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Chemical Reactions

Chemical Reactions. Chapter 4. Electrolytes and Nonelectrolytes. Ionic Theory of Solutions  certain substances produce freely moving ions when they dissolve in water, and these ions conduct an electric current in an aqueous solution

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Chemical Reactions

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  1. Chemical Reactions Chapter 4

  2. Electrolytes and Nonelectrolytes • Ionic Theory of Solutions  certain substances produce freely moving ions when they dissolve in water, and these ions conduct an electric current in an aqueous solution • Electrolyte  a substance that dissolves in water to give an electrically conducting solution • The ions are charged and mobile • Can be strong or weak • Ex: NaCl • In general  ionic solids that dissolve in water are electrolytes • Nonelectrolytes  a substance that dissolves in water to give a non-conducting or very poor conducting solution. • Sucrose, C12H22O11 

  3. Strong and Weak Electrolytes • Strong Electrolyte  an electrolyte that exists in solution almost entirely as ions • NaCl  Na+(aq) + Cl-(aq) • Weak electrolyte  an electrolyte that dissolves in water to give a relatively small percentage of ions • Are generally molecular substances • NH3 (aq) + H20 (l) NH4+ (aq) + OH- (aq) 

  4. Solubility Rules • Soluble  dissolves in water • Insoluble  does not • Table 4.1 on page 128 or hand out 

  5. Molecular and Ionic Equations • Molecular Equations  where the reactants and products are written as if they were molecular substances, even though they may actually exist in solution as ions. • Ca(OH)2 (aq) + Na2CO3 (aq)  CaCO3 (s) + 2NaOH (aq) • The molecular equation closely describes what you actually do in the laboratory or in a industrial process • Molecular equations don’t tell you what’s happening at the ion level 

  6. Complete Ionic Equations • Complete ionic equations  when strong electolytes are written as separate ions in the solution • Ca(OH)2 (aq) + Na2CO3 (aq)  CaCO3 (s) + 2NaOH (aq) • Using the above equation each reactant is soluble in water and can be written using freely moving ions • The product CaCO3 is insoluble and left alone where NaOH is soluble and again written as ions • Ca2+(aq) + 2OH-(aq) + 2Na+(aq) + CO32- (aq)  CaCO3 (s) + 2Na+(aq) + 2OH- (aq) 

  7. Net Ionic Equations • Net ionic equations  an ionic equation form which spectator ions have been canceled • Spectator ions  ions in an ionic equation that does not participate in the reaction • Ca2+(aq) + 2OH-(aq) + 2Na+(aq) + CO32- (aq)  CaCO3 (s) + 2Na+(aq) + 2OH- (aq) • Net ionic: Ca2+(aq) + CO32- (aq)  CaCO3 (s) 

  8. Example • Write an net ionic equation for the molecular equation: • 2HClO4(aq) + Ca(OH)2(aq)  Ca(ClO4)2(aq) + H2O(l) • HC2H3O2(aq) + NaOH(aq)  NaC2H3O2(aq) + H2O(l) • 

  9. Homework • Pg 165: • 2, 3, 5, 7, 11, 12, 22, 26, 27, 30, 31, 34, 36, 38, 39, 42, 46, 50, 51, 53, 55, 59, 60, 64, 66, 73, 77, 79, 82, 83, 84, 91, 94, 100, 104

  10. Precipitation Reactions • Precipitate  an insoluble solid compound formed during a chemical reaction in solution • To predict a precipitate you need to know if the products are insoluble or not • If all reactants and products are soluble then no reaction occurs • Exchange or (metathesis) reaction  is between compounds that, when written as a molecular equation , appears to involve the exchange of parts between the 2 reactants • MgCl2 + AgNo3  

  11. Examples • Aqueous sodium chloride and iron(II) nitrate • Aqueous aluminum sulfate and sodium hydroxide 

  12. Acid-Base Reactions • Properties • Acids – have a sour taste • Bases – have a bitter taste and soapy feel • Acid – Base indicator  is a dye used to distinguish between acidic and basic solutions by means of a color change • Acids  a substance that produces hydrogen ions, H+, when dissolved in water • HNO3 (aq)  H+ (aq) + NO3- (aq) • Bases  a substance that produces hydroxide ions, OH-, when dissolved in water • NaOH (aq) + H2O (l)  NH4+ (aq) + OH- (aq) 

