Chemical Reactions

1. Chemical Reactions

2. Chemical Reactions Reactants: A starting substance in a chemical reaction. Products: A substance formed in a chemical reaction.

3. Reactants Products Reactants are always on the left of the arrow. Products are always on the right of the arrow. What does the arrow mean? Yields or produces

4. For Review: Law of Conservation of Mass Are atoms ever created or destroyed in a chemical reaction? NO!!!

5. Definitions Chemical Equation: The arrow separates the formulas of the reactants from the formulas of the products. Skeleton Equation: A chemical equation that doesn’t indicate the relative amounts of the reactants and products involved. Symbols you will see: (s), (l), (g), and/or (aq)

6. More Definitions: Catalyst: Substance that speeds up the rate of a reaction but is not used up in the reaction. If a catalyst is used, the symbol or formula is written above the arrow.

7. Evidence of Chemical Reaction • Energy change • Temp increase/decrease • Light emitted • Precipitate produced • Unexpected color change • Gas released/bubbles

8. Types of Reactions There are 5 types of reactions. 1. Combination (a.k.a. synthesis) 2. Decomposition 3. Single replacement 4. Double replacement 5. Combustion

9. Combination (Synthesis) Start with TWO elements/compounds and merge to form ONE product Two elements combine A + B C Mg(s) + O2(g) MgO(s) Element/cmpd + element/cmpd Compound Example: Complete the following combination reaction: K(s) + Cl2(g) 2 2 2 2 KCl

10. Combination, cont. Nonmetal oxides react with water to form acids SO2(g) + H2O(l) H2SO3 (aq) Metallic oxides react with water to from bases CaO(s) + H2O(l) Ca(OH)2

11. Synthesis Reactions Here is another example of a synthesis reaction

12. Practice Predict the products. Write and balance the following synthesis reaction equations. Sodium metal reacts with chlorine gas Na(s) + Cl2(g)  Solid Magnesium reacts with fluorine gas Mg(s) + F2(g)  Aluminum metal reacts with fluorine gas Al(s) + F2(g) 

13. Decomposition: A compound breaks down into a two or more simpler substances. C A + B CaCO3(s) CaO(s) + CO2(g) *Most decomposition reactions require energy in the form of heat, light or electricity. Cmpd element/cmpd + element/cmpd

14. Decomposition Rxn. • Start with ONE cmpd and it breaks into TWO pieces (two elements or two smaller cmpds) • Ionic Salt  M + NM NaCl  • Carbonate salt  metal oxide + CO2 CaCO3  • Chlorate/perchlorate salt  metal chloride + O2 KClO3 

15. Decomposition Reactions Another view of a decomposition reaction:

16. Practice Predict the products. Then, write and balance the following decomposition reaction equations: Solid Lead (IV) oxide decomposes PbO2(s)  Aluminum nitride decomposes AlN(s) 

17. Single Replacement One element replaces a second element in a compound. A + BC AC + B Element + Compound Compound + Element

18. Single Displacement Rxns. • Element reacts with a cmpd to produce a new element and a new cmpd. • Metal + cmpd  new metal + new cmpd *AgNO3(aq) + Cu (s) • NM + cmpd  new NM + new cmpd Cl2(g) + NaI (aq)  • Metal + water  H2 + metal hydroxide (base) Li (s) + HOH (l)  • Use REACTIVITY SERIES to make sure these rxns will occur • Element must be more reactive than the ion in the cmpd in order to replace it

19. Single Replacement, cont. Whether or not a metal will displace another metal depends on its reactivity to other metals. A metal must be more reactive to “kick out” another metal. If it is not, there is no reaction. Activity Series of Metals. Reactivity decreases as you go down, so metals will only replace another one below it. A nonmetal can also replace another nonmetal from a compound. This is usually limited to halogen group. Activity of halogens decreases as you go down the column (F is more reactive than I).

