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Chapter 4 and 15 ( Sections 15.1, 15.3, 15.4)

Reactions in Aqueous Solutions and Acids and Bases. Chapter 4 and 15 ( Sections 15.1, 15.3, 15.4). General Properties. What are the properties of a solution?. Homogeneous mixture 2 Components: Solute – is dissolved (smaller amount) Solvent – does dissolving (larger amount)

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Chapter 4 and 15 ( Sections 15.1, 15.3, 15.4)

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  1. Reactions in Aqueous Solutions and Acids and Bases Chapter 4 and 15(Sections 15.1, 15.3, 15.4)

  2. General Properties What are the properties of a solution? • Homogeneous mixture • 2 Components: • Solute – is dissolved (smaller amount) • Solvent – does dissolving (larger amount) • Aqueous when the solvent is water

  3. Measurement of Volume What units am I allowed to use? • mL and L are the most common, but IB will not use these • IB will commonly use: • dm3 for L • cm3 for mL • Negative superscript means you invert the unit • dm-3 = 1/dm3 • cm-3 = 1/cm3

  4. Concentration What is concentration and why is it important? • Concentration is a measure of how much solute is dissolved in the solvent • Tells us how much of a reactant is present, allows us to do stoichiometry • Use square brackets when expressing concentration • Ex. [H+]

  5. Concentration How can we represent moles in a solution? • Molarity (M) • 16 M, say 16 molar • Formula • High M = concentrated • Low M = dilute

  6. Molarity • I dissolved 29.22 g of sodium chloride in 1000 mL of water. • How many moles of NaCl? • (29.22 g)/(58.44 g/mol) = 0.5000 mol • What is the volume? • 1000 mL = 1 L • What is its molarity? • (0.5000 mol)/(1 L) = 0.5 M

  7. Molarity Practice • What is the molarity of a barium chloride solution that has 40.0 g of solute dissolved in 5.0 L of water? • Solute  BaCl2 (208.23 g/mol) • Solvent  H2O • Answer  0.038 M

  8. Serial Dilution How can we make a series of solutions starting with a concentrated solution? • Called serial dilution • Start with a concentrated “stock” solution • Use the molarity ratio to figure out your measurements • M1V1 = M2V2 • 1 = initial (have) • 2 = final (wanted) • NOTE: You take the initial volume, transfer it and dilute it to the final volume with solvent

  9. Serial Dilution • Let’s take a 0.5 M salt solution. • You now need 1.0 L of a 0.125 M salt solution. • What are our initial conditions? • M1 = 0.5 M • M2 = 0.125 M • V2 = 1.0 L • Rearrange the formula to solve for the initial volume. • V1 = M2V2/M1 • V1 = (0.125 M)(1.0 L)/(0.5 M) • V1 = 0.3 L

  10. Serial Dilution Let’s take the 0.125 M salt solution. You now need 1.0 L of a 0.100 M salt solution. What are the initial conditions? Rearrange the formula and find the missing component.

  11. So, Johnny was diligently working in the lab trying to generate a little nucleation. But things weren't going so well and he's just not having any luck. Then all of a sudden his lab partner fumbles in, accidently knocking Johnny's beaker of silver nitrate into some potassium chloride which spills all over Johnny. 'Heavens to Betsy!' Johnny gleefully proclaims as a beautiful white solid of silver chloride materializes. And that's why, the legend goes, if you're not part of the solution, you're part of the precipitate.

  12. Precipitation Reaction What is the difference between soluble and insoluble? Is a precipitate soluble or insoluble? Soluble – will dissolve in a particular solvent and insoluble will not Insoluble solid product that is formed from a chemical reaction Process of creating a precipitate is called nucleation

  13. Precipitation Reactions • Sample 1  The reaction of lead(II) nitrate with potassium iodide yields a yellow precipitate of lead(II) iodide. • Sample 2 Lead(II) nitrate reacts with sodium hydroxide to form a dense white precipitate of lead(II) hydroxide. • Sample 3  A diffuse white precipitate of aluminum hydroxide is formed when aluminum chloride reacts with sodium hydroxide.

