1 / 107

Chapter 9 (9.1-9.4)

Chapter 9 (9.1-9.4). Chemical Reactions in Aqueous Solutions. 1. Observing and Predicting Reactions. How do we know whether a reaction occurs? What observations indicate a reaction has occurred? In your groups, make a list of changes that indicate a chemical reaction has occurred. 2.

sorena
Download Presentation

Chapter 9 (9.1-9.4)

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. Chapter 9 (9.1-9.4) Chemical Reactions in Aqueous Solutions 1

  2. Observing and Predicting Reactions • How do we know whether a reaction occurs? What observations indicate a reaction has occurred? • In your groups, make a list of changes that indicate a chemicalreaction has occurred. 2

  3. Precipitate Formation Cr3+ + 3OH-  Cr(OH)3 Ag+ + Cl- AgCl Pb2+ + CrO42-  PbCrO4 3 return

  4. Cesium in a bathtub Temperature Change Brainiac Thermite movie Ba(OH)2.8H2O + NH4Cl Thermite: Al + Fe2O3 4 return

  5. Smoke 2Al + 3Br2 2AlBr3 P4 + 5O2P4O10 2Na + Cl22NaCl 5 return

  6. Solid Decomposition • (NH4)2Cr2O7(s) Cr2O3(s) + 4H2O(g) + N2(g) • CuSO4.5H2O CuSO4 + 5H2O(g) • 2NI3 N2 + 3I2 Heat CuSO4 NI3 6 return

  7. Gas Bubbles Cu + 2H+ Cu2+ + H2(g) Na + 2H2O  NaOH + H2(g) Mg + 2HCl  MgCl2 + H2 7 return

  8. Crystal Formation/Solid Deposition Cu + 2Ag+ Cu2+ + 2Ag Zn + Sn2+ Zn2+ + Sn 8 return

  9. Explosion Cesium in a bathtub Exploding Whale Building Demolition 9 return

  10. Properties of Aqueous Solutions • Substances behave differently when they are placed in water, specifically ionic versus co- valent compounds. • One breaks apart in water, the other does not. • Which one is more likely to be pulled apart by water molecules? • Electrolytes are ionic and strong acid solutions (e.g., GatoradeTM); Nonelectrolytes are covalent compounds (e.g., sugar); weak electrolytes are in between.

  11. Electrolytes in Aqueous Solutions Strong/Weak Electrolytes

  12. Electrolytic Properties • Strong electrolyte: substance that, when dissolved in water, results in a solution that can conduct electricity (NaCl) • Weak electrolyte: substance that is a poor conductor of electricity when dissolved in water (CH3COOH – vinegar) • Nonelectrolyte: substance that doesn’t conduct electricity when dissolved in water (CH3OH – methanol)

  13. Electrolytes in Aqueous Solutions Strong/Weak Electrolytes

  14. Electrolytes • 3 substances: A2X, A2Y, and A2Z. Which one is the strongest electrolyte? The weakest? (Water has been omitted for clarity.)

  15. Properties of Aqueous Solutions • Most reactions in general chemistry take place in an aqueous environment. What does that mean? • Terms: • Solution: homogeneous mixture of two or more substances • Solute: substance present in smaller amount • Solvent: substance present in greater amount • Aqueous solution: solvent is water

  16. Ways Reactions Occur • Three general categories: • Precipitation: insoluble (solid) product is formed from aqueous solutions • Acid-base neutralization: acid and base react to form water and a salt (ionic compound) • Oxidation-Reduction: electrons are transferred between atoms in reaction • Combination • Decomposition • Single-replacement (metal or hydrogen)

  17. Products of Precipitation Reactions • Precipitation reactions always begin with two ionic compounds. • Example: NaCl (aq) + AgNO3 (aq)  ? • Draw these compounds in two separate aqueous environments. What are the possible products when they are combined?

  18. NO3- NO3- NO3- Cl- Cl- Cl- Predict Products and Types of Reactions +  Na+ Na+ Ag+ Ag+ Write formulas of products (based on charges), predict phases (Solubility Rules), and balance the equation. Na+ Ag+

  19. Solubility Rules 19 If not covered by the rules, it is probably insoluble.

  20. Solubility Rules • Determine if the following ionic compounds will be soluble (aq) or insoluble (s) in water: • K2CO3 • BaSO4 • PbI2 • NaClO4 • Ag2S • (NH4)3PO4 • Cu(OH)2

