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William L Masterton Cecile N. Hurley Edward J. Neth University of Connecticut Chapter 4 Reactions in Aqueous.

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Presentation on theme: "William L Masterton Cecile N. Hurley Edward J. Neth University of Connecticut Chapter 4 Reactions in Aqueous."— Presentation transcript:

1 William L Masterton Cecile N. Hurley Edward J. Neth University of Connecticut Chapter 4 Reactions 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 O 2-, 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 MgCl 2 (s) → Mg 2+ (aq) + 2 Cl -1 (aq) The concentrations of ions are related to each other by the formula of the compound: Molarity of MgCl 2 = Molarity of Mg 2+ Molarity of Cl -1 = 2 X Molarity of MgCl 2 Total number of moles of ions per mole of MgCl 2 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 NH 4 + cations All nitrates All chlorides, except for AgCl All sulfates, except for BaSO 4

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 CaCO 3 from solutions of CaCl 2 and Na 2 CO 3 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 → Ca 2+ (aq) + 2 OH - (aq)

30 Strong Acids and Bases

31 Weak Bases Weak bases ionize only partially NH 3 (aq) + H 2 O ⇌ NH 4 + (aq) + OH - (aq) CH 3 NH 2 (aq) + H 2 O ⇌ CH 3 NH 3 + (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: H 2 SO 4 + 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) → Zn 2+ (aq) + H 2 (g) Consider two half equations: Zn loses two electrons Zn (s) → Zn 2+ (aq) + 2 e - H + gains an electron 2H + (aq) + 2 e - → H 2 (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 good oxidizing 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 a.Group 1 metals are +1 b.Group 2 metals are +2 c.Oxygen is always -2 d.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

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