CHEMISTRYThe Central Science 9th Edition Chapter 4Aqueous Reactions and Solution Stoichiometry David P. White Chapter 4

  2. General Properties of Aqueous Solutions • Electrolytic Properties • Aqueous solutions, solutions in water, have the potential to conduct electricity. • The ability of the solution to conduct depends on the number of ions in solution. • There are three types of solution: • Strong electrolytes, • Weak electrolytes, and • Nonelectrolytes. Chapter 4

  3. General Properties of Aqueous Solutions Electrolytic Properties Chapter 4

  4. General Properties of Aqueous Solutions • Ionic Compounds in Water • Ions dissociate in water. • In solution, each ion is surrounded by water molecules. • Transport of ions through solution causes flow of current. Chapter 4

  5. General Properties of Aqueous Solutions • Molecular Compounds in Water • Molecular compounds in water (e.g., CH3OH): no ions are formed. • If there are no ions in solution, there is nothing to transport electric charge. Chapter 4

  6. General Properties of Aqueous Solutions • Strong and Weak Electrolytes • Strong electrolytes: completely dissociate in solution. • For example: • Weak electrolytes: produce a small concentration of ions when they dissolve. • These ions exist in equilibrium with the unionized substance. • For example: Chapter 4

  7. Precipitation Reactions • When two solutions are mixed and a solid is formed, the solid is called a precipitate. Chapter 4

  8. Precipitation Reactions

  9. Precipitation Reactions • Exchange (Metathesis) Reactions • Metathesis reactions involve swapping ions in solution: • AX + BYAY + BX. • Metathesis reactions will lead to a change in solution if one of three things occurs: • an insoluble solid is formed (precipitate), • weak or nonelectrolytes are formed, or • an insoluble gas is formed. Chapter 4

  10. Precipitation Reactions • Ionic Equations • Ionic equation: used to highlight reaction between ions. • Molecular equation: all species listed as molecules: • HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) • Complete ionic equation: lists all ions: • H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  H2O(l) + Na+(aq) + Cl-(aq) • Net ionic equation: lists only unique ions: • H+(aq) + OH-(aq)  H2O(l) Chapter 4

  11. Acid-Base Reactions • Acids • Dissociation = pre-formed ions in solid move apart in solution. • Ionization = neutral substance forms ions in solution. • Acid= substances that ionize toform H+ in solution(e.g. HCl,HNO3,CH3CO2H, lemon, lime, vitamin C). • Acids with one acidic proton are called monoprotic (e.g., HCl). • Acids with two acidic protons are called diprotic (e.g., H2SO4). • Acids with many acidic protons are called polyprotic. Chapter 4

  12. Acid-Base Reactions • Bases • Bases = substances that react with the H+ ions formed by acids (e.g. NH3, Drano™, Milk of Magnesia™). Chapter 4

  13. Acid-Base Reactions • Strong and Weak Acids and Bases • Strong acids and bases are strong electrolytes. • They are completely ionized in solution. • Weak acids and bases are weak electrolytes. • They are partially ionized in solution. Chapter 4

  14. Acid-Base Reactions • Identifying Strong and Weak Electrolytes • Water soluble and ionic = strong electrolyte (probably). • Water soluble and not ionic, but is a strong acid (or base) = strong electrolyte. • Water soluble and not ionic, and is a weak acid or base = weak electrolyte. • Otherwise, the compound is probably a nonelectrolyte. Chapter 4

  15. Acid-Base Reactions Identifying Strong and Weak Electrolytes Chapter 4

  16. Acid-Base Reactions • Neutralization Reactions and Salts • Neutralization occurs when a solution of an acid and a base are mixed: • HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) • Notice we form a salt (NaCl) and water. • Salt = ionic compound whose cation comes from a base and anion from an acid. • Neutralization between acid and metal hydroxide produces water and a salt. Chapter 4

  17. Acid-Base Reactions • Acid-Base Reactions with Gas Formation • Sulfide and carbonate ions can react with H+ in a similar way to OH-. • 2HCl(aq) + Na2S(aq)  H2S(g) + 2NaCl(aq) • 2H+(aq) + S2-(aq)  H2S(g) • HCl(aq) + NaHCO3(aq)  NaCl(aq) + H2O(l) + CO2(g) Chapter 4

