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Aqueous Solutions Pt. 2

Aqueous Solutions Pt. 2. Reaction Types. Precipitation Reactions Cations and anions come together to form insoluble ionic compounds AgNO 3 + NaCl → AgCl + NaNO 3 Ag + + Cl - → AgCl Neutralization Reactions Acid Base reactions forming a salt and water HCl + NaOH → H 2 O + NaCl

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Aqueous Solutions Pt. 2

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  1. Aqueous Solutions Pt. 2

  2. Reaction Types • Precipitation Reactions • Cations and anions come together to form insoluble ionic compounds • AgNO3 + NaCl → AgCl + NaNO3 • Ag+ + Cl- → AgCl • Neutralization Reactions • Acid Base reactions forming a salt and water • HCl + NaOH → H2O + NaCl • H+ + OH- → H2O • Reduction Oxidation Reactions (Redox) • Electrons are transferred changing oxidation states

  3. Oxidation Reduction Reactions • Electrons are transferred from one element to another changing oxidation states • Ca(s) + 2HCl (aq) → CaCl2(aq) + H2 (g) • Ca + 2H+ → Ca2+ + H2 • If electrons are lost charge increases - oxidized • If electrons are gained charge decreases – reduced • Called oxidation because reactions with oxygen were characterized first • 2Ca(s) + O2(g) → 2CaO(s)

  4. Oxidation Numbers • Use oxidation numbers to determine which elements have lost or gained electrons • Oxidation number • Actual charge of atom if it was separated into ions • Hypothetical charge assigned by set of rules

  5. Rules • When atom is in elemental form its oxidation number is 0 • H2, P4, Ca • When an ion is monatomic its oxidation number is equal to the charge on the ion • K+ = +1, S2- = -2

  6. Rules Cont. • Nonmetals usually have negative oxidation numbers but they can be positive • The oxidation number of oxygen is usually -2 except in peroxides O22- • H2O, H2O2 • The oxidation number of fluorine is -1 in all compounds. Other halogens have a -1 ox # in binary compounds but a positive ox # when they are in oxyanions • HF, HCl, HClO3 • The oxidation number of hydrogen is +1 when attached to nonmetals and -1 when attached to metals • HI, NaH

  7. Assigning Oxidation Numbers • Write the oxidation number of sulfur in each compound. • H2S, S8, SCl2, Na2SO3, SO42- • Write the oxidation state of the bold faced element. • P2O5, NaH, Cr2O72-, SnBr4, BaO2

  8. Oxidation of Metals by Acids and Salts • Displacement reactions • Single replacement reactions • A + BX → AX + B • Zn + 2HBr → ZnBr2 + H2 • Mn + Pb(NO3)2 → Mn(NO3)2 + Pb • Metals react with acids to produce salts and hydrogen gas • Mg + 2HCl → MgCl2 + H2 • Mg + 2H+ → Mg2+ + H2

  9. Redox Reactions • Write the balanced molecular and net ionic equations for the following reactions. Identify what has been reduced and oxidized. • Aluminum and hydrobromic acid • Magnesium and cobalt (II) sulfate

  10. Activity Series • The activity series allows us to determine whether a metal will be oxidized by an acid or a particular salt. • Listed in order of decreasing ease of oxidation.

  11. Activity Series Cont. • Elements can be oxidized by ions of elements below them. • Cu + 2Ag+ → Cu2+ + 2Ag • Only metals above H2 are able to react with acids to form H2. • Ni + 2HCl → NiCl2 + H2 • Ni + 2H+ → Ni2+ + H2

  12. Activity Series • The activity series allows us to determine whether a metal will be oxidized by an acid or a particular salt. • Listed in order of decreasing ease of oxidation.

  13. Activity Series • The activity series allows us to determine whether a metal will be oxidized by an acid or a particular salt. • Listed in order of decreasing ease of oxidation.

