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AP Chemistry Unit 1 Exam Review: Chemical Reactions & Solution Stoichiometry

Prepare for your AP Chemistry exam with a review of key concepts in chemical reactions, solution stoichiometry, electrolytes, solubility rules, net ionic equations, molarity, dilution equations, titration, types of reactions, and gases. Understand ideal gas laws, stoichiometry, and more to ace your exam.

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AP Chemistry Unit 1 Exam Review: Chemical Reactions & Solution Stoichiometry

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  1. Ch. 4 Chemical Reactions & Solution Stoichiometry See Unit 2 Student Topic List (available on the blog!)

  2. Unit 1 Exam Analysis • 20 minutes individual review. • 10 minutes peer review….for half credit (based on reflection)! • Learning is a PROCESS! • For credit you must correct your answer and provide a brief reflection of why your answer choice was incorrect and how you corrected it. (3 to 4 sentences per question)

  3. AP Ch. 4 (Reactions and SolnStoich-- Exam Objectives • I can apply the following concepts… • Electrolytes and Ions • Soln Stoichiometry (LR/XS with solutions) • SolnStoich goes along with solubility rules • Solubility Rules! • Net Ionic Equations (NIEs), spectator ions, etc and how to write NIEs • Molarity=mol solute/L soln (what you need to calculate it) • M1V1=M2V2 “dilution equation”. Used to calculate dilutions or back calculate concentrated molarities • Titration (Redox bleach lab and what titration is and how to classify a titration rxn) • Types of Reactions: • SD • DD or precipitation • Combustion • Redox • Decomposition • Acid-Base (know strong acid/base lists & how to write strong/weak in NIEs) • **DO NOT WORRY ABOUT COMPLEX IONS** • Types of Reactions and how to predict reaction products and writing NIEs w DD • REDOX reactions and how to balance in acid and base (memorize rules in order!) • Oxidation state….know all rules on sheet. Do not forget diatomic elements are zero. • LEO goes GER • Copper Reactions Lab

  4. Next… • AP Ch. 4 (Reactions and SolnStoich) and Ch. 5 Gases Exam Objectives • I can apply the following concepts… • Kinetic Molecular Theory and Gases • Ideal vs. Real Gases • What conditions are ideal for gases? • PV=nRT PV=(m/MM)RT • Know how to convert C to K • Density and gases and MW kitty cat • Graham’s Law of Effusion • Molar volume of any gas at STP= 22.4 L/mol • STP= 0°C (273 K) and 1 atm=760 mm Hg = 760 torr • Combined Gas Law P1V1 =P2V2 • T1 T2 • Boyles Law P1V1=P2V2 • Charles’ Law V1/T1 = V2/T2 • Gay Lussac’s Law P1/T1 = P2/T2 • Avogadro’s Law or n is proportional to L or volume • Molar Volume of a Gas Lab • Dalton’s Law of Partial Pressures (Ptotal = P1 + P2 + P3 + P4 …..) • Mole fraction (it is really a mole decimal and no units) • Barometric pressure and manometers • Kinetic energy and as a function of temperature • **Sig Figs and Units!**

  5. Water, the common solvent • Water is not a linear molecule. It is bent at an angle of about 105° • Electrons are not evenly distributed around the atoms in water. The molecule is polar because the charges are not distributed symmetrically. • Image: https://www.google.com/search?q=water+molecule&client=firefox-a&hs=mA9&rls=org.mozilla:en-US:official&channel=sb&source=lnms&tbm=isch&sa=X&ei=p8f6U8OmO42AygTxnYHACg&ved=0CAgQ_AUoAQ&biw=1280&bih=621#facrc=_&imgdii=_&imgrc=iPDWj07nrlW_zM%253A%3BmkqaXgwGuHe1ZM%3Bhttp%253A%252F%252Fimages.quickblogcast.com%252F5%252F2%252F9%252F2%252F1%252F319511-312925%252FWaterMole.jpg%253Fa%253D84%3Bhttp%253A%252F%252Fmccartyhome.com%252F%253Fp%253D11%3B1400%3B617

  6. Like dissolves like. The following classes of molecules, in general, are miscible: • Polar and Ionic • Polar and Polar • Nonpolar and nonpolar • Ionic salts dissolve in water. Compounds that contain only carbon and hydrogen are nonpolar. • Predict whether each pair of substance will mix. Explain. • NaNO3 and H2O • C6H14 and H2O • I2 and C6H14 • I2 and H2O