  13. Bronsted – Lowry Acids and Bases • Acid Donates a proton to another species in a proton-transfer rxn • HNO3 (aq) + H2O(l)  H3O+ (aq) + NO3- (aq) • Base  the species that accepts a proton in a proton-transfer rxn • NH3(aq) + H2O(l)  NH4+ (aq) + OH- (aq) 

  14. Strong Acids and Bases • See table 4.3 on page 139 • Strong acid  is an acid that ionizes completely in water, it is a strong electrolyte • HCl, HNO3 • Weak acid  is an acid that only partly ionizes in water, it is a weak electrolyte • HCN, HF • Strong base  is a base that is present in aqueous solution entirely as ions, one of which is OH-, is a strong electrolyte • The hydroxides of group IA and IIA except beryllium hydroxide • NaOH • Weak base  is a base that only partly ionizes in water, is a weak electrolyte • NH3 

  15. Neutralization Reactions • Neutralization reaction  is a reaction of an acid and base that result in an ionic compound and water • The ionic compound is neutral and normally a salt • HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l) • In each rxn there is a proton transfer where the hydroxide ion latches to the hydrogen ion producing water. • Because water is so soluble it provides the driving force of the rxn 

  16. Homework • Pg 165: • 2, 3, 5, 7, 11, 12, 22, 26, 27, 30,31, 34, 36, 38, 39, 42, 46, 50, 51, 53, 55, 59, 60, 64, 66, 73, 77, 79, 82, 83, 84, 91, 94, 100, 104

  17. Example • Write the molecular and net ionic equation for the neutralization of nitrous acid, HNO2, by sodium hydroxide, NaOH. 

  18. Acids and Bases • Acids that only have one acidic hydrogen atom per molecule such as HCl are called monoprotic acids • Polyprotic acids  are acids that have two or more acidic hydrogen atoms per molecule. • H3PO4(aq) + NaOH (aq)  NaH2PO4 (aq) + H2O (l) • Salts such as NaH2PO4 that have acidic hydrogen atoms and can undergo neutralization with bases are called acid salts 

  19. Acid-base rxns with gas formation • Some salts, (mostly carbonates, sulfites and sulfides), react with acids to form a gas product • Na2CO3(aq) + 2HCl (aq)  2NaCl (aq) + H2O (l) + CO2(g) • See table 4.4 • 

  20. Common Redox Reactions • Common redox rxns: • Combination (synthesis) • 2 substances combine to form a single substance • Ex: 2Na + Cl2 2NaCl • Decomposition • One substance is broken into 2 or more substances • Ex: 2KClO3 2KCl + 3O2 • Combustion • When a substance reacts with oxygen. With a hydrocarbon the products are CO2 and water • Ex: 2C4H10 + 13O2 8CO2 + 10H2O • Ex: 4Fe + 3O2 2Fe2O3 

  21. Common Redox Reactions • Displacement Reactions: • Single replacement  an element reacts with a compound, displacing an element from it • Ex: Cu + 2AgNO3  Cu(NO3)2 + 2Ag • Double replacement  2 compounds react, where the cations displace each other. • Ex: Pb(NO3)2+ 2NaBr  PbBr2 + 2NaNO3 • Similar to precipitation reactions

  22. Homework • Pg 165: • 2, 3, 5, 7, 11, 12, 22, 26, 27, 30,31, 34, 36, 38, 39, 42, 46, 50, 51, 53, 55, 59, 60, 64, 66, 73, 77, 79, 82, 83, 84, 91, 94, 100, 104

  23. Oxidation – Reduction Reactions • Oxidation – Reduction (redox) Reactions occur when electrons have been transferred • Oxidation numbers  the actual charge of the atom if it exists as a monatomic ion or a hypothetical charge assigned to the atom in the substance by simple rules • Assigning oxidation numbers: table 4.5, pg 147 • Oxo # of an atom in an element is 0 • Oxo # of an atom in monatomic ions is the charge of the ion • Oxo # of oxygen is -2 except H2O2 where it is -1 • Hydrogen is +1 (except with a binary compound of a metal it is -1: CaH2) • Halogens have a -1 (except in binary compounds with another halogen above it) • The oxidation number of compounds add up to 0. • The oxidation number of polyatomic ions add up to the charge of the ion 

