Chapter Eight Chemical Reactions

1. Chapter Eight Chemical Reactions

2. Section 8.1Chemical Equations

3. Chemical Reactions • A chemical reaction is simply a chemical change • Atoms rearrange themselves to form new compounds

4. Chemical Indicators • Production of a gas (bubbles) • Production of a solid (precipitate) • Color Change • Production of heat/light (energy) • Noise (ex: sizzling) • Only the evidence of a NEW substance after chemical analysis is 100% proof of chemical reaction

5. Chemical Equations • A systematic way to write a chemical reaction using symbols • Shows compounds involved in the reaction • States of matter • Changes in energy (endo or exo) • Quantifies reaction (ratios between compounds)

6. Chemical Equations “and” (Separates reactants or products) “Yields” or “Produces” Ex: Hydrogen gas and oxygen gas react together to produce water Ex: 2H2(g) + O2(g) 2H2O(l) Reactants Products

7. Another Example… Cu(s) + 2AgNO3(aq) 2Ag(s) + Cu(NO3)2(aq) ΔH = -484 kJ States of matter (s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous Coefficients Tell us the mole ratio between reactants/products. Enthalpy Tells if the reaction gives off energy (exo) or needs energy (endo)

8. Some Special Notes • At standard conditions, here are the following states of matter: • Metal atoms are solids • Ionic Compounds are typically solid or aqueous (dissolved in water) • Nonmetals are gases • Bromine/Mercury are liquid

9. Some Special Notes • Metals are listed as a single element • (i.e. Iron  Fe) • Carbon can be listed as C(graphite) or C(diamond) • Remember diatomics (i.e. F2, Cl2, H2) • Phosphorus can be P4 and Sulfur can be S8 • A species written over the arrow indicates a “helping” substance, though doesn’t affect chemicals involved • Ex: Δ H2O

10. Types of Reactions • Six general types of reactions • Synthesis/Combination • Decomposition • Combustion • Single Replacement (Displacement) • Double Replacement (Displacement) • Acid/Base Neutralization

11. Synthesis/Combination • Reaction between two or more substances (reactants) and making ONE PRODUCT (compound) • Form: A + B  AB • Example: 2Na + Cl2  2NaCl

12. Decomposition • Reaction of ONE REACTANT breaking down into its elementary substances • Form: AB  A + B • Example: H2O  H2 + O2

13. Combustion • Carbon-based compound reacting w/ oxygen (burning) • “CXY” could be a hydrocarbon or also contain oxygen • Products are ALWAYS CO2 and H2O • Form: CXY + O2 CO2 + H2O • Example: CH4 + 2O2  CO2 + 2H2O

14. What are the following reactions? • What type of reaction are the following? • N2 + H2 NH3 • C2H8 + O2  CO2 + H2O • H2 + O2  H2O • KClO3  KCl + O2 • Ag2O  Ag + O2 • S8 + O2  SO3 • CH4 + O2  CO2 + H2O • H2O  H2O + O2

15. Single Replacement • Chemical reaction where one element replaces another element in a compound • “Like” replaces “like” (ex: metals switch) • An ELEMENT and a COMPOUND for REACTANTS • Form: A + BC  B + AC • Example: Al + CuCl2  AlCl3 + Cu

16. Single Displacement A B C

17. Not Every SR reaction will happen • You must look at the Activity Series to determine if a cationic replacement reaction will happen. • Ex: Zn(s) + HCl(aq)  H2 + ZnCl2 Cu(s) + HCl(aq)  No Reaction

18. Double Displacement/Precipitation • Chemical reaction where two elements (or polyatomic ions) in different compounds switch places • TWO COMPOUNDS for REACTANTS • Form: AB + CD  CB + AD • Example: Na2SO4 + AlCl3  Al2(SO4)3 + NaCl

19. Double Displacement/Precipitation A B C D

20. Acid Base Neutralization • Special type of Double Replacement • Where an Acid (containing H+) and a Base (containing OH-) react to produce • A Salt (ionic compound) and • Water (HOH or H2O) • Ex: HCl + NaOH  NaCl + HOH (H2O)

21. What type is it? • Determine the type of reaction. **Predict the products for the starred reactions • **Na2S + ZnCl2 • Mg(OH)2  • **H2 + Cl2  • C6H6 + O2  • K + Fe(OH)3  • F2(aq) + KI(aq)  • **HF(aq) + NaOH(aq)  • **C10H8 + O2 

22. Balancing Equations H2 + O2 H2O Law of Conservation of Matter/Mass: Matter (atoms) cannot be created or destroyed In other words, same # of atoms on both sides Ex: Balance the above chemical equation

23. DIVIDE & CONQUER! Steps for Balancing Equations Step 1: List elements on both sides • Step 2: Write down the number of atoms for all the elements • Step 3: Use coefficients to balance atoms • DO NOT CHANGE THE SUBSCRIPTS!!!!! Ex: __FeS +__HCl  __FeCl2 + __H2S

24. Tips That Make Balancing Easier • Here are some tips to balancing some difficult equations • Are there any “2:3” ratios? Do those elements first • Ex: AlCl3 + ZnBr2 AlBr3 + ZnCl2 • Balance metals first • If you have the same polyatomic on both sides, treat the polyatomic as ONE • Ex: Na3PO4 + CaCl2  Ca3(PO4)2 + NaCl • Na2CO3  Na2O + CO2 • Try to balance Oxygen and Hydrogen at the end.

