CH 4: Chemical Reactions

# CH 4: Chemical Reactions

## CH 4: Chemical Reactions

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1. CH 4: Chemical Reactions Renee Y. Becker Valencia Community College CHM 1045

2. Solutions • Solute – solid in liquid or lowest mass quantity of substance • Solvent- liquid solute is dissolved in or highest mass quantity of substance

3. Solution Concentrations • Concentration: allows us to measure out a specific number of moles of a compound by measuring the mass or volume of a solution. • Molarity(M) = Moles of Solute Liters of Solution moles = M•L L = moles/M

4. Example: Solution Concentrations • How many moles of solute are present in 125 mL of 0.20 M NaHCO3?

5. Example: Solution Concentrations • How many grams of solute would you use to prepare 500.00 mL of 1.25 M NaOH?

6. Solution Concentrations • Dilution: the process of reducing a solution’s concentration by adding more solvent. Moles of solute(constant) = Molarity  Volume Mi • Vi = Mf • Vf Vf = (Mi • Vi) / Mf Mf = (Mi • Vi) / Vf

7. Example: Solution Concentrations • What volume of 18.0 M H2SO4 is required to prepare 250.0 mL of 0.500 M H2SO4?

8. Example: Solution Concentrations • What is the final concentration if 75.0 mL of 3.50 M glucose is diluted to a volume of 400.0 mL?

9. SolutionStoichiometry • Titration:a technique for determining the concentration of a solution • Standard solution: known concentration • If you have a known volume of standard solution and use it to titrate a known volume of an unknown concentrated solution you can calculate to find the number of moles in the unknown and therefore find it’s concentration

10. Titration • When doing a titration you add titrant (standard solution) to the analyte (unknown concentration solution) until the endpoint or the equivalence point is reached. This point is when you have equal moles of titrant and analyte, from the volume of the titrant and analyte used and the molarity of the titrant, you can find the molarity of the analyte • Endpoint- based on an indicator • Indicator- a substance that changes color in a specific pH range • Equivalence point- not based on an indicator, usually a pH meter • Use Manalyte• Vanalyte = Mtitrant • Vtitrant

11. Example: SolutionStoichiometry • A 25.0 mL sample of vinegar (dilute CH3CO2H) is titrated and found to react with 94.7 mL of 0.200 M NaOH. What is the molarity of the acetic acid solution? NaOH(aq) + CH3CO2H(aq) CH3CO2Na(aq) + H2O(l)

12. SolutionStoichiometry

13. Oxidation–Reduction Reactions • Assigning Oxidation Numbers:All atoms have an “oxidation number” regardless of whether it carries an ionic charge. 1. An atom in its elemental state has an oxidation number of zero. Elemental state as indicated by single elements with no charge. Exception: diatomics H2 N2 O2 F2 Cl2 Br2 and I2

14. Oxidation–Reduction Reactions 2.An atom in a monatomic ion has an oxidation number identical to its charge.

15. Oxidation–Reduction Reactions 3.An atom in a polyatomic ion or in a molecular compound usually has the same oxidation number it would have if it were a monatomic ion. A. Hydrogen can be either +1 or –1. B. Oxygen usually has an oxidation number of –2. In peroxides, oxygen is –1. C. Halogens usually have an oxidation number of –1. • When bonded to oxygen, chlorine, bromine, and iodine have positive oxidation numbers.

16. Oxidation–Reduction Reactions 4. The sum of the oxidation numbers must be zero for a neutral compound and must be equal to the net charge for a polyatomic ion. A.H2SO4 neutral atom, no net charge SO42- sulfate polyatomic ion [SO4]2- [Sx O42-] = -2 X + -8 = -2 X = 6 so sulfur has an oxidation # of +6

17. Oxidation–Reduction Reactions B.ClO4– , net charge of -1 [ClO4]-1 [Clx O42-] = -1 X + -8 = -1 X = 7 so the oxidation number of chloride is +7