  7. Dissociation of Ionic Salts in water • Model: • NaI(s) • Ba(OH)2(s) • Practice: • CaCl2(s) • Fe(NO3)3(s)

  8. Solution Vocabulary and ELECTROLYTES Strong Electrolyte • high conductivity • total dissociation • examples: strong acids (HCl); salts (NaCl); strong bases (Ba(OH)2) Weak Electrolyte • low conductivity • partial dissociation • examples: weak organic acids (acetic acid); weak bases (NH3) Non Electrolyte • no conductivity • close to zero dissociation • examples: sugar, AgCl • Solute dissolves into the solvent (if same phase as solvent it is the one in lesser amount) • Solvent dissolves the solute (if same phase as solvent it is the one in greater amount) • Aqueous Solution water is the solvent

  9. Composition of Solutions Molarity (M) • moles of solute per liter of solution. • M = (moles of solute) / (liter of solution) *Watch Units!*

  10. Example A • Calculate the molarity of a solution prepared by dissolving 11.85 g of solid KMnO4 in enough water to make 750.mL of solution. • Solution Steps: • 1. Convert from grams to moles. • 2. Divide moles by volume (in Liters!) to get molarity. • **GUESS METHOD!**

  11. Example B • Calculate the mass of NaCl needed to prepare 175 mL of a 0.500M NaCl solution.

  12. Example C • How many mL of solution are necessary if we are to have a 2.48 M NaOH solution that contains 31.52 g of the dissolved solid?

  13. Example D • Calculate the molarity of all the ions in each of the following solutions. • **Hint: always write out the dissociation equation** • 0.25 M Ca(OCl2) • 2 M CrCl3

  14. Example E • Determine the molarity of Cl- ion in a solution prepared by dissolving 9.82 g of CuCl2 in enough water to make 600. mL of solution.

  15. Where did water come from?

  16. Dilutions • An important part of chemistry is to be able to prepare dilute solutions from more concentrated “stock” solutions. • Remember— • “Moles of solute after dilution = moles of solute before dilution” • Thus— • M1V1 = M2V2

  17. Example f • What volume of 12 M hydrochloric acid must be used to prepare 600. mL of a 0.30 M HCl solution?

  18. Example g • What volume of 9.0 M sodium hydroxide must be used to prepare 1.2 L of a 1.0 M NaOH solution?

  19. For next time… • Lab Revisions Due: • Pre-lab for Copper Chemistry Lab (in black binder) • Abstract of Procedures and Identifying Chemical Equations! • Continue…Ch. 4 Reading • Practice Problems • **Unit 2 Review Packet**

  20. Copper Lab Groups

  21. Types of Reactions Ch. 4 Part II

  22. Types of Chemical Reactions (3) • Precipitation Reactions • Acid-Base Reactions • Oxidation-Reduction Reactions

  23. NIE POGIL • Do all reactant species participate in a reaction?

  24. Precipitation Reactions • A precipitation reaction can also be called a double displacement reaction. • When ionic compounds dissolve in water, the resulting solution contains the separated ions. • If a solid forms from a combination of selected ions in solution, the solid must contain an anion part and cation part, and the net charge on the solid must be zero. • There are some simple solubility rules to help predict the products of reactions in aqueous solutions.

  25. Simple Rules for Solubility • 1. Most nitrates (NO3-) salts are soluble. • 2. Most salts containing the alkali metal ions (Li+, Na+, K+, Cs+, Rb+) and the ammonium ion (NH4+) are soluble. • 3. Most chloride, bromide, and iodide salts are soluble. Notable exceptions are salts containing the ions Ag+, Pb+, and Hg2+ • 4.Most sulfate salts are soluble. Notable exceptions are BaSO4, PbSO4, HgSO4, and CaSO4 • 5.Most hydroxide salts are only slightly soluble. The important soluble hydroxides are NaOH and KOH. The compounds Ba(OH)2, Sr(OH)2, and Ca(OH)2 are slightly soluble. • 6.Most sulfide (S2-), carbonate (CO32-), chromate (CrO42-), and phosphate (PO43-) salts are only slightly soluble.

  26. Solubility Memorization Requirements for AP Exam All sodium, potassium, ammonium, and nitrate salts are soluble in water. **Memorization of “solubility rules” is beyond the scope of this course and the AP Exam.** Rationale: Memorization of solubility rules does not deepen understanding of the big ideas.

