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Mushing along in Chapter 4

Mushing along in Chapter 4. Currently moving through Chapter 4. Sort of breaks down like this… Three classes of reactions (that take place in water) Acid Base Precipitation Oxidation Reduction (Redox). Applications Redox chem in everyday life

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Mushing along in Chapter 4

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  1. Mushing along in Chapter 4 • Currently moving through Chapter 4. Sort of breaks down like this… • Three classes of reactions (that take place in water) • Acid Base • Precipitation • Oxidation Reduction (Redox). • Applications • Redox chem in everyday life • Titrations (with which you have some experience) • Oodles of calculations involved with these reactions

  2. Recall from Chapter 3 • Important calculation chapter—sets tone • Molar (Molecular, atomic) mass • % composition (and empirical formula) • Molarity (moles/liter)—and diltions (mv = m2v2) • Stoichiometric coefficients (balnc’d chem eq) • G—mol—mol—g calcs (really mol—mol!) • Limiting reagents • Theoretical yields • Actual and % yields

  3. Aqueous (water) chemistry • Chapter 4 gives an intro to water chem—after all, ‘tis most abundant molecule on earth’s surface… • Much of the “interesting” aq. chemistry involves ionic solutions (water and +/- ions) • Cations--+ charged • Anions …you get the idea.

  4. What’s so interesting? • This separation of charge (+/-) resulted in a ‘flow’ of electric charge—electricity • Simple schematic of electric flow in a solution…note that anode is where anions flow (to + charged plate)

  5. What factors influence charge flow • Simple—concentration of ions!! • OK, so which compounds produce ions? • Three types of compounds…electrolytes • Strong electrolytes—ALL particles are ions • Weak electrolytes—small % of ionic particles • Non-electrolytes—zero % ionic particles • Strong electrolytes conduct, non-electrolytes do NOT, weak…you can guess…

  6. Compounds & electrolyte classes • Soluble ionic compounds (NaCl) are SE’s • A few molecular compounds (essentially the acids we’ve already discussed) SE’s • Most molecular compounds are weak to non electrolytic • Organic compounds also weak/non (carboxylic acids/amines exceptions…weak

  7. CH3CO2H(aq) H1+(aq) + CH3CO21-(aq) Electrolytes in Aqueous Solution Strong Electrolytes: Compounds that dissociate to a large extent into ions when dissolved in water. KCl(aq) K1+(aq) + Cl1-(aq) Weak Electrolytes: Compounds that dissociate to a small extent into ions when dissolved in water. The double headed arrow denotes rxns that ‘go both ways’ and the ‘larger’ arrow shows the favored reaction direction

  8. Electrolytes in Aqueous Solution Strong Acids: hydrochloric acid, hydrobromic acid, hydroiodic acid, perchloric acid, nitric acid, sulfuric acid.

  9. Electrolytes in Aqueous Solution Ionic Compounds

  10. Electrolytes in Aqueous Solution Weak acids

  11. Electrolytes in Aqueous Solution Molecular Compounds (other than any strong or weak electrolytes)

  12. Electrolytes illustrated

  13. [Ion] Calculations in solution • What is the concentration of NaCl in a solution where 1 mol of NaCl is dissolved in 1 L of water? • Albuquerque? No, but it is a trick question… • Concentration is zero. All the salt dissolves, Na+/Cl- • Concentrations of Na+ and Cl-, though, are 1.0 mol/L or 1.0 M • What about salts with two ions per formula unit? K2CO3

  14. More than one ion per formula unit • In 0.010 M K2SO4: • two moles of K+ ions are formed for each mole of Na2SO4 in solution, so [K+] = 0.020 M. • one mole of SO42– ion is formed for each mole of Na2SO4 in solution, so [SO42–] = 0.010 M. • What if there’s more than one source of K? • Mixture of KCl and K2SO4 • An ion can have only one concentration in a solution, even if the ion has two or more sources. • Ion concentration is the sum of [ ] from each source

  15. Example 4.1 Calculate the molarity of each ion in an aqueous solution that is 0.00384 M Na2SO4 and 0.00202 M NaCl. In addition, calculate the total ion concentration of the solution. After a while, you will become familiar enough that these [ ]’s are calculated by inspection—for example Highest [ ] of H+ in the following— 0.10 M HCl 0.10 M CH3COOH (acetic acid) 0.10 M H2SO4 0.15 M NH3 (ammonia)

  16. How do you know if something dissolves • Even NaCl isn’t infinitely soluble in water (36 g/100 mL) • If a compound dissolves to give a solution > 0.010 M it’s viewed as soluble • When two salts are mixed in solution, it’s possible to prepare a precipitate (a compound INSOLUBLE in water) • Just like the table of polyatomic ions from the previous chapter, there is a table of solubility guidelines that help us predict the solubility of a given compound

  17. With these guidelines we can predict precipitation reactions. • Is there a GOOD way to cut through this without straight memrztion? • YES!! Think about the ions as CHARGED species • Ions with a charge of + or – 1 are GENERALLY soluble • Ions with + - 2 charge tend to be INSOLUBLE • WHY? ATTRACTION.