  21. Solubility Rules Answers • Determine if the following ionic compounds will be soluble (aq) or insoluble (s) in water: • K2CO3 soluble (aq) • BaSO4 insoluble (s) • PbI2 insoluble (s) • NaClO4 soluble (aq) • Ag2S insoluble (s) • (NH4)3PO4 soluble (aq) • Cu(OH)2 insoluble (s)

  22. Molecular, Ionic, and Net Ionic Equations • There are 3 ways to represent ppt reactions: • As whole compounds (molecular equation) • As ionic species (ionic equation) – more accurate • As participants in reaction (net ionic equation) • Any aqueous ionic substance is written as a compound (e.g., AgNO3), but this isn’t accurate. What does this look like in water? • Solids, liquids, and gases remain as compounds. 22

  23. Formation of Silver Chloride • Molecular equation: • NaCl(aq) + AgNO3(aq) AgCl(s) + NaNO3(aq) • Ionic equation (write separate ions for soluble (aq) compounds): • Na+(aq) + Cl-(aq) + Ag+(aq) +NO3-(aq) AgCl(s) + Na+(aq) + NO3-(aq) • Net ionic equation (cancel any identical ion on both sides of the equation, called spectator ions): Ag+(aq) + Cl-(aq) AgCl(s) • Note: s, l, and g stay together!!!!! Ag+/Cl- ions

  24. Chemistry humor, ha ha!

  25. Precipitation Reactions • Reaction of lead (II) nitrate and potassium iodide. What is the precipitate? • Write the molecular, ionic, and net ionic equations.

  26. Acid-Base Reactions • Acid: substance that breaks apart in water to form H+(e.g., HCl, HNO3, CH3COOH, lemon, lime, vitamin C). • HA(aq) H+(aq) + A-(aq) • Base: substance that breaks apart in water to form OH- (e.g., NH3, DranoTM, Milk of MagnesiaTM) • MOH(aq) M+(aq) + OH-(aq) acid base strong acid weak acid 26

  27. Common Acids and Bases • You need to KNOW these!!! • Strong acids: HCl, HBr, HI, HClO4, HClO3, H2SO4, HNO3 • Strong bases: LiOH, KOH, NaOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2 27

  28. Na+ H+ Na+ Cl- Cl- OH- OH- H+ Acid-Base Neutralization • What does “neutralization” mean? • How might these two solutions react when combined? +

  29. Acid-Base Neutralization • Neutralization reaction: reaction between acid and base; products are salt (ionic compound) and water • HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) • Acid + base  salt + water • What are the ionic and net ionic equations for these reactions? • HF(aq) + NaOH(aq) • KOH (aq) + H2SO4 (aq) 

  30. Acid-Base Reaction HCl Mg(OH)2 Acid + Base → H2O + Salt Water is created from H+ and OH- 2HCl(aq) + Mg(OH)2(s)→ 2H2O(l) + MgCl2(aq) The salt is created from spectator ions

  31. Complete/Balance These Equations • Complete and balance these equations. Write ionic and net ionic equations, if applicable. • Na2S(aq) + CuCl2(aq) • MgSO4(aq) + BaCl2(aq)  • KNO3(aq) + CaCl2(aq)  • CuSO4(aq) + NaOH(aq)  • HF(aq) + NaOH(aq) • KOH(aq) + H2SO4(aq) 31

  32. Solubility Rules 32 If not covered by the rules, it is probably insoluble.

  33. Complete/Balance These Equations • Na2S(aq) + CuCl2(aq) 2 NaCl(aq) + CuS(s) • Cu2+(aq) + S2-(aq)  CuS(s) • MgSO4(aq) + BaCl2(aq)  MgCl2(aq) + BaSO4(s) • Ba2+(aq) + SO42-(aq)  BaSO4(s) • 2 KNO3(aq) + CaCl2(aq)  2 KCl(aq) + Ca(NO3)2(aq) • No reaction • CuSO4(aq) + 2 NaOH(aq) Cu(OH)2(s) + Na2SO4(aq) • Cu2+(aq) + 2 OH-(aq) Cu(OH)2(s) • HF(aq) + NaOH(aq)H2O (l) + NaF (aq) • H+ (aq) + OH- (aq)  H2O (l) • 2KOH(aq) + H2SO4(aq) 2H2O (l) + K2SO4 (aq) • H+ (aq) + OH- (aq)  H2O (l) 33

  34. Group Quiz #19 • Determine the products of the reaction. Identify the phase of each compound, and balance the equation. Also write the ionic and net ionic equations. • Molecular: Na2S + Cr(NO3)3 • Complete Ionic: • Net Ionic:

  35. Oxidation-Reduction Reactions • Oxidation-Reduction (redox) reactions: electron-transfer reactions • When iron rusts, it loses electrons to form a cation, oxygen gain electrons to form an anion: 4 Fe(s) + 3 O2(g) 2 Fe2O3(s) • Use oxidation number rules to determine gain and loss of electrons. • Oxidation numbers are assigned as if elements in compounds completely transferred electrons (like in ionic compounds).

