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Chapter 4 Reactions in Aqueous Solution. Outline. Solute Concentrations: Molarity Precipitation Reactions Acid-Base Reactions Oxidation-Reduction Reactions. Review. In Chapter 3, we learned about chemical reactions Most reactions were between pure gases, liquids and solids

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Chapter 4 Reactions in Aqueous Solution


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    1. Chapter 4Reactions in Aqueous Solution

    2. Outline • Solute Concentrations: Molarity • Precipitation Reactions • Acid-Base Reactions • Oxidation-Reduction Reactions

    3. Review • In Chapter 3, we learned about chemical reactions • Most reactions were between pure gases, liquids and solids • No solvent was used

    4. Background: 1. Many reactions occur in aqueous solutions: • Three common types of reactions in solution: Precipitation, Acid-base, and Oxidation-reduction • Concentration of solutions is measured in units of Molarity • Water is the universal solvent

    5. 2. Definitions • Solution: homogeneous mixture of a solvent and a solute (not always in same phase – solid/liquid/gas) • Aqueous: dissolved in water • Anion: a negatively charged ion • Ex O2-, CN-1 • Cation: positively charged ion • Ex H+ • Electrodes: measure “electron” flow in a solution

    6. Solute Concentrations - Molarity • Definition of molarity • Molarity = moles of solute/liters of solution • Symbol is M • Square brackets are used to indicate concentration in M • [Na+] = 1.0 M Example: 1.5 moles of NaCl are dissolved to make 250mL aqueous solution.

    7. Additivity • Masses are additive; volumes are not • The total mass of a solution is the sum of the mass of the solute and the solvent • The total volume of a solution is not the sum of the volumes of the solute and solvent • Molarity as a conversion: Use: # moles = 1 Liter

    8. Volumetric Glassware • Volumetric pipets, burets and flasks are made so that they contain an exact volume of liquid at a given temperature • Preparing solutions with concentrations in M involves using volumetric glassware

    9. Figure 4.1 – Preparation of Molar Solution

    10. Example 4.1

    11. Dissolving Ionic Solids • When an ionic solid is dissolved in a solvent, the ions separate from each other • MgCl2 (s) → Mg2+ (aq) + 2 Cl-1 (aq) • The concentrations of ions are related to each other by the formula of the compound: • Molarity of MgCl2 = Molarity of Mg2+ • Molarity of Cl-1 = 2 X Molarity of MgCl2 • Total number of moles of ions per mole of MgCl2 is 3

    12. Example 4.2

    13. Solubility: • Soluble compounds that dissolve • Insoluble compounds that do not dissolve

    14. Precipitation • Precipitation in chemical reactions is the formation of a solid where no solid existed before reaction • Precipitation is the reverse of solubility, where a solid dissolves in a solvent to produce a solution

    15. Precipitates • Precipitates are called insoluble – they do not dissolve in solution • Precipitation of an insoluble solid • Mix a solution of nickel(II) chloride with one of sodium hydroxide • A solid forms: Ni(OH)2(s)

    16. Figure 4.4

    17. Figure 4.3 – Precipitation Diagram

    18. Solubility Trends • Mostly soluble • Compounds of Group 1 and NH4+ cations • All nitrates • All chlorides, except for AgCl • All sulfates, except for BaSO4

    19. Solubilities Trends • Mostly insoluble • Carbonates and phosphates, except for the Group I and ammonium • Hydroxides, except for the Group 1, Group 2 and ammonium

    20. Simple Solubility Rules: SAP (compounds containing sodium, ammonium, and potassium are soluble) CAN (chlorate, acetate, and nitrate containing compounds are soluble)

    21. Example 4.3

    22. Net Ionic Equations • Consider the precipitation of CaCO3 from solutions of CaCl2 and Na2CO3 Formula Equ. Ioinic Equ. Net Ionic Equ.