25. Tips Specific to Combustion • Balance carbons first • Balance hydrogens second • Balance oxygens last • If you end up having an odd # of oxygens on the product side, put a decimal coefficient in front of oxygen on reactant side that will make it balance, THEN DOUBLE EVERY COEFFICIENT to make whole #s • Ex: ___C2H6 + ___O2 ___CO2 + ___H2O

26. Some practice • __Li3PO4 + __SrCl2  __Sr3(PO4)2 + __LiCl • __Fe2O3 + __H2SO4 __Fe2(SO4)3 + __H2O • __C5H12 + __O2 __CO2 + __H2O

27. Group Quiz #1 • Balance the following equations and state what kind of reaction it is: • _____ N2 + _____ O2 _____ N2O • ____KClO3  ____KCl + ____ O2 • __ Al(OH)3 + __ H2SO4 __ Al2(SO4)3 + __ H2O

28. Section 8.2Combustion Analysis

29. Combustion Analysis • Determining the empirical formula of an organic compound

30. Steps • 1) Convert your known quantities of CO2 and H2O into grams of C & H only • 2) Determine how many grams of oxygen (if any) are present in the compound as well. • 3) Turn your grams of C, H, & O into moles of C, H, & O, respectively. • 4) Compare your 3 answers from step 3 and divide ALL 3 answers by the lowest # you get (this is your ratio, and therefore your subscripts)

31. Example • 18.8 g of an unknown organic substance produced 27.6 g of carbon dioxide and 11.3 g of water. • What is the empirical formula?

32. Section 8.3Calculations with Balanced Chemical Equations

33. Stoichiometry • Used to predict amount of product produced or how much reactants were used • Remember: Coefficients tell us MOLE ratio N2(g) + 3H2(g)  2NH3(g) • How many moles of ammonia are created when 2 moles of nitrogen are used? 6 moles of hydrogen? • How many moles of nitrogen are needed to make 4 moles of ammonia?

34. Stoichiometry • What mass of lithium nitride is produced when 75.0 g of lithium metal react with excess nitrogen? 6Li(s) + N2(g) 2Li3N(s) • What is the mass of O2 necessary to react with 5.71 g Al? 4Al(s) + 3O2(g)  2Al2O3(s)

35. Group Quiz #2 • Given the following equation: 2 KClO3 ---> 2 KCl + 3 O2 • How many moles of O2 can be produced by letting 12.00 moles of KClO3 react? (watch sig figs!!!) • Given the following equation: 2 K + Cl2 ---> 2 KCl • How many grams of KCl are produced from 2.50 g of K and excess Cl2 ?

36. Section 8.4Limiting Reactants

37. Limiting Reactants Although we want all reactants to be completely used up, in reality there is usually one reactant that is used up before the other and thus LIMITS the reaction and how much can be produced.

38. Limiting Reactants • Limiting Reactant: The reactant that is completely consumed (runs out first) • Limits the reaction • Nothing is left over • Excess Reatant: The reactant left over • Not completely consumed

39. Stoichiometric Steps • Though every stoichiometry problem is unique, the same general steps are used: • Convert given chemical(s) into MOLES • Use MOLE RATIO to convert from given chemical to desired chemical *After this step is where you would determine limiting reactant IF it applies to the problem • Convert desired chemical into DESIRED UNIT (ex: grams, liters, etc.)

40. Example Problems • How many moles of NH3 can be produced from the combination of 3.0 moles N2 and 1.5 moles H2? • What mass of water is produced when you react 12.4 g of H2 with 13.5 g of oxygen? • If you have 11.5 L of nitrogen reacting with 16.5 L of hydrogen, what volume of ammonia is produced at STP? (At STP 1 mole of any gas = 22.41 L)

41. Group Quiz #3 • If you have 1.22 g of O2 reacting with 1.05 g of H2, what mass of water will you produce?

42. Reaction Yields • When amount of product is calculated using stoichiometry, this is known as theoretical yield • What you theoretically SHOULD get • When you perform the experiment in the lab, you typically get less, this is known as actual yield

43. Percent Yield Actual Yield x 100 = % Yield Theoretical Yield • Ex: Magnesium burns in air. If you burn 6.73 grams of Magnesium: • How many grams of product would you produce? • What would be your percent yield if you did the experiment and got 10.7g of the product?

44. Group Quiz #4 • Suppose you react sulfur dioxide with oxygen to produce sulfur trioxide. If you react 12.4 g of sulfur dioxide with 3.45 g of oxygen, how many grams of sulfur trioxide will you produce? • If you did this lab and yielded 13.4 grams of sulfur trioxide, what is your percent yield?

45. Section 8.5Periodic Trends in Reactivity of the Main Group Elements

46. Trends of Reactivity • Group 1A metals tend to be highly reactive toward oxygen, water, and acid. • M + H2O  MOH + H2 • Lithium forms oxides (Li + O2  Li2O) • The other alkali metals form oxides or peroxides (O22- ion) (Na + O2  Na2O2) • K, Rb, and Cs can also form superoxides (containing the superoxide ion O2-) • Group 2A metals are less reactive

47. Group 2A Continued • Be doesn’t react with water, Mg reacts slowly to steam, and Ca, Sr & Ba react strongly with cold water to form hydrogen gas and hydroxides • Ca + H2O  Ca(OH)2 + H2 • React w/ oxygen (Be and Mg only at high temps) • React with acids form hydrogen gas

48. 3A elements • Boron (a metalloid) is unreactive to oxygen and water; all other elements are metals and tend to be reactive • Aluminum will readily form aluminum oxide when exposed to air (oxygen)