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Chemical Reactions

Physical Science Chapter 13. Chemical Reactions. Types of Reactions. There are four types of chemical reactions we will talk about: Combination reactions _____________ reactions Single replacement reactions ________________ reactions

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Chemical Reactions

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  1. Physical Science Chapter 13 Chemical Reactions

  2. Types of Reactions • There are four types of chemical reactions we will talk about: • Combination reactions • _____________ reactions • Single replacement reactions • ________________ reactions • You need to be able to identify the type of reaction and predict the product(s)

  3. Steps to Writing Reactions • Some steps for doing reactions • Identify the type of reaction • Predict the product(s) using the type of reaction as a model • Balance it Don’t forget about the diatomic elements! (BrINClHOF) For example, Oxygen is O2 as anelement. In a compound, it can’t be a diatomic element because it’s not an element anymore, it’s a compound!

  4. 1. Combination reactions • Combination reactions occur when two substances (generallyelements) combine and form a compound. (Sometimes these are called synthesis or addition reactions.) reactant + reactant  1 product • Basically: A + B  AB • Example: 2H2 + O2  2H2O • Example: C+ O2  CO2

  5. Combination Reactions • Here is another example of a combination reaction

  6. Practice • Predict the products. Write and balance the following combination 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) 

  7. 2. Decomposition Reactions • Decomposition reactions occur when a compound breaks up into the elements or in a few to simpler compounds • 1 Reactant  Product + Product • In general: AB  A + B • Example: 2 H2O  2H2 + O2 • Example: 2 HgO  2Hg + O2

  8. Decomposition Reactions • Another view of a decomposition reaction:

  9. Decomposition Exceptions • Carbonates and chlorates are special case decomposition reactions that do not go to the elements. • Carbonates (CO32-) decompose to carbon dioxide and a metal oxide • Example: CaCO3  CO2 + CaO • Chlorates (ClO3-) decompose to oxygen gas and a metal chloride • Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2 • There are other special cases, but we will not explore those in Chemistry I

  10. 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) 

  11. Chemical Reaction Review • In each of the following animations determine if a chemical reaction did occur • On what evidence did you base your decision Animation Link

  12. Chemical Reaction Review • Identify the type of reaction and whether it is exothermic or endothermic • Reaction #1 • Reaction #2 • Write the balanced decomposition reaction for water. Is there more, less, or the same amount of hydrogen gas produced when compared to oxygen? Based on the chemical equation what is the ratio of hydrogen to oxygen gas produced? Animation check of prediction

  13. Chemistry in Auto Safety • Airbag Deployment Video • Airbag Chemistry

  14. Practice Identify the type of reaction for each of the following synthesis or decomposition reactions, and write the balanced equation: N2(g) + O2(g) BaCO3(s)  Co(s)+ S(s)  NH3(g) + H2CO3(aq)  NI3(s)  Nitrogen monoxide (make Co be +3)

  15. Energy and Rate of Reaction • Arrhenius believed that for molecules to react upon collision, they must become "activated," known as the Activation Energy.

  16. Heat Flow During a Reaction • Exothermic • Combustion of gasoline • Cellular Respiration • Endothermic • Example: Photosynthesis

  17. Factors Influencing Reaction Rates • Increase surface area of reactants • Increase temperature • Increase reactant concentration • Addition of a catalyst

  18. Catalysts • It is not always practical or convenient to increase reaction rates by increasing the temperature or other reaction factors • A Catalyst is a substance that speeds up a reaction without being consumed during by the reaction • Catalysts are used to speed up reactions without changing the temperature

  19. Catalysts at Work • Catalysts are used in a huge variety of ways because they can enhance reaction rates by many orders of magnitude! • In general, they work by lowering the activation energy to a reaction.

  20. Catalysts at Work • Activity: • Draw an activation energy graph for a chemical reaction • Draw an activation energy curve on the graph that would represent the impact of using a catalyst

  21. Impact of a Catalyst

  22. Catalyst Example 2N2O(g)  2N2(g) + O2(g)

  23. Catalyst Example - Enzymes

  24. The Chemistry of Acids and Bases Physical Science Chapter 13

  25. Acid and Bases

  26. Acid and Bases

  27. Acid and Bases

  28. Acids Have a sour taste. Vinegar is a solution of acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Bases Have a bitter taste. Feel slippery. Many soaps contain bases.

