1 / 46

Chapter 8

Chapter 8. Chemical Equations. Parts of a Chemical Equation. O 2 (g) + Fe(s)  Fe 2 O 3 (s) Products – chemicals produced in a chemical reaction. ( products) Reactants – chemicals reacted together in a chemical reaction. Products = Fe 2 O 3 Reactants = O 2 and Fe

asanderson
Download Presentation

Chapter 8

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 8 Chemical Equations

  2. Parts of a Chemical Equation • O2(g) + Fe(s)  Fe2O3(s) • Products – chemicals produced in a chemical reaction. (products) • Reactants – chemicals reacted together in a chemical reaction. • Products = Fe2O3 • Reactants = O2 and Fe •  - yields or produces

  3. Parts of a Chemical Equation • O2(g) + Fe(s)  Fe2O3(s) • (g) = gas • (s) = solid • (l) = liquid • examples: water, Hg, Br2, organics • (aq) = aqueous • A homogeneous solution. Most of what we use in class

  4. Parts of a Chemical Equation • O2(g) + Fe(s)  Fe2O3(s) • Why O2 not O and why Fe not Fe2? • Oxygen is diatomic. • Meaning that in order to be at its lowest energy, oxygen will bond with other oxygen. • Diatomic Elements – H, N, O, F, Cl, Br, I and probably (At) • Trivial others: S8 and P4

  5. Let’s Sing There are seven Diatomic Elements That you must know. (repeat) Fluorine, Chlorine Bromine, Iodine Nitrogen, Oxygen Hydrogen (repeat)

  6. Other Symbols Things on the reaction arrow • Catalyst – speeds up a reaction. • Inhibitor – slows down a reaction • Neither are parts of the products or reactants. • Energy – in the form of heat, light or electricity. Probably heat! NH4Cl Δ

  7. Example 1 • When solid mercury (II) sulfide is heated with oxygen, liquid mercury metal and gaseous sulfur dioxide are produced.

  8. Example 2 • Oxygen gas can be made by heating potassium chlorate in the presence of the catalyst manganese dioxide. Potassium chloride is left as a solid residue.

  9. Balancing Equations • Balanced Equation – an equation that has the same number and type of atoms on the product and reactant sides. • We balance equations by putting in coefficients. • A coefficient is a multiplier. • You MAY NOT change subscripts, unless they are yours and you made a mistake initially when you wrote the equation.

  10. Why do we balance equations? • Law of conservation of matter – matter is neither created nor destroyed.

  11. Example 3 Δ KClO3 O2 + KCl K Cl O

  12. Example 4 Al + O2 Al2O3 Al O

  13. Example 5 AgNO3 + Cu  Cu(NO3)2 + Ag Ag N O Cu

  14. Example 5 (again) AgNO3 + Cu  Cu(NO3)2 + Ag Ag NO3 Cu

  15. General Hints • In general you will want of get rid of odd numbers. • Also, work with combined elements before working on lone elements.

  16. Types of Reactions Chapter 8

  17. Five Common Types of Reactions • Synthesis • (Direct combination or combination) • Decomposition • Single Replacement • (Single Displacement) • Double Replacement • (Double Displacement) • Combustion

  18. Synthesis • Two or more elements or compounds combine to form a new compound. • General Form: A + B  AB

  19. Specific Examples • Cl2 + Al 

  20. Specific Examples • Sr + N2 

  21. Decomposition • A compound breaks down into two or more elements and/or compounds. • General Form: • AB  A + B

  22. Specific Examples • MgO 

  23. Specific Examples • CoBr3 

  24. Specific Examples • Carbonates break down into their metal oxide and carbon dioxide. • Na2CO3

  25. Specific Examples • Hydroxides break down into their metal oxide and water. • NaOH 

  26. Single Replacement • An element replaces another element inside a compound. • General Form: A + BC  B + AC or A + BC  C + BA

  27. Activity Series of Metals • An activity series of metals needs to be consulted to determine if a reaction will take place or not. • A more reactive metal will replace a less reactive metal in a compound. • Less commonly, there is also an activity series of non-metals as well.

  28. Activity Series of Metals Lithium Most Reactive Potassium Barium Calcium Sodium Magnesium Aluminum Zinc Iron Nickel Tin Lead Hydrogen Copper Mercury Silver Gold Least Reactive

  29. Activity Series of Halogens Fluorine Most Reactive Chlorine Bromine Iodine Least Reactive

  30. Specific Examples • Au + NaCl 

  31. Specific Examples • Cu + AgNO3 

  32. Specific Examples • Li + H2O 

  33. Specific Examples • Al + Mg(NO3)2 

  34. Double Replacement • One element out of a compound replaces another element out of a compound. • General Form AB + CD  AD + CB

  35. How do I know a reaction will occur? • That is a topic for our next section, for this chapter we will assume they all work. • Live the LIE!

  36. Evidence of a Double Replacement Reaction • A precipitate forms. • Insoluble solid • A gas is produced. • Water is produced.

  37. Specific Examples • KCl + AgNO3

  38. Specific Examples • HCl + NaOH 

  39. Specific Examples • Na2CO3 + Cd(NO3)2 

  40. Specific Examples • AlPO4 + Zn(ClO2)2

  41. Specific Examples • HBr + K2CO3

  42. “Tooth Fairies” • NH4OH = NH3 + H2OH2CO3 = H2O + CO2

  43. Combustion • A hydrocarbon reacts with oxygen to produce carbon dioxide and water. Heat and light are also produced. • General Form: • CxHyOz + O2 CO2 + H2O

  44. Specific Examples • Combusting propane (C5H12) • C5H12 + O2 CO2 + H2O

  45. Specific Examples • Combusting sugar (C12H22O11) • C12H22O11 + O2 CO2 + H2O

  46. Specific Examples • Tricky Balancing • CH5 + 

More Related