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Solutions

Solutions. Solute – what is dissolved Solvent – what the solute is dissolved into. Aqueous solutions have water as the solvent. Water solubility depends on:. Polarity of solute and solvent.

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Solutions

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  1. Solutions Solute – what is dissolved Solvent – what the solute is dissolved into. Aqueous solutions have water as the solvent.

  2. Water solubility depends on: • Polarity of solute and solvent. • Relative attractions of the ions for each other as opposed to the ions attraction for the water molecules

  3. Strong electrolytes • A solution of a strong electrolyte is able to conduct a current efficiently • The ionic substance completely ionizes. Soluble salts, strong acids, strong bases

  4. Weak Electrolytes • Solutions that do not conduct electricity efficiently. • Only a small amount of the ionic compound ionizes Insoluble or slightly soluble salts, weak acids or bases

  5. Nonelectrolytes • Solutions that do not conduct electricity. • Solute is polar but not ionic so no ions can be produced. • Molecular compounds in solution.

  6. Molarity (M) • M = moles of solute liters of solution

  7. A solution of ethanol is prepared by dissolving 75.0 mL of ethanol (density = 0.79g/mL) in enough water to make 250.0 mL of solution. What is the molarity of this solution?

  8. Calculate the concentration of all ions present in a solution made from 1.00g of K2SO4 in 250.0 mL of solution.

  9. How would you prepare 500 mL of a 0.500M solution of sodium carbonate from the pure solid?

  10. Dilutions • Adding water to a stock or concentrated solution in order to prepare a lower molarity solution. • Moles of solute before dilution =moles of solute after dilution

  11. M=molarity; V=volume • M1xV1 = moles of solute before dilution • M2xV2 = moles of solute after dilution • M1xV1 = M2xV2

  12. How would you prepare 250 mL of 6.0M sulfuric acid from concentrated (18M) sulfuric acid?

  13. Precipitation Reactions • When two solutions are mixed an insoluble product is formed. • To predict the product of these reactions solubility rules must be known. • Table 4.1 p. 152

  14. Write the balanced molecular, the complete ionic, and the net ionic equations for the following • FeSO4(aq) + KCl(aq) • Al(NO3)3(aq) + Ba(OH)2(aq) • CaCl2(aq) + Na2SO4(aq) • K2S(aq) + Ni(NO3)2(aq

  15. Solution Stoichiometry • Identify reactants and products • Balance net ionic equation • Calculate moles of reactants • Determine limiting reactant • Calculate moles of product • Convert to grams or other unit

  16. What volume of 0.100M Na3PO4 is required to precipitate all the lead(II) ions from 150.0 mL of 0.250M Pb(NO3)2 ?

  17. What mass of barium sulfate can be produced when 100.0 mL of a 0.100M solution of barium chloride is mixed with 100.0 mL of a 0.100M solution of iron(III) sulfate?

  18. Acid –Base Reactions • Acid- proton donor • Base – proton acceptor • Acid + base salt + water • Neutralization reaction.

  19. Write the balanced molecular, the complete ionic, and the net ionic equations for the following • HNO3(aq) + Al(OH)3 • HC2H3O2(aq) + KOH(aq) • Ca(OH)2(aq) + HCl(aq)

  20. Hydrochloric acid (75.0 mL of 0.250M) is added to 225.0 mLof 0.0550M Ba(OH)2 solution. What is the concentration of the excess H+ or OH- ions left in solution?

  21. Oxidation-Reduction Reactions • Reactions where one or more electrons are transferred. • The oxidation state or number of an element will change during the reaction. • Rules for assigning oxidation numbers – Table 4.2 p.167

  22. Species Oxidized - increase in oxidation state , lose electron(s) and act as the reducing agent. • Species Reduced - decrease in oxidation state, gain electron(s) and act as the oxidizing agent.

  23. Assign oxidation states to all atoms • UO22+ • As2O3 • NaBiO3 • Cl2 • Mg2P2O7

  24. Specify which of the following are redox reactions, identify the oxidizing agent, the reducing agent, the species being oxidized and the species being reduced • Cu(s)+ 2Ag+(aq) 2Ag(s) + Cu2+(aq) • HCl(g) + NH3(g)  NH4Cl(s) • SiCl4(l)+ 2Mg(s) 2MgCl2(s) + Si(s)

  25. Balancing Redox Equations • Write separate half reactions • Balance each half reaction • Elements except H and O • Oxygen using water • Hydrogen using H+ • Charge using electrons

  26. Balancing Redox Equations • Equalize electrons transferred by multiplying half reactions by whole numbers. • Add half reactions together • If the reaction is in a basic environment neutralize the H+ ions with OH- ions.

  27. Balance the following in acid solution • Cr2O7(aq)+ Cl-(aq) Cr3+(aq) + Cl2(g) • Pb(s)+PbO2(s)+H2SO4(aq)PbSO4(s) • Mn2+(aq)+NaBiO3(s)Bi3+(aq)+MnO4(aq)

  28. Balance the following in basic solution • Cr(s)+CrO42-(aq) Cr(OH)3(s) • MNO4-(aq)+ S2-(aq) MnS(s) + S(s) • CN-(aq)+MnO4-(aq)CN-(aq)+MnO2(s)

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