18. Example: Oxidation–Reduction Reactions Assign oxidation numbers to each atom in the following substances: A. CdS F. VOCl3 B. AlH3 G. HNO3 C. Na2Cr2O7 H. FeSO4 D. SnCl4 I. Fe2O3 E. MnO4– J. V2O3

19. Quiz What is the oxidation number of arsenic in AsO43- ?

20. Electrolytes in Solution • Electrolytes:Dissolve in water to produce ionic solutions. • Nonelectrolytes:Do not form ions when they dissolve in water. a) NaCl sol’n conducts electricity, completes circuit (charged particles) b) C6H12O6 does not

21. Electrolytes in Solution • Dissociation: The process by which a compound splits up to form ions in the solution.

22. Electrolytes in Solution • Strong Electrolyte:Total dissociation when dissolved in water. • Weak Electrolyte:Partial dissociation when dissolved in water.

23. Types of Reactions • Precipitation • Acid-base neutralization • Oxidation-reduction (redox) • Metathesis Reactions (double replacement)

24. Types of Chemical Reactions • Precipitation Reactions:A process in which an insoluble solid precipitate drops out of the solution. • Most precipitation reactions occur when the anions and cations of two ionic compounds change partners. (double replacement) Pb(NO3)2(aq) + 2 KI(aq) 2 KNO3(aq) + PbI2(s)

25. Solubility Rules & Precipitation • Allow you to predict whether a reactant or a product is a precipitate. • Soluble compounds are those which dissolve to more than 0.01 M. • There are three basic classes of salts:

26. Solubility Rules & Precipitation 1. Salts which are always soluble: • All alkali metal salts: Cs+, Rb+, K+, Na+, Li+ • All ammonium ion (NH4+) salts • All salts of the NO3–, ClO3–, ClO4–, C2H3O2–,and HCO3–ions

27. Solubility Rules & Precipitation 2. Salts which are soluble with exceptions: • Cl–, Br–, I– ion salts except with Ag+, Pb2+, & Hg22+ • SO42– ion salts except with Ag+, Pb2+, Hg22+, Ca2+, Sr2+, & Ba2+

28. Solubility Rules & Precipitation 3. Salts which are insoluble with exceptions: • O2– & OH– ion salts except with the alkali metal ions, and Ca2+, Sr2+, & Ba2+ ions • CO32–, PO43–, S2–, CrO42–, & SO32– ion salts except with the alkali metal ions and the ammonium ion • If not listed the compound is probably insoluble

29. Example: Solubility Rules & Precipitation • Predict the solubility of: (a) CdCO3 (b) MgO (c) Na2S (d) PbSO4 (e) (NH4)3PO4 • Write the balanced reaction and predict whether a precipitate will form for: (a) NiCl2 (aq) + (NH4)2S (aq) (b) Na2CrO4 (aq) + Pb(NO3)2 (aq) (c) AgClO4 (aq) + CaBr2 (aq)

30. Solubility Rules & Precipitation

31. Equations • Molecular equation – as the reaction is • Complete ionic equation – all broken up into ions(only aqueous solutions) • Net ionic equation – cancel out spectator ions

32. Net Ionic Equations for Precipitation Reactions • Write net ionic equation for the following reaction: 2 AgNO3(aq) + Na2CrO4(aq)  Ag2CrO4(s) + 2 NaNO3(aq) 1. Is it balanced? If not do it! 2. Separate all except solids into ions (complete ionic equation) 3. Cancel out spectator ions on both sides 4. Rewrite (net ionic equation)

33. Types of Chemical Reactions • Acid–Base Neutralization:A process in which an acid reacts with a base to yield water plus an ionic compound called a salt. • The driving force of this reaction is the formation of the stable water molecule. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

34. Acid–Base Concepts • Arrhenius Acid: A substance which dissociates in water to form hydrogen ions (H+). • Arrhenius Base: A substance that dissociates in, or reacts with, water to form hydroxide ions (OH–). Limitations: Has to be an aqueous solution and doesn’t account for the basicity of substances like NH3.