  29. Some Properties of Acids • Produce H+ (as H3O+) ions in water (the hydronium ion is a hydrogen ion attached to a water molecule) • Taste sour • Corrode metals • Electrolytes • React with bases to form a salt and water • pH is less than 7 • Turns blue litmus paper to red “Blue to Red A-CID”

  30. Acid Nomenclature Review Binary  Ternary An easy way to remember which goes with which… “In the cafeteria, youATEsomethingICky”

  31. Acid Nomenclature Flowchart

  32. Acid Nomenclature Review • HBr (aq) • H2CO3 • H2SO3 hydrobromicacid  carbonicacid  sulfurousacid

  33. Name ‘Em! • HI (aq) • HCl (aq) • H2SO3 • HNO3 • HIO4

  34. Some Properties of Bases • Produce OH- ions in water • Taste bitter, chalky • Are electrolytes • Feel soapy, slippery • React with acids to form salts and water • pH greater than 7 • Turns red litmus paper to blue “Basic Blue”

  35. Some Common Bases NaOH sodium hydroxide lye KOH potassium hydroxide liquid soap Ba(OH)2 barium hydroxide stabilizer for plastics Mg(OH)2 magnesium hydroxide “MOM” Milk of magnesia Al(OH)3 aluminum hydroxide Maalox (antacid)

  36. Acid/Base definitions • Definition #1: Arrhenius (traditional) Acids – produce H+ ions (or hydronium ions H3O+) Bases – produce OH- ions (problem: some bases don’t have hydroxide ions!)

  37. Arrhenius acid is a substance that produces H+ (H3O+) in water Arrhenius base is a substance that produces OH- in water

  38. The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion.Under 7 = acid 7 = neutralOver 7 = base

  39. pH of Common Substances

  40. Calculating the pH pH = - log [H+] (Remember that the [ ] mean Molarity) Example: If [H+] = 1 X 10-10pH = - log 1 X 10-10 pH = - (- 10) pH = 10 Example: If [H+] = 1.8 X 10-5pH = - log 1.8 X 10-5 pH = - (- 4.74) pH = 4.74

  41. Try These! Find the pH of these: 1) A 0.15 M solution of Hydrochloric acid 2) A 3.00 X 10-7 M solution of Nitric acid

  42. pH testing • There are several ways to test pH • Blue litmus paper (red = acid) • Red litmus paper (blue = basic) • pH paper (multi-colored) • pH meter (7 is neutral, <7 acid, >7 base) • Universal indicator (multi-colored) • Indicators like phenolphthalein • Natural indicators like red cabbage, radishes

  43. Paper testing • Paper tests like litmus paper and pH paper • Put a stirring rod into the solution and stir. • Take the stirring rod out, and place a drop of the solution from the end of the stirring rod onto a piece of the paper • Read and record the color change. Note what the color indicates. • You should only use a small portion of the paper. You can use one piece of paper for several tests.

  44. pH meter • Tests the voltage of the electrolyte • Converts the voltage to pH • Very cheap, accurate • Must be calibrated with a buffer solution

  45. pH indicators • Indicators are dyes that can be added that will change color in the presence of an acid or base. • Some indicators only work in a specific range of pH • Once the drops are added, the sample is ruined • Some dyes are natural, like radish skin or red cabbage

  46. 3. Single Replacement Reactions • Single Replacement Reactions occur when one element replaces another in a compound. • A metal can replace a metal (+) OR a nonmetal can replace a nonmetal (-). • element + compound product + product A + BC  AC + B (if A is a metal)OR A + BC  BA + C (if A is a nonmetal) (remember the cation always goes first!) When H2O splits into ions, it splits into H+ and OH- (not H+ and O-2 !!)

  47. Single Replacement Reactions • Another view:

  48. Single Replacement Reactions • Write and balance the following single replacement reaction equation: • Zinc metal reacts with aqueous hydrochloric acid Zn(s) + HCl(aq) ZnCl2 + H2(g) Note: Zinc replaces the hydrogen ion in the reaction 2

  49. Single Replacement Reactions • Sodium chloride solid reacts with fluorine gas NaCl