  24. Examples • Find the oxidation number of Mn ion KMnO4 and in K2MnO4 • What is the oxidation number of Cr in Cr2O72-?

  25. Oxidation – Reduction Reactions • Half-rxn is one of two parts of a redox rxn. • Oxidation  loses electrons • Fe(s)  Fe2+(aq) + 2e- • Reduction  gains electrons • Cu2+(aq) + 2e- Cu(s) • “Leo the lion says ger”

  26. Oxidation – Reduction Reactions • Oxidizing agent  a species that oxidizes another species; it is itself reduced • Reducing agent  is a species that that reduces another species; it is itself oxidized • Fe(s) + Cu2+(aq)  Fe2+(aq) + Cu(s)

  27. Balancing Redox Reactions • Balance charge as well as atoms • assign oxidation numbers to determine what is being oxidized/reduced. • write the oxidation and reduction half-reactions. • balance the charge of each half-reaction by adding electrons. • adjust coefficients so that electrons lost/gained are equal. • add the half-reactions together (the electrons will cancel out) • Mn + FeCl3 Fe + MnCl2 • 3( Mn MnCl2 + 2e-) (oxidation) • 2( FeCl3 + 3e- Fe) (reduction) • 3Mn + 2FeCl3 3MnCl2 + 2Fe (balanced)

  28. Example • Mg(s) + N2(g) Mg3N2(s)

  29. Homework • Pg 165: • 2, 3, 5, 7, 11, 12, 22, 26, 27, 30,31, 34, 36, 38, 39, 42, 46, 50, 51, 53, 55, 59, 60, 64, 66, 73, 77, 79, 82, 83, 84, 91, 94, 100, 104

  30. Solutions • Solute  substance that dissolves • Solvent  the liquid • Molar Concentration (Molarity: M)  the moles of solute dissolved in one liter of solution • M = • A sample of NaNO3 weighing 0.38g is placed in a 50.0-mL volumetric flask. The flask is then filled with water. What is the molarity of the solution?

  31. Solutions • Diluting Solutions  when diluting solutions we use • M1V1 = M2V2 • Ex: A stock solution of NaCl is 6.00 M. How much of the stock solution is needed to prepare 1.00 L of saline solution which is 0.154 M NaCl?

  32. Quantitative Analysis • Quantitative analysis  the determination of the amount of a substance or species present in a material. • Gavimetric Analysis  were the amount of a species is determined by converting the species to a product than can be isolated completely and weighed • Precipitation reactions

  33. Example • A soluble silver compound was analyzed for the percentage of silver by adding sodium chloride solution to precipitate the silver ions as silver chloride. If 1.583 g of silver compound gave 1.788 g of silver chloride, what if the mass percentage of silver in the compound? • Can use percentage silver in silver chloride to find grams of silver in product and in the compound

  34. Example • In a study of calcium ion present in blood, the calcium ion in a solution was precipitated with sodium oxalate. If 50.0 mL of solution gave 0.358 g of calcium oxalate, CaC2O4, what was the molarity of Ca2+ ions in solution? • Convert grams of calcium oxalate to moles. 1 mol calcium oxalate  1 mol Ca2+

  35. Quantitative Analysis • Volumetric analysis  molarity is used as a conversion factor to determine the volume of solution that is equivalent to a given mass of solute. • Zinc sulfide reacts with hydrochloric acid. How many mL of 0.00512 M HCl solution are required to react with 0.0392 g ZnS?

  36. Quantitative Analysis • Titration  determines the amount of a substance by adding a measured volume of another reactant with known concentration until the reaction is complete • 5H2O2 + 2KMnO4 + 6H+ 8H2O + 5O2 + 2K+ + 2Mn2+ • What is the mass percentage of H2O2 in a solution if 57.5 g of solution required 389 mL of 0.0534 M KMnO4?

  37. Homework • Pg 165: • 2, 3, 5, 7, 11, 12, 22, 26, 27, 30, 31, 34, 36, 38, 39, 42, 46, 50, 51, 53, 55, 59, 60, 64, 66, 73, 77, 79, 82, 83, 84, 91, 94, 100, 104

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