1 / 37

Chapter 4

Chapter 4. Types of Chemical Reactions And Solution Stoichiometry. Section 4.1: Water, The Common Solvent.

nissim-hewitt
Download Presentation

Chapter 4

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 4 Types of Chemical Reactions And Solution Stoichiometry

  2. Section 4.1:Water, The Common Solvent Hydration of an ionic compound will occur when the partial positive end of a water becomes attracted to the anions in the compound; likewise for the partial negative center of the water and the cations. Solubility depends on the strength of the intermolecular attractions between the ions and water, as well as the intramolecular attractions of the cations and anions of the compound. NH4NO3(s) NH4+(aq) + NO3-(aq)

  3. Soluble Alcohols ex: C2H5OH Sugars ex: C6H12O6 Ionic compounds ex: NaCl, KOH, LiBr Insoluble Fats ex: bacon grease Oils ex: cooking oil Non-Polar Substances ex: turpentine What can dissolve in H2O? WHY? Because of intermolecular forces: the OH group on the sugars and alcohols is particularly attractive to a water molecule. Generally speaking: “Like Dissolves Like”

  4. Section 4.2:Strong and Weak Electrolytes • Solute + Solvent = Solution • Strong electrolytes conduct electricity • Weak electrolytes barely conduct electricity • Conductivity depends upon ionization

  5. Strong Electrolytes Soluble salts Strong acids Strong bases All of these dissociate completely in water. Weak Electrolytes Weak acids Weak bases All of these partially dissociate in water HCl  H+ + Cl- NaOH  Na+ + OH- HC2H3O2H+ + C2H3O2 Non-Electrolytes are completely molecular substances in water (not even a little dissociation); Non polar substances.

  6. Section 4.3:Composition of Solutions • Concentration is measured in molarity, molality, and many others. • Concentration DOES NOT directly express the number of ions present in a solution. M=moles solute liters solution MgCl2 Mg2+ + 2Cl- 1.0 M 1.0 M 2.0 M

  7. Sample Problems • Calculate the number of moles of Cl- ions in 1.75 L of 1 x 10-3 M ZnCl2. A chemist needs 1.0 L of 0.20 M K2Cr2O7 solution. How much solid K2Cr2O7 must be weighed out to make this solution?

  8. Standard Solution: a solution whose concentration is accurately known. • Example: 0.1022 M HCl; 1.003 M NaOH • Creating dilutions • Chemical analysis of a compound • Theoretical Calculations What would you do to prepare a standard solution? In your answer, include specific pieces of glassware, techniques, or equipment you should use. ANSWER NOW

  9. Dilutions • Dilution is the process used to make the solution less concentrated. moles before dilution = moles after dilution Because M =mol/L, V1(M1) = V2(M2) Lab Technique: Use a pipet to deliver the correct amount of original solution to a volumetric flask. Add some water, swirl. Fill to line, invert.

  10. You have a large quantity of 1.5 M NaOH solution available. Dilute this to 100.0mL of a 0.05 M solution. Submit your calculations and store your final product for use in our first lab. DO NOW

  11. Section 4.4:Types of Chemical Reactions There are more than just these few types, but in this chapter we will cover… • Precipitation • Acid-base • Oxidation-Reduction

  12. Section 4.5:Precipitation Reactions • Precipitation Reactions (double displacement) • Forms a solid precipitate from aqueous reactants. • Color of precipitate can help in identification • Solubility rules help BUNCHES MORE…

  13. Solubility RULES • All compounds containing alkali metal cations and the ammonium ion are soluble. • All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are soluble. • All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb2+, and Hg2+. • All sulfates are soluble except those containing Hg2+, Pb2+, Sr2+, Ca2+, and Ba2+. • All hydroxides are only slightly soluble, except those containing an alkali metal, Ca2+, Ba2+,and Sr2+. NaOH and KOH are the most soluble hydroxides. • All compounds containing PO43-, S2-, CO32-, and SO32- are only slightly soluble except for those containing alkali metals or the ammonium ion.

  14. Practice Predicting • Potassium nitrate and barium chloride • Sodium sulfate and lead (II) nitrate • Potassium hydroxide and iron (III) nitrate

  15. ALL REACTIONS SHOULD BE WRITTEN IN NET IONIC FORM

  16. Section 4.7:Stoichiometry of Precipitation Reactions • Stoichiometry in a precipitation reaction is performed just like stoichiometry for a molecular reaction. • You need to know which ion comes from which molecular formula.

