Types of Chemical Reactions

1. Types of Chemical Reactions 6.2 Assigning oxidation numbers for individual atoms of monatomic and polyatomic ions 6.3 Identifying the nomenclature of ionic compounds, binary compounds, and acids 6.4 Classifying chemical reactions as composition, decomposition, single replacement, or double replacement AHSGE Reading 4.2 Demonstrate the ability to preview and predict. Other AOD C.6.1 Define stoichiometry, reactants, and products.

2. What are some types of chemical reactions? • Synthesis • Decompostion • Single Replacement/Single Displacement • Double Replacement/Double Displacement • Combustion • Complete combustion yields ________. • Incomplete combustion yields __________.

3. “Classes” of Reactions • Precipitation reactions • Acid-base reactions • Oxidation-reduction reactions • Almost all reaction types can be put into one of these three “classes”.

4. Precipitation Reactions • What do you think this reaction would involve? • AKA, double replacement/double displacement • Example: K2CrO4(aq) + Ba(NO3)2(aq) → ??? • What does (aq) mean? • What are some other designations? • What is the name of K2CrO4? Ba(NO3)2? • What would the products be in a double displacement reaction? • Would either product be a solid (precipitate)?

5. Three Ways to Write the Reaction Equation • Molecular equation: shows the overall balanced reaction (reactants and products) • Complete ionic equation: shows a balanced equation of the actual FORMS of the reactants and products in the reaction • Net ionic equation: balanced equation showing ONLY the ions involved in the reaction. • Ions that remain in solution (as ions) are not shown. (“spectator ions”)

6. K2CrO4(aq) + Ba(NO3)2(aq) → ??? • Molecular equation: K2CrO4(aq) + Ba(NO3)2(aq) → 2 KNO3 (aq) + BaCrO4(s) • Complete ionic equation: 2K+(aq) + CrO42-(aq) + Ba2+(aq) + 2NO3-(aq) → BaCrO4(s) + 2K+(aq) + 2NO3-(aq) • Net ionic equation: • What is the only CHANGE in this rxn? • CrO42-(aq) + Ba2+(aq) → BaCrO4(s)

7. Examples • See Sample Exercise 4.9, p.155. • P. 182 (29a): • BaCl2(aq) + Na2SO4(aq) →BaSO4(s) + 2NaCl(aq) • Net ionic equation: Ba2+(aq) + SO42-(aq) → BaSO4(s) • What happened to the “Cl2” and the “Na2” part of the molecular equation? • What would the complete ionic equation look like?

8. Assignment • P.182: 30, 34, 36 • Special note on #34: You are MAKING UP a chemical reaction that will produce the specified solid. • It should be a precipitation (double displacement) reaction.

9. Acid-Base Reactions • Acid-Base definitions: • Arrhenius: • Acid – a substance that produces H+ when dissolved in water • Base - a substance that produces OH- when dissolved in water • Bronsted-Lowry: • Acid – a proton donor • Base – a proton acceptor

10. Predicting Acid-Base Reactions • What element do acid chemical formulas start with? • Most bases contain hydroxide ions. • Using Arrhenius’ definitions, H+ + OH-→ ?? • So Acid-Base reactions yield _________.

11. Example: NaOH + HCl • What is the acid? Base? • Molecular equation: NaOH +HCl → NaCl + H2O (What’s missing from this equation???) • Complete ionic equation: Na+(aq)+ OH-(aq)+ H+(aq) + Cl-(aq) → Na+(aq)+ Cl-(aq) + H2O(??) • Net ionic equation: ???? • Assignment: P.183 (46, 48)

12. Oxidation-Reduction (Redox) Reactions • Def: reactions in which one or more electrons are transferred • Example: 2Na(s) + Cl2(g) → 2NaCl(s) • What is the “charge” on Na(s)? • On Cl in Cl2? • On Na in Nacl? • On Cl in NaCl?

13. Redox Reactions, continued… • Includes most reactions involving energy production (i.e., photosynthesis, combustion). • Example: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + energy • How do we calculate the charges on elements in covalent molecules to determine a transfer of electrons?

14. Oxidation States (aka, Oxidation Numbers) • Def: the imaginary charges the atoms in a covalent molecule would have if the shared electrons were assigned to the atom with the most electronegativity • Rules for Assigning Oxidation States (Table 4.2, p.167)

15. Rules for Assigning Oxidation States • Give the oxidation states for all the atoms in the following molecules: • O2 • CH4 • CO2 • H2O • SF6 • NO3- • Fe3O4

16. CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + energy • For which atoms did the oxidation state change? • C went from -4 to +4. • This change shows that each C atom LOST 8 electrons. • O went from 0 to -2. • This shows that each O atom ______ ____ electrons.