  14. General Solubility Rules for Ionic Compounds in Water (pg. 113)

  15. Solubility Practice • Please determine whether the following are soluble or insoluble in water and why. • AgNO3 • NaOH • RbClO3 • AgI • CaSO4 Soluble (nitrate) Soluble (alkali metal ex) Soluble (chlorate) Insoluble (silver ion ex) Insoluble (calcium ion ex)

  16. Precipitation Reactions What is the definition of dissociation? How is an ionic equation different from a regular molecular equation? • Compounds separate into their respective ions with appropriate charges • Ex. NaCl → Na1+ + Cl1- • Ionic equation shows charged ions as if they were dissociated

  17. Ionic Equation Example • Molecular: • Pb(NO3)2(aq)+ 2NaI(aq) → PbI2(s)+ 2NaNO3(aq) • What is the precipitate? • Lead (II) Iodide because it is the only solid product • Ionic: • Pb2+ + 2NO31- + 2Na1+ + 2I1- → PbI2 + 2Na1+ + 2NO31- • All of the aqueous compounds stay in ionic form • What is a spectator ion? • An ion that has nothing to do with the overall reaction • Ex. The sodium and nitrate ions were not involved • What is the net ionic equation? • An ionic equation that removes the spectator ions • Ex. Pb2+ + 2I1- → PbI2

  18. Ionic Equation Practice • Al(NO3)3(aq) + NaOH(aq) → Al(OH)3 + NaNO3 • What product is the precipitate? • What is the ionic equation? • What is the net ionic equation?

  19. Ionic Equation Practice • Please determine the net ionic equation. • AgNO3 (aq) + KCl (aq) AgCl + KNO3 • Mg(NO3)2 (aq) + Na2CO3 (aq)  MgCO3 + NaNO3

  20. Electrolytic Properties What is the difference between an electrolyte and a non-electrolyte? Electrolytes conduct electricity when placed in water and non-electrolytes do not Strong electrolytes conduct electricity better than weak ones All acids and bases are electrolytes Table 4.1 on Page 111

  21. Acid-Base Chemistry Using the Arrhenius definition, what are acids and bases? • Acids produce H+ ions in water • Bases produce OH- ions in water

  22. Acid-Base Chemistry Using the Bronsted-Lowry definition, what are acids and bases? • Acids are any species that can donate a proton (H+ ions) in solution • Proton Donors • 1 – monoprotic acid • 2 – diprotic acid • 3 – triprotic acid • Bases are any species that accept a proton (H+ ions) in solution • Proton Acceptors

  23. Acid-Base Chemistry • Using the Arrhenius definition, what happens when you combine an acid and a base? • Examples: • HCl + KOH  • HNO3 + KOH  • HBr + Al(OH)3 • Neutralization (irreversible) • Produces salt and water • A salt is an ionic (metal/nonmetal) compound that uses ions other than hydrogen and hydroxide • KCl + H2O • KNO3 + H2O • AlBr3 + H2O

  24. Acid-Base Chemistry • What are the products of the following neutralization reactions? • HCl + NaOH → • HClO4 + NH4OH → • H2SO4 + 2KOH → • NaCl + H2O • NH4ClO4 + H2O • K2SO4 + 2H2O

  25. Acid-Base Chemistry • Using the Bronsted-Lowry definition, what happens when you combine an acid and a base? • Proton transfer (reversible) • Produces two conjugates • Conjugate base – acid after proton lost • Conjugate acid – base after proton is gained

  26. Acid-Base Chemistry • What does amphoteric mean? • What are conjugate acid/base pairs? • A substance can act as an acid or a base • Ex. Water • OH- is called hydroxide • H3O+ is called hydronium • Pairing up the: • Acid/Conjugate Base • Base/Conjugate Acid