35. Acid–Base Concepts • Brønsted Acid:Can donate protons (H+) to another substance. • Brønsted Base:Can accept protons (H+) from another substance. (NH3)

36. Example : Conjugate acid-base pairs For the following reactions label the acid, base, conjugate acid, and conjugate base. CH3CO2H(aq) + H2O(l) H3O+(aq) + CH3CO2-(aq) NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)

37. Quiz Which of the following is a Bronsted-Lowry base but not an Arrhenius base? • NaOH • NH3 • Mg(OH)2 • KOH

38. Acids and Bases • Strong acid - st. electrolyte, almost completely dissociates in water • HCl, H2SO4, HNO3, HClO4, HI, HBr • Weak acid - wk. electrolyte, does not dissociate well in water • HF, HCN, CH3CO2H • Strong base - st. electrolyte, almost completely dissociates in water • Metal hydroxides • Weak base - does not dissociate well in water • NH3

39. Types of Chemical Reactions • Metathesis Reactions: These are reactions where two reactants just exchange parts. (double replacement) AX + BYAY + BX BaCl2(aq) + K2SO4(aq) BaSO4(s) + 2 KCl(aq) This is also a ppt reaction, if I ask you what type of reaction is it, what is the best answer??

40. Types of Chemical Reactions • Oxidation–Reduction (Redox) Reaction:A process in which one or more electrons are transferred between reaction partners. • The driving force of this reaction is the decrease in electrical potential. Mg(s) + I2(g) MgI2(s) Oxidation :Mg0 Mg2+ + 2 electrons Reduction: I20 + 2 electrons  I21-

41. Quiz Which of the following is not an acid-base neutralization reaction? • HCl(aq) + NaOH(s) NaCl(aq) + H2O(l) • 2 HF(aq) + Mg(OH)2(aq) MgF2(aq) + 2 H2O(l) • Pb(NO3)2(aq) + 2 KI(aq)  PbI2 (s) + 2 KNO3(aq)

42. Oxidation–Reduction Reactions • Redox reactions are those involving the oxidation and reduction of species. • Oxidation and reduction mustoccur together. They cannot exist alone. Fe2+ + Cu0 Fe0 + Cu2+ Reduced: Iron gained 2 electrons Fe2+ + 2 e  Fe0 Oxidized: Copper lost 2 electrons Cu0 Cu2+ + 2e • Remember that electrons are negative so if you gain electrons your oxidation # decreases and if you lose electrons your oxidation # increases

43. Oxidation–Reduction Reactions Fe2+ + Cu0 Fe0 + Cu2+ • Fe2+ gains electrons, is reduced, and we call it an oxidizing agent • Oxidizing agent is a species that can gain electrons and this facilitates in the oxidation of another species. (electron deficient) • Cu0 loses electrons, is oxidized, and we call it a reducing agent • Reducing agent is a species that can lose electrons and this facilitates in the reduction of another species. (electron rich)

44. Quiz Which is a reduction half reaction? • Fe  Fe2+ + 2e • Fe2+ Fe3+ + 1e • Fe  Fe3+ + 3e • Fe3+ + 1e  Fe2+

45. Example: Oxidation–Reduction Reactions For each of the following, identify which species is the reducing agent and which is the oxidizing agent. • Ca(s) + 2 H+(aq) Ca2+(aq) + H2(g) • 2 Fe2+(aq) + Cl2(aq) 2 Fe3+(aq) + 2 Cl–(aq) C) SnO2(s) + 2 C(s) Sn(s) + 2 CO(g)

46. Oxidation–Reduction Reactions

47. Balancing Redox Reactions • Half-Reaction Method:Allows you to focus on the transfer of electrons. This is important when considering batteries and other aspects of electrochemistry. • The key to this method is to realize that the overall reaction can be broken into two parts, or half-reactions. (oxidation half and reduction half)

48. Balancing Redox Reactions • Balance the following using the half reaction method Mg(s) + N2(g) Mg3N2(s) Ca(s) + Cl(g) CaCl2(s) FeI3(aq) + Mg(s)  Fe(s) + MgI2(aq) H2(g) + Ag+(aq)  Ag(s) + H+(aq)