  27. Predicting Reaction Products • To begin, focus on the ions in solution before any reaction occurs. • Complete and balance the following reactions, determining in each case, if a precipitate is formed. • KCl(aq) + Pb(NO3)2(aq) • AgNO3(aq) + MgBr2(aq) • Ca(OH)2(aq) + FeCl3(aq) • NaOH(aq) + HCl(aq)

  28. Describing Reactions in Solution • Formula Equation (overall reaction – does not identify ions) 2AgNO3(aq) + Na2S(aq) Ag2S(s) + 2NaNO3(aq) • Complete Ionic Equation (includes ions in solution) 2Ag+(aq) + 2NO3-(aq) + 2Na+(aq) + S2-(aq) Ag2S(s) + 2NO3-(aq) + 2Na+(aq) • Net Ionic Equation (identifies species that undergo chemical change) 2Ag+(aq)+ S2-(aq) Ag2S(s)

  29. Writing Equations for Reactions Example • Write the formula, complete ionic, and net ionic forms for each of the following equations. • Aqueous nickel (II) chloride reacts with aqueous sodium hydroxide to give a nickel (II) hydroxide precipitate and aqueous sodium chloride. • Solid potassium metal reacts with water to give aqueous potassium hydroxide and hydrogen gas. • Aqueous sodium hydroxide reacts with aqueous phosphoric acid to give water and aqueous sodium phosphate.

  30. Stoichiometry of Precipitation Reactions • Steps: • 1. Identify the species present in the combined solution, and determine what reaction occurs. • 2. Write the balanced net ionic equation for the reaction. • 3. Calculate the moles of reactants. • 4. Determine which reactant is limiting. • 5. Calculate the moles of product or products, as required. • 6. Convert to grams of other units, as required. **No two examples are alike, the PROCESS to solving will be similar**

  31. Solution Stoichiometry Example • What mass of Fe(OH)3 is produced when 35. mL of a 0.250 M Fe(NO3)3 solution is mixed with 55 mL of a 0.180M KOH solution? • Individual Practice

  32. Check for UnderstandingPrecipitation Stoichiometry Please complete the following on a sheet of paper to turn in. • When aqueous solutions of Na2SO4 and Pb(NO3)2 are mixed, PbSO4 precipitates. Calculate the mass of PbSO4 formed when 1.25 L of 0.0500 M Pb(NO3)2 and 2.00 L of 0.0250 M Na2SO4 are mixed.

  33. Acid-Base Reactions • The Bronsted-Lowry Definitions • Acid is a proton donor. • Base is a proton acceptor. • An acid-base reaction is often called a neutralization reaction.

  34. Acid-Base Example • How many mL of a 0.800M NaOH solution is needed to just neutralize 40.00mL of a 0.600M HCl solution? (Hint: Write down the reaction and convert to moles, and relate moles of acid to moles of base.) • Individual Example Problems 65 page 171 – 8 Minutes

  35. Acid-Base Titrations • Volumetric analysis • A technique for determining the amount of a certain substance by doing a titration. • Titration • Involves delivery (from a buret) of a measured volume of a solution of known concentration (the titrant) into a solution containing the substance being analyzed (analyte). • Equivalence point (Stoichiometric point) • Point in titration where enough titrant has been added to react exactly with the analyte • Indicator • A substance added at the beginning of the titration that changes color at (or very near) the equivalence point • Point at which indicator actually changes color is the endpoint of the titration.

  36. Acid-base titration (Standardizing) • You wish to determine the molarity of a solution of sodium hydroxide. To do this, you titrate a 25.00mL aliquot of your sample, which has had 3 drops of phenolphthalein indicator added so that it is pink, with 0.1067 M HCl. The sample turns clear after the addition of 42.95 mL of the HCl. Calculate the molarity of your NaOH solution. • Individual Example Problems 63 and 65 page 174 – 10 Minutes

  37. Oxidation Reduction Reactions • Reactions in which one or more electrons are transferred (redox reactions) • Example: H2(g) + Cl2(g) 2HCl(g) Electrons are transferred from the hydrogen to the chlorine. • Table 4.2 Rules for Assigning Oxidation States (pg 171) • Remember: The sum of the oxidation states in a neutral compound must equal zero.