  27. Acid-Base Chemistry • What are the conjugate acid/base pairs in the following equation? • NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(aq) • Written as: • H2O/ OH- • NH3/ NH4+

  28. Conjugate Acids and Bases • What are the conjugate bases of these acids? • HNO3 • H2O • H3O+ • H2SO4 • HBr • HCO3- • What are the conjugate acids of these bases? • OH- • H2O • HCO3- • SO42- • ClO4-

  29. Acid-Base Chemistry What is the connection between weak and strong electrolytes? • Weak acids and bases are weak electrolytes • Do not dissociate completely, reversible reaction, so not many ions in solution • Ex. CH3COOH ↔ CH3COO- + H+ • Strong acids and bases are strong electrolytes • Dissociate completely, so more ions are in the solution • Ex. HCl → Cl- + H+

  30. 6 Rules of Oxidation States Elements (whether monatomic, diatomic or polyatomic) have an oxidation state of 0 Group 1 metals have an oxidation state of 1+, Group 2 metals are 2+ and aluminum is 3+ Oxygen has an oxidation state of 2-, except when in peroxide form then it is 1- Hydrogen has an oxidation of 1+, except when bonded to Group 1 or 2 metals then it is 1- Fluorine always has an oxidation state of 1- The sum of the oxidation states in a neutral compound are equal to zero, but in a polyatomic ion it is equal to the charge

  31. Oxidation State Review • Please determine the oxidation state of each atom: • HCl • H2O2 • MgCl2 • H3PO4 • NaH • Fe • H2 • H1+ and Cl1- • H1+ and O1- • Mg2+ and Cl1- • H1+, P5+ and O2- • H1- and Na1+ • Fe0 • H20

  32. Oxidation-reduction Introduction • What are oxidation-reduction (redox) reactions? • Reactions where electrons are transferred between elements • If electrons are exchanged, oxidation states change • If one element gains some electrons, another element must lose some

  33. Redox What is reduction? What is oxidation? How can we remember? • Gain of electrons • Loss of electrons • Two ways: • LEO GER -or- • OIL RIG

  34. Confusing Redox What are redox agents?

  35. Redox Practice • Please determine the element that has been reduced and oxidized as well as the oxidizing and reducing agents: • S + O2 → SO2 • 2N2O → 2N2 + O2 • Fe + H2SO4 → FeSO4 + H2

  36. Redox What is a half reaction? • A reaction that shows whether electrons(e-) were lost or gained • Oxidation reaction – half reaction, electrons lost are in the product • Ex. Ca → Ca2+ + 2e- • Reduction reaction – half reaction, electrons gained in the reactants • Ex. Ca2+ + 2e- → Ca

  37. Redox • Please balance the following equation. • Cu(s) + Ag+1(aq) Cu+2(aq) + Ag(s)

  38. Redox • Please balance the following equation. • Li(s) + Pb+2(aq) Li+1(aq) + Pb(s)

  39. Cleaning Up What does disproportionation mean? What is the main difference between acid-base and redox reactions? • At least one element is both oxidized and reduced in the same reaction • All based on transfers: • Acid-Base – protons • Redox – electrons

  40. Types of Reactions What are the five types of reactions? • Decomposition (breakdown) • C → A + B • Synthesis (combination) • A + B → C • Combustion • A + O2 → B + Water • Single Displacement • A + BC → AC + B • Double Displacement • AB + CD → AC + BD

  41. Displacement Reactions What are some “special” types of displacement reactions? • Hydrogen Displacement • Many metals will replace a H in water or acid • Creates hydrogen gas (H2) • Metal Displacement • Metal replacing another metal • Check activity series, related to electrode potential (Table 14) • The lower the potential the better the reducing agent • A metal will displace another metal that has a higher potential • Ex. Zn will replace Cu

  42. Displacement Reactions What are some “special” types of displacement reactions? • Halogen Displacement • F2 > Cl2 > Br2 > I2 • The higher the potential the better the oxidizing agent • A nonmetal will displace another nonmetal that has a lower potential

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