  18. Oxidation-Reduction Reactions • Oxidation and Reduction • When a metal undergoes corrosion it loses electrons to form cations: • Ca(s) +2H+(aq)  Ca2+(aq) + H2(g) • Oxidized: atom, molecule, or ion becomes more positively charged. • Oxidation is the loss of electrons. • Reduced: atom, molecule, or ion becomes less positively charged. • Reduction is the gain of electrons. Chapter 4

  19. Oxidation-Reduction Reactions Oxidation and Reduction Chapter 4

  20. Oxidation-Reduction Reactions Oxidation and Reduction

  21. Oxidation-Reduction Reactions • Oxidation Numbers • Oxidation number for an ion: the charge on the ion. • Oxidation number for an atom: the hypothetical charge that atom would have if it was an ion. • Oxidation numbers are assigned by a series of rules: • If the atom is in its elemental form, the oxidation number is zero. E.g., Cl2, H2, P4. • For a monoatomic ion, the charge on the ion is the oxidation state. Chapter 4

  22. Oxidation-Reduction Reactions • Oxidation Numbers • Nonmetal usually have negative oxidation numbers: • Oxidation number of O is usually –2. The peroxide ion, O22-, has oxygen with an oxidation number of –1. • Oxidation number of H is +1 when bonded to nonmetals and –1 when bonded to metals. • The oxidation number of F is –1. • The sum of the oxidation numbers for the atom is the charge on the molecule (zero for a neutral molecule). Chapter 4

  23. Oxidation-Reduction Reactions • Oxidation of Metals by Acids and Salts • Metals are oxidized by acids to form salts: • Mg(s) +2HCl(aq)  MgCl2(aq) + H2(g) • During the reaction, 2H+(aq) is reduced to H2(g). • Metals can also be oxidized by other salts: • Fe(s) +Ni2+(aq)  Fe2+(aq) + Ni(s) • Notice that the Fe is oxidized to Fe2+ and the Ni2+ is reduced to Ni. Chapter 4

  24. Oxidation-Reduction Reactions • Activity Series • Some metals are easily oxidized whereas others are not. • Activity series: a list of metals arranged in decreasing ease of oxidation. • The higher the metal on the activity series, the more active that metal. • Any metal can be oxidized by the ions of elements below it. Chapter 4

  25. Concentrations of Solutions • Molarity • Solution = solute dissolved in solvent. • Solute: present in smallest amount. • Water as solvent = aqueous solutions. • Change concentration by using different amounts of solute and solvent. • Molarity: Moles of solute per liter of solution. • If we know: molarity and liters of solution, we can calculate moles (and mass) of solute. Chapter 4

  26. Concentrations of Solutions • Molarity Chapter 4

  27. Concentrations of Solutions • Dilution • We recognize that the number of moles are the same in dilute and concentrated solutions. • So: • MdiluteVdilute = moles = MconcentratedVconcentrated Chapter 4

  28. Solution Stoichiometry and Chemical Analysis • There are two different types of units: • laboratory units (macroscopic units: measure in lab); • chemical units (microscopic units: relate to moles). • Always convert the laboratory units into chemical units first. • Grams are converted to moles using molar mass. • Volume or molarity are converted into moles using M = mol/L. • Use the stoichiometric coefficients to move between reactants and product. Chapter 4

  29. Solution Stoichiometry and Chemical Analysis Chapter 4

  30. Solution Stoichiometry and Chemical Analysis • Titrations

  31. Solution Stoichiometry and Chemical Analysis • Titrations • Suppose we know the molarity of a NaOH solution and we want to find the molarity of an HCl solution. • We know: • molarity of NaOH, volume of HCl. • What do we want? • Molarity of HCl. • What do we do? • Take a known volume of the HCl solution, measure the mL of NaOH required to react completely with the HCl. Chapter 4

  32. Solution Stoichiometry and Chemical Analysis • Titrations • What do we get? • Volume of NaOH. We know molarity of the NaOH, we can calculate moles of NaOH. • Next step? • We also know HCl + NaOH  NaCl + H2O. Therefore, we know moles of HCl. • Can we finish? • Knowing mol(HCl) and volume of HCl (20.0 mL above), we can calculate the molarity. Chapter 4

  33. End of Chapter 4:Aqueous Reactions and Solution Stoichiometry Chapter 4

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