1 / 27

Back from Break? Hope yours was well

Back from Break? Hope yours was well. Exams returned and covered—overall, I was impressed with the class. Chapter 10 problems (a few, but they go pretty quickly) posted below. Intermolecular Forces (van der Waal forces, polar…etc). 32-43, 101, 116. Types of reaction—only three.

tudor
Download Presentation

Back from Break? Hope yours was well

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Back from Break? Hope yours was well • Exams returned and covered—overall, I was impressed with the class. • Chapter 10 problems (a few, but they go pretty quickly) posted below. • Intermolecular Forces (van der Waal forces, polar…etc). • 32-43, 101, 116

  2. Types of reaction—only three.

  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-electrolyts 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. Electrolytes illustrated

  8. [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

  9. 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

  10. 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)

  11. 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 (like Cu and OH in the lab you did this week) it’s possible to prepare a precipitate (a compound that is INSOLUBLE in water) • compound

  12. 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.

  13. 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

  14. 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)

  15. 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) ?

  16. Reactions of Acids and Bases:Strong and Weak Acids • Strong acids are strong electrolytes; completely ionized in water. In water: HCl(g) → H+(aq) + Cl–(aq) No HCl in solution, only H+ and Cl– ions. • Weak acidsare weak electrolytes. Some of the dissolved molecules ionize; the rest remain as molecules. • In water: CH3COOH(l)  H+(aq) + CH3COO–(aq) Just a little H+ forms. Some acids have more than one ionizable hydrogen atom. They ionize in “steps” (more in Chapter 15). H2SO4→ H+ + HSO4– HSO4– H+ + SO42–

  17. Just a little OH– forms. Most of the weak base remains in the molecular form. Strong and Weak Bases • Strong bases: Most are ionic hydroxides (Group IA and IIA, though some IIA hydroxides aren’t very soluble). • Weak bases: Like weak acids, they ionize partially. Ionization process is different. • Weak bases formOH– by accepting H+ from water … NH3 + H2O → NH4++ OH– CH3NH2 + H2O → CH3NH3++ OH– methylamine methylammonium ion

  18. Acid–Base Reactions:Net Ionic Equations • In the reaction above, the HCl, NaOH, and NaCl all are strong electrolytes and dissociate completely. • The actual reaction occurs between ions. HCl + NaOH H2O + NaCl Na+ and Cl– are spectator ions. H+ + Cl– + Na+ + OH– H2O + Na++ Cl– H++ OH– H2O A net ionic equation shows the species actually involved in the reaction.

  19. Example 4.2 Barium nitrate, used to produce a green color in fireworks, can be made by the reaction of nitric acid with barium hydroxide. Write (a) a complete-formula equation, (b) an ionic equation, and (c) a net ionic equation for this neutralization reaction.

  20. Big Picture of “Redox” Chemistry • Redox is EVERYWHERE!! • Respiration—breathe in, breathe out… • Fermentation—sugar to wine/beer • Combustion (the original definition of oxidation/reduction) • Bleaching (blonde in a bottle) • Metallurgy—oxidation of a metal • Corrosion…ick from the roads/sea • So…what IS Redox? Let’s find out.

  21. 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.

  22. 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?)

  23. 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.

  24. 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…

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

  26. 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.

  27. 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?

More Related