    23. Spectator Ions: ions that remain soluble on the products side of the reaction Net ionic equations - follow the rules for equations • Atoms must balance • Charges must balance • Show only the ions that react

    24. Example 4.4

    25. Example 4.5 - Precipitation Stoichiometry

    26. Acids and Bases • Everyday life includes contact with many acids and bases

    27. Strong and Weak Acids and Bases • Strong acids ionize completely to H+ • HCl (aq) → H+ (aq) + Cl- (aq) • In a solution of 1.0 M HCl, there is 1M H+ and 1M Cl- • No HCl is left un-ionized • Other strong acids ionize in similar fashion

    28. Weak Acids • Weak acids ionize only partially • HB (aq) ⇌H+ (aq) + B- (aq) • HF (aq) ⇌H+ (aq) + F- (aq) • Commonly, weak acids are 5% ionized or less; double headed arrow means the reaction is moving in both directions

    29. Strong Bases • Strong bases ionize completely to OH- • NaOH (s) →Na+ (aq) + OH- (aq) • Ca(OH)2→Ca2+ (aq) + 2 OH- (aq)

    30. Strong Acids and Bases

    31. Weak Bases • Weak bases ionize only partially • NH3 (aq) + H2O ⇌ NH4+ (aq) + OH- (aq) • CH3NH2 (aq) + H2O ⇌ CH3NH3+ (aq) + OH- (aq) • Commonly, weak bases are 5% ionized or less

    32. Strong Acid – Strong Base Reactions: Neutralization Reaction: Double replacement reaction, one product will always be water; best to write as H(OH) Example: H2SO4 + NaOH 

    33. Strong Acids and Bases: Must be memorized: Strong Acids: Br I Cl SO NO ClO 4,3,4 Strong Bases: hydroxides of group I except the first 1(H) and group II except the first 2(Be and Mg)

    34. Example 4.6

    35. Acid-Base Titrations • Commonly used to determine the Molarity of a solution

    36. Titrations • Titrant (in the buret) • Know concentration • Know volume • Analyte (in the Erlenmeyer flask) • Know volume or mass • Unknown concentration

    37. Titrations Indicator: Dye solution that changes color at a set pH Equivalence Point: the place in the titration where the number of moles of acid and moles of base in the flask are equal Endpoint: the place in the titration where the color changes

    38. Figure 4.7 – An Acid-Base Titration

    39. Example 4.7

    40. Acids and Metals • Many metals will react with acids, producing hydrogen gas

    41. Oxidation-Reduction Reactions • Short name: Redox reactions • Electron exchange • Oxidation is a loss of electrons; increase charge • Reduction is a gain of electrons; decrease charge

    42. Reaction of Zinc with an Acid • Zn (s) + 2 H+ (aq) → Zn2+ (aq) + H2 (g) • Consider two half equations: • Zn loses two electrons • Zn (s) →Zn2+ (aq) + 2 e- • H+ gains an electron • 2H+ (aq) + 2 e-→ H2 (g)

    43. Principles: • Oxidation and reduction must occur together • The total number of electrons on each side of the equation must be equal; no net change

    44. Cause and Effect • Something must cause the zinc to lose two electrons • This is the oxidizing agent – the H+ • Something must cause the H+ to gain two electrons • This is the reducing agent – the Zn

    45. Reducing Agents • Reducing agents become oxidized • We know that metals commonly form cations • Metals are generally reducing agents

    46. Oxidizing Agents • We know that many nonmetals form anions • To form an anion, a nonmetal must gain electrons • Many nonmetals are goodoxidizing agents

    47. Rules Governing Oxidation Numbers 1. The oxidation number of an element that is alone (including diatomic elements) is zero. 2. The oxidation number of a element in a monatomic ion is the charge on the ion 3. Certain elements have the same oxidation number in most compounds • Group 1 metals are +1 • Group 2 metals are +2 • Oxygen is always -2 • Hydrogen is always +1 4. Oxidation numbers sum to zero (compound) or to the charge (polyatomic ion)

    48. Example 4.8

    49. Redox Reactions and Oxidation Numbers • Oxidation is an increase in oxidation number • This is the same as a loss of electrons (LEO) • Reduction is a decrease in oxidation number • This is the same as a gain of electrons (GER)

    50. Example: • Which element is being oxidized and which is being reduced? Fe  Fe+2 + 2e- F + 1e-  F-1