17. Oxidation and Reduction • Oxidation: an INCREASE in oxidation state (LOSES e-s) • Reduction: a DECREASE in oxidation state (GAINS e-s) • If an atom is oxidized, it is called the reducing agent (electron donor). • If an atom is reduced, it is called the _______ agent (electron acceptor). • OIL RIG: Oxidation Involves Loss; Reduction Involves Gain

18. CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + energy • Which element is oxidized in this equation? • Which element is reduced? • Which compound is the oxidizing agent? • Which compound is the reducing agent? • Example: 2Al(s) + 3I2(s) → 2AlI3(s) • Assignment: Pp.183-184 (58, 60, 62)

19. Balancing Redox Reactions • If it does not balance easily using the normal rules, use the half-reaction method. • Uses two separate reaction equations: • Oxidation half-reaction: balance the substance being oxidized • Reduction half-reaction: balance the substance being reduced • Example: Ce+4(aq) + Sn+2(aq) → Ce+3(aq) + Sn+4(aq) • Sn+2(aq) → Sn+4(aq) ---- __________ half-reaction • Ce+4(aq) → Ce+3(aq) --- __________ half-reaction • But how would you balance this????

20. Actual Balancing Steps • Varies by whether the reaction occurs in acidic or basic solutions! • Acidic solutions (solns) follow steps on p.172 • Basic solns follow acidic steps, THEN add steps 2-4 on p.177.

21. Steps in Acidic Solutions • Write the two half-reactions. • For each half-reaction! • Balance all elements EXCEPT H and O. • Balance O by adding H2O. • Balance H by adding H+. • Balance any charge by adding e-s. • Multiple either of the half-reactions by an integer to have equal number of e-s in both half-reactions, if needed. (You want the e-s to cancel in the next step.) • Add the half-reactions, cancelling any identical species. • Check that both elements AND charges are balanced in final equation.

22. Example: Zn(s) + NO3-(aq) → Zn+2 + NO2(g) • Half-reactions: • Oxidation: Zn(s) → Zn+2(aq) • Reduction: NO3-(aq) → NO2(g) • Balancing steps: • Already done  • NO3-(aq) → NO2(g) + H2O(l) • 2H+(aq) +NO3-(aq) → NO2(g) + H2O(l) • e- + 2H+(aq) + NO3-(aq) → NO2(g) + H2O(l) AND Zn(s) → Zn+2(aq) + 2e-

23. Example (continued..): Zn(s) + NO3-(aq) → Zn+2 + NO2(g) • Which half-reaction do we need to multiply, and by what integer? • 2 [e- +2H+(aq) + NO3-(aq) → NO2(g) + H2O(l)] = 2e- +4H+(aq) + 2NO3-(aq) → 2NO2(g) + 2H2O(l) • Zn(s) → Zn+2(aq) + 2e- • Add the two equations together: 4H+(aq) + 2NO3-(aq) + Zn(s) → 2NO2(g) + 2H2O(l) + Zn+2(aq) • Check the balance of elements and charges!

24. Try This One: MnO4-(aq) + Zn(s) → Mn+2(aq) + Zn+2(aq) • Write the half-reactions. • Balancing steps: • Mn and Zn • Add __________ to balance O. • Add __________ to balance H. • Add __________ to balance any charges. • Multiply to get equal number of ____ in both half-reactions. • Add them together. • Check the elements and charges!

25. Steps in Basic Solutions • Start with the steps used for acidic solutions. • Add OH- (equal to the final amount of H+) to BOTH sides of the equation. • Eliminate equal amounts of water from both sides of the equation. • Check balance of elements and charges!

26. Example (#65(b)): Cl2(g) → Cl-(aq) + OCl-(aq) • Acidic steps get you to: 2H2O(l) + 2Cl2(g) → 2Cl-(aq) + 2OCl-(aq) + 4H+(aq) • Add OH- to BOTH sides. How many??? 4OH-(aq) + 2H2O(l) + 2Cl2(g) → 2Cl-(aq) + 2OCl-(aq) + (4H+(aq) + 4OH-) • 4OH-(aq) + 2Cl2(g) → 2Cl-(aq) + 2OCl-(aq) + 2H2O(l) • Check balance of elements and charges! • What is one more thing we can do to this final equation????

27. Assignment • P.184: 64, 66