  17. Sample problem • Calculate the mass of solid NaCl needed to add to 1.5 L of 0.1 M silver nitrate solution to precipitate all Ag+ ions in the form of AgCl. Net Ionic Eq: Ag+ + Cl- AgCl

  18. General Format • Write the Net Ionic Equation • Calculate the moles present • Identify the Limiting Reactant* • Use Mole Ratio(s) • Fancy-fy your answer (put in correct units)

  19. Try Me! • What mass of precipitate will be produced when 50.0 mL of 0.200M aluminum nitrate is added to 200.0 mL of 0.100 M potassium hydroxide?

  20. Section 4.8: Acid-Base Reactions Acids yield H+ Bases yield OH - • Definitions of acid and base vary. • Arrhenius and Bronsted/Lowry are common theories. • Acid-Base rxns are called NEUTRALIZATIONS Bases are proton acceptors Acids are proton donors

  21. Strong Acid-Strong Base (HCl) (NaOH) Both dissociate completely H+ + OH- H2O Na+ and Cl- are spectators. Weak Acid - Strong Base (HC2H3O2) (KOH) Acetic acid will not dissociate KOH will completely HC2H3O2 + OH- H2O + C2H3O2- K+ is a spectator.

  22. Stoichiometry sample • What volume of 0.100 M HCl is needed to neutralize 25 mL of 0.35 M NaOH? H+ + OH- H2O

  23. Titrations…define me! • Volumetric analysis • Titration • Titrant • Analyte • Equivalence point • Indicator • Endpoint

  24. To complete a successful titration… • The reaction between the titrant and the analyte should be known (you should know WHAT substances you have) • The equivalence point should be marked accurately (you should use the right indicator) • Volume of the titrant needed to reach the equivalence point should be recorded accurately (you should use a buret!)

  25. Effective Indicator Ranges

  26. Titration Try Me Calc 1 A 50.0 mL sample of a sodium hydroxide solution is to be standardized. 1.3009 M of KHP (potassium hydrogen phthalate, KHC8H4O4) is used as the titrant. KHP has one acidic hydrogen. 41.20 mL of the KHP solution is used to titrate the sodium hydroxide solution to the endpoint. What is the resulting concentration of the analyte?

  27. Titration Try Me Calc 2 • How many milliliters of a 0.610 M sodium hydroxide solution are needed to neutralize 20.0 mL of a 0.245 M sulfuric acid solution?

  28. Norton Tutorial • Go to the website http://www.wwnorton.com/college/chemistry/chemistry3/ch/17/chemtours.aspx • Find the tutorial on Acid/Base ionization. • Complete the tutorial question form.

  29. Section 4.9: Redox Reactions What is it?? -A reaction that occurs in conjunction with a transfer of electrons. We assign oxidation states to individual atoms in a reaction to observe the change in electrons. Oxidation states are written with the +/- sign before the quantity. Ion charges are written with the +/- sign behind the quantity.

  30. Assigning Oxidation States

  31. Oxidation= an increase in the oxidation state • Reduction = a decrease in the oxidation state oxidation 2Na(s) + Cl2(g) 2NaCl(s) reduction

  32. Metal Atom • Oxidized Substance: • Loss of electrons • Oxidation state increases • Gets Smaller • Called the Reducing Agent Other Atom e- The metal is oxidized and the other substance is reduced. Other Ion Metal Ion • Reduced Substance: • Gain of electrons • Oxidation state decreases • Gets Bigger • Called the Oxidizing Agent

  33. Section 4.10: Balancing Redox How To, in Acid: • Write the ½ reactions • Balance the non-H and non-O atoms • Balance O by adding H2O where needed • Balance H by adding H+ where needed • Balance charge using e- • Multiply by coefficients until both e- are equal for each ½ reaction • Add the ½ reactions together (cancel stuff)

  34. Redox Sample Problem Balance: MnO4- + Fe2+  Fe3+ + Mn2+ 2+ +2 3+ +3 Check Charges! +7 +2 8+ 1- 2+ 0 x5! x5!

  35. Redox Try Me Problem As2O3(s) + NO3- H3AsO4 + NO(g)

  36. How To, in Base: • Repeat steps from old procedure • Cancel out H+ by adding OH- ions • Re-write H+ and OH- as water • Add ½ reactions together (and cancel stuff)

  37. Redox Try Me Problem 2 Balance, in base: Ag(s) + CN- + O2 Ag(CN)2-

More Related