  18. All -1 anions SOLUBLE NO3 ClO4…perchlorate CH3COO- acetate, anion of acetic acid Halides (Cl, Br, ) a couple of exceptions All +1 cations soluble Na, K, Rb, Cs NH4+ + and – 2 SO42-, sulfates soluble Ex. Ba, Sr (+2), Pb +2, Hg2(+2) CO3, S, (both -2), PO4 (-3) are insoluble (unless +1) Makes sense, high charge OH- insoluble (strange -1) Group 2 cations (Mg, Ca) are slight/moderately soluble Another way to look at it

  19. Reactions that Form Precipitates • There are limits to the amount of a solute that will dissolve in a given amount of water. • If the maximum concentration of solute is less than about 0.01 M, we refer to the solute as insoluble in water. • When a chemical reaction forms such a solute, the insoluble solute comes out of solution and is called a precipitate.

  20. Writing equations of pptn reactions • Several ways to write these reactions • Complete equation (one with all ions present) • Net-ionic equation (eliminates spectator ions) • Spectator ions—ions that appear in same form on both sides of a chemical equation • Write the complete and net-ionic equation for the AB reaction of HCl/NaOH. • HCl + NaOH  NaCl + H2O (complete) • H+ + Cl- + Na+ + OH-  Na+ + Cl- + H2O (Ionic form) • H+ + OH-  H2O (NET IONIC Equation)

  21. Example 4.4 Predict whether a precipitation reaction will occur in each of the following cases. If so, write a net ionic equation for the reaction. • Na2SO4(aq) + MgCl2(aq) ? • (NH4)2S(aq) + Cu(NO3)2(aq) ? • K2CO3(aq) + ZnCl2(aq) ?

  22. Example –sponsored by Campbell’s Soup! One cup (about 240 g) of a certain clear chicken broth yields 4.302 g AgCl when excess AgNO3(aq) is added to it. Assuming that all the Cl– is derived from NaCl, what is the mass of NaCl in the sample of broth? write the net ionic equation (predict precipitate?) simple calculations we’ve already covered

  23. Reactions Involving Oxidation/Reduction • Second type of reaction in Chapter 4 • Can be complex, but the “goals” of redox for this semester • ID (and balance) simple Redox rxns • Recognize if any redox eq. Is properly balanced • Oxidation: Loss of electrons • Reduction: Gain of electrons • Both oxidation and reduction must occur simultaneously. • A species that loses electrons must lose them to something else (something that gains them). • A species that gains electrons must gain them from something else (something that loses them). • AFTER ALL—you can’t just vanish an electron (cons. of mass) • Historical: “oxidation” used to mean “combines with oxygen”; the modern definition is much more general.

  24. Oxidation Numbers • An oxidation number is the charge on an ion, or a hypothetical charge assigned to an atom in a molecule or polyatomic ion. • You will hear me call this the oxidation ‘state’…it’s the same thing (like molar mass/formula mass) • Examples: in NaCl, the oxidation number of Na is +1, that of Cl is –1 (the actual charge). • In CO2 (a molecular compound, no ions) the oxidation number of oxygen is –2, because oxygen as an ion would be expected to have a 2– charge. • The carbon in CO2 has an oxidation number of +4 (Why?)

  25. A “7-stepper” Rules for Assigning Oxidation Numbers • For the atoms in a neutral species—an isolated atom, a molecule, or a formula unit—the sum of all the oxidation numbers is 0. • For the atoms in an ion, the sum of the oxidation numbers is equal to the charge on the ion. • In compounds, the group 1A metals all have an oxidation number of +1and the group 2A metals all have an oxidation number of +2. • In compounds, the oxidation number of fluorine is –1. • In compounds, hydrogen has an oxidation number of +1. ex. NaH • In most compounds, oxygen has an oxidation number of –2. peroxides • In binary compounds with metals, group 7A elements have an oxidation number of –1, group 6A elements have an oxidation number of –2, and group 5A elements have an oxidation number of –3.

  26. Oxidation Numbers of Nonmetals • The maximum oxidation number of a nonmetal is equal to the group number. • For nitrogen, +5. • For sulfur, +6. • For chlorine, +7. • The minimum oxidation number is equal to the (group number – 8). • Note goofyness of some elements…

  27. Example 4.7 What are the oxidation numbers assigned to the atoms of each element in (a) KClO4 (b) Cr2O72– (c) CaH2(d) Na2O2(e) Fe3O4

  28. Identifying Oxidation–Reduction Reactions • In a redox reaction, the oxidation number of a species changes during the reaction. • Oxidation occurs when the oxidation number increases (species loses electrons). • Reduction occurs when the oxidation number decreases (species gains electrons). • If any species is oxidized or reduced in a reaction, that reaction is a redox reaction. • Examples of redox reactions: displacement of an element by another element; combustion; incorporation of an element into a compound, etc.

  29. Oxidizing and Reducing Agents • An oxidizing agent causes another substance to be oxidized. • The oxidizing agent is reduced. • A reducing agent causes another substance to be reduced. • The reducing agent is oxidized. Mg + Cu2+ Mg2+ + Cu What is the oxidizing agent? What is the reducing agent?

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