20. Single Replacement Reactions Another view:

21. Single Replacement Practice Complete the following reactions: 1. Zn + H2SO4 2. Sn + NaNO3 3. Cl2 + NaBr ZnSO4 + H2 No Reaction NaCl + Br2

22. Single Replacement Practice Complete the following reactions: 1. Mg + KCl 2. Ag + HCl 3. l2 + KF 4. Zn + LiCl 5. Ca + PbO2

23. Double Replacement Involves an exchange of positive ions between two reacting compounds. AB + CD AD + CB Cmpd + Cmpd Cmpd + Cmpd • For a double replacement reaction to occur, one of the following is usually true: • 1. A precipitate is formed. • 2. A gas is produced and bubbles out of solution. • 3. One product is a molecular compound such as water.

24. Double Replacement Reactions Think about it like “foil”ing in algebra, first and last ions go together + inside ions go together Example: AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq) Another example: K2SO4(aq) + Ba(NO3)2(aq)  KNO3(aq) + BaSO4(s) 2

25. Practice Predict the products HCl(aq) + AgNO3(aq)  CaCl2(aq) + Na3PO4(aq)  Pb(NO3)2(aq) + BaCl2(aq)  FeCl3(aq) + NaOH(aq)  H2SO4(aq) + NaOH(aq)  KOH(aq) + CuSO4(aq) 

26. Combustion An element or a compound reacts with oxygen (O2) often producing energy as heat and light. X + O2 XO Often involves hydrocarbons: C and H C#H#, C#H#O# or C#H#OH • Complete combustion of a hydrocarbon • looks like this: • C3H8 + O2(g) CO2 + H2O

27. Combustion Reactions In general: CxHy + O2  CO2 + H2O Products in combustion are ALWAYS carbon dioxide and water. (although incomplete burning does cause some by-products like carbon monoxide) Combustion is used to heat homes and run automobiles (octane, as in gasoline, is C8H18)

28. Identify the following reactions • S + O2 → SO3 • CaCl2 + Na3PO4  Ca3(PO4)2 + NaCl • K + Cl2 → KCl • H2O2 → H2O + O2 • Na + H2O → NaOH + H2 • KClO3 → KCl + O2 • C2H6 + O2 → CO + H2O • KI + Cl2  KCl + l2 • CH4 + O2 → CO2 + H2O • N2 + H2 → NH3 • P + O2 → P4O10 • Fe(OH)3 → Fe2O3 + H2O

29. Practice Predict the products. Identify the type of reaction. S + O2  CaCl2 + Na3PO4)  Pb(NO3)2 + BaCl2  KI + Cl2  CH4 + O2  C3H8 + O2  CaCO3  Na + H2O 

30. Balancing Combustion Reactions • When balancing combustion reactions, start with Carbon and balance Oxygen last. • If you end up needing an odd number of diatomic oxygen, multiply the coefficient in front of the hydrocarbon by 2 and rebalance.

31. Here is what I mean… Balance the following equation: C3H8 + O2 CO2 + H2O C4H10 + O2 CO2 + H2O C4H10 + O2 CO2 + H2O Combustion Reactions 5 3 4 4 5 6.5 13 2 8 10

32. Combustion Practice Combust the following compounds: C6H6 CH3OH Be sure to balance these! 2 + 15 O2 12 CO2 + 6 H2O 2 + 3 O2 2 CO2 + 4 H2O

33. Balancing Equations Balanced Equation: Each side of the equation has the same number of atoms of each element. Why balance an equation? To obey the law of conservation of mass.

34. How do we balance? Use Coefficients: “The big numbers in front.” Place the coefficient in front of the symbols for the respective parts. If no number is written, it is understood to be one.

35. Rules for Balancing 1.Determine the correct formulas If states are given, include these in parenthesis. 2.Write the formula for reactants on left and products on the right. Separate with a yield sign. If there are two or more reactants or products, separate with a plus sign.

36. Rules for Balancing, cont. 3. Count the number of atoms of each element. A polyatomic ion appearing unchanged on both sides of the equation is counted as a single unit. Example: CaS(s) + NaOH(aq) Ca(OH)2(aq) + Na2S(s) Versus: NaHCO3(s)+HCl(aq) NaCl(aq)+H2O(l)+ CO2(g)

37. Rules for Balancing, cont. 4. Balance the elements one at a time. 5. Check to be sure each atom or polyatomic ion is balanced. 6. Make sure all coefficients are in the lowest possible ratio.