  38. Oxidation States • Assign oxidation states to each of the atoms in the following compounds: Hint: Check that oxidation states sum up to the charge on the compound/ion. • CaF2 • C2H6 • H2O • ICl5 • KMnO4 • SO42-

  39. Practice- Oxidation States Fe2O3 + 2Al Al2O3 + 2Fe THE IRON GAINED ELECTRONS. IT HAS BEEN REDUCED. THE ALUMINUM LOST ELECTRONS. IT HAS BEEN OXIDIZED. OIL – Oxidation Involves Loss (of electrons) RIG- Reduction Involves Gain (of electrons)

  40. Tricky…Tricky… • Oxidizing agent • Something that is reduced (it causes something else to oxidize) • Reducing agent • Something that is oxidized (it causes something else to reduce) Fe2O3 + 2Al Al2O3+ 2Fe • Iron was reduced. • Aluminum was oxidized. • Iron (II) oxide was the oxidizing agent. • Aluminum was the reducing agent.

  41. OIL RIG Practice • For each reaction, identify the atoms that undergo reduction or oxidation (and charges). Also, list the oxidizing and reducing agents. 2H2(g) + O2(g) 2H2O(g) Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) 2AgCl(s) + H2(g) 2H+(aq) + 2Ag(s) + 2Cl-(aq)

  42. Balancing Oxidation-Reduction Equations (Half-Reaction Method) • Steps: Acidic • 1. Write separate equations for the oxidation and reduction half-reactions. • 2. For each half-reaction • Balance all the elements except hydrogen and oxygen. • Balance oxygen using H2O. • Balance hydrogen using H+. • Balance the charge using electrons (e-). • 3. If necessary, multiply one or both balanced half-reactions by an integer equalize the number of electrons transferred in the two half-reactions. • 4. Add the half-reactions, and cancel identical species. • 5. Check that the elements and charges are balanced.

  43. Balancing Oxidation-Reduction Equations (Half-Reaction Method) • Steps: Basic • The same as acidic and then… • 6. Add a number of OH- ions to both sides so that you just balance excess H+ ions. • 7. H+ and OH- will form H2O on the side with excess H+. Free OH- will appear on one side of the equation. • 8. Double check atoms and charges!

  44. Balancing Redox Example - Acid Cu(s) + HNO3(aq) Cu2+(aq) + NO(g) • Copper is being oxidized: • Nitrogen is being reduced:

  45. Balancing Redox Reactions - Base Cr2O72-(aq) + 2NO(g) + 6H+(aq) 2Cr3+(aq) + 2NO3-(aq) + 3H2O(l) **Already balanced in acid**

  46. Practice Time! • Be sure to show your work!

  47. Chemistry of Copper Lab - _______ • Pre-lab: REQUIRED FOR ENTRY INTO LAB!! • Read lab handout thoroughly. • Set-up data organization • What data are you collecting for each step? • Complete procedure component of lab • Safety: Safety goggles at ALL times! • Concentrated acids will be used!! • Can cause severe burns • No touching your face! • Wash your hands after completing lab. • Read thoroughly…carefully pour acids and remove hot beakers from hot plate with tongs! (A hot beaker looks like a cold one!) • Avoid congestion around the fume hood (Step 5 should be done in fume hood!). • Keep lab stations cleaned through multistep procedure.

  48. Step 1: Copper Lab Task • Label beaker with tape • Block • Student group initials

  49. Bleach Redox Lab • PRE-LAB DUE (____)!! • Read handout thoroughly! (in black binder) • Complete ABSTRACT of procedure (note –safety precautions) • Complete the pre-lab questions • Hint: Dilution, Concentration, Stoichiometry • Required for entry into lab. • Also for _________… • Continue practice with Precipitation, Acid-Base, and redox reactions. (Think about quiz level) • Begin reading Ch.5 Gases and notes! (packet available)

  50. Bleach Redox Lab • Pre-lab • Safety: • Bleach is a 5% solution of sodium hypochlorite. This solution is a corrosive liquid; it can cause skin burns. The solution reacts with acid and with heat to evolve chlorine gas. The solution is moderately toxic by ingestion and inhalation. Keep away from skin and clothing. Hydrochloric acid solution is toxic by ingestion or inhalation. It is corrosive to skin and eyes. • CAUTION: Adding hydrochloric acid to bleach may cause chlorine gas to be given off. Step 6 is completed in the fume hood! • Wear your goggles at ALL times! Do NOT touch your face. Wash your hands with soap before leaving the lab.

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