  14. Using the Activity Series • Will an aqueous solution of iron(II) chloride oxidize magnesium metal. If so write the molecular and ionic equation. • Which of the following metals will be oxidized by lead(II) nitrate • Zn, Cu, Fe

  15. Reaction Predictions • Questions to ask when predicting how reactions will proceed. • What are the reactants? • Are they electrolytes or non-electrolytes? • Are they acids and bases? • If they are electrolytes will they produce a precipitate, water or gas? • If not, can they react in an oxidation reduction reaction?

  16. Concentrations of Solutions • Concentration – ([ ]) amount of solute dissolved in a quantity of solvent or solution • Molarity (M) – moles of solute / liters of solution • Make 250mL of a 1.00 M CuSO4 • Measure .250 mol of CuSO4 • Place in volumetric flask/ beaker • Add water until final volume if 250mL

  17. Molarity Practice • What is molarity of a solution made by dissolving 23.4g of Na2SO4 in enough water for the final volume to be 125mL? • What is the molarity when 5.0 grams of has been dissolved in 100mL of water?

  18. Concentration of Electrolytes • When ionic solid dissolves the [ ] of ions depends on the formula • 1M NaCl → 1M Na+, 1M Cl- • 1M Na2SO4 → 2M Na+, 1M SO42- • Concentration can be expressed in terms of initial compound or ions • What is the [ ] of each ion present in 0.025M calcium nitrate and 0.015M potassium carbonate?

  19. Molarity and Moles • Molarity allows us to relate moles of solute and L of solution • How many grams of Na2SO4 are required to make 350mL of 0.500M Na2SO4 solution? • How many grams of Na2SO4 are there in 15mL of 0.50M Na2SO4? • How many mL of 0.50M Na2SO4 are needed to provide 0.038mol of the salt?

  20. Dilutions • Purchased solutions come in high concentrations to lower [ ] we need to dilute them • The amount of solute does not change but volume does • Moles solute before dilution = Moles of solute after dilution • How do make 250mL of 0.0100M CuSO4 from a 1M stock solution? • M1V1=M2V2

  21. Dilution Practice • How many mL of 3.0M H2SO4 are need to make 450mL of 0.10M H2SO4? • What volume of 2.50M Pb(NO3)2 contains 0.500mol of Pb2+? • How many mL of 5.0M K2Cr2O7 must be diluted to make 250mL of a 0.10M stock solution? • If 10mL of a 10.0M stock solution of NaOH is diluted to 250mL, what is the resulting [ ]?

  22. Solution Stoichiometry

  23. Solution Stoichiometry Practice • How many grams of calcium hydroxide are needed to neutralize 25.0mL of 0.100M HNO3? • How many grams of sodium hydroxide are needed to neutralize 20.0mL of 0.150M H2SO4? • How many liters of 0.500M HCl are needed to react completely with 0.100mol of lead(II) nitrate making lead(II) chloride?

  24. Titrations • Find [ ] of a solute use a technique called titration • Combine a sample with a reagent of a known [ ] (standard solution) • Reaction is run until it is complete = equivalence point • Identified using indicators such as phenolphthalein • Dyes change color to show the reaction is completer

  25. Titrations • Use titration to find [ ] of HCl • [NaOH] = 0.100M • Take 20.00mL sample of acid • Add dye • Add NaOH until color change

  26. Titration Example • How many grams of chloride are in a sample of water if 20.2mL of a 0.100M silver ion solution is needed to react with all of the chloride in the sample? If the sample has a mass of 10.0g, what is the percentage of chloride present in the sample?

  27. Titration Practice • A sample of iron is dissolved in acid producing iron(II). The sample is titrated with 47.20mL of 0.02240M permanganate. The reaction proceeds as follows: • How many moles of permanganate were added? • How many moles of iron(II) were in the sample? • How many grams of iron were in the sample? • If the sample had a mass of 0.8890g what was the percentage of iron present in the sample?

  28. Integrated Practice • A 70.5mg sample of potassium phosphate react with 15.0mL of 0.050M silver nitrate producing a precipitate. • Write the molecular and net ionic equations. • What is the limiting reagent? • Calculate the theoretical yield of the precipitate

  29. Homework • 37, 42, 44, 46, 47, 48, 51, 55, 58, 62, 68, 71, 78, 94, 97

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