  36. Assigning Oxidation Numbers • 1) An atom (or molecule) in its elemental state has an oxidation number of 0. • 2) An atom in a monatomic ion (Na+, Cl-) has an oxidation number identical to its charge. • 3a) Hydrogen has an oxidation number of +1, unless it is combined with a metal, in which case it has an oxidation number of –1. • 3b) Oxygen usually has an oxidation number of -2. Oxygen in peroxides (O22-) has an oxidation number of -1. 36

  37. Rules continued • 3c) Halogens usually have an oxidation number of -1 (except when bonded to oxygen or in polyatomic ions). • 4) The sum of oxidation numbers is 0 for a neutral compound and is equal to the net charge for a polyatomic ion. (Example: NaCl = 0, SO42- = -2) • 4a) For binary ionic compounds, the position of the element in the periodic table may be useful: • Group IA: +1; Group IIA: +2; Group VIIA: –1; Group VIA: –2; Group VA: –3 37

  38. Example • H2SO4 • H = +1; O = –2 • S is unknown, so leave this for last. • The overall charge on this compound is 0. • Use algebra to solve for S: • 2(+1) + 1(x) + 4(-2) = 0 • Solve for each element: MgCr2O7 38

  39. Assigning Oxidation Numbers • Determine values of the oxidation number of each element in these compounds or ions: H2O SO2 CCl4 H2O2 Fe3(PO4)2 MnO4- NaNO3 KClO4 39

  40. Assigning Oxidation Numbers • Determine values of the oxidation number of each element in these compounds or ions: H2O H: +1, O: -2 SO2S: +4, O: -2 CCl4C: +4, Cl: -1 H2O2H: +1, O: -1 Fe3(PO4)2Fe: +2, P: +5, O: -2 MnO4-Mn: +7, O: -2 NaNO3Na:+1, N:+5, O: -2 KClO4K: +1, Cl: +7, O: -2 40

  41. Oxidation-Reduction Reactions • Oxidized: atom, molecule, or ion becomes more positively charged • Loss of electrons is oxidation (LEO) • Reduced: atom, molecule, or ion becomes less positively charged (reduced charge) • Gain of electrons is reduction (GER) • Or: OIL RIG (oxidation is loss; reduction is gain) Redox 1 Redox 2

  42. Assigning Oxidation Numbers

  43. Oxidation-Reduction Reactions • The substance oxidized causes the other substance to be reduced and is called the reducing agent. • The substance reduced causes the other substance to be oxidized and is called the oxidizing agent. • 4 Fe(s) + 3O2(g) 2Fe2O3(s)

  44. Oxidation-Reduction Reactions • Identify the element or ion oxidized/reduced. Also identify the oxidizing agent and the reducing agent. • Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) Reduced Oxidized

  45. Group Quiz #20 • Identify the oxidation number of each element in the compounds or ions below: • Ba(ClO3)2 • SO32- • For the reaction below, identify what has been oxidized and reduced; identify the oxidizing agent and the reducing agent. • Cu(s) + 2AgNO3(aq)  2Ag(s) + Cu(NO3)2(aq)

  46. Redox Reactions • Combination (1 product) • Na(s) + Cl2(g) • Decomposition (1 reactant) – usually give off gases • CuCO3(s) • Single Replacement (or Displacement) (start and end with an element and a compound) • Zn(s) + HCl(aq) 46

  47. Metals versus Nonmetals • Metals tend to react with water to form bases: • 2Na (s) + 2H2O  2NaOH + H2 • MgO (s) + H2O  Mg(OH)2 • Nonmetals tend to react with water to form acids: • 2F2 (g) + 2H2O  4HF + O2 • CO2 (g) + H2O  H2CO3 Alkali metals + water

  48. CO2 (s) + H2O (l)  H2CO3 (aq)

  49. Combination Reactions • element + element  compound • H2(g) + O2(g) • metal + nonmetal  ionic compound • Na(s) + Cl2(g) • nonmetal + nonmetal  covalent compound • C(s) + O2(g) • Why are these redox reactions? 49

  50. Combination Reactions • Mg + O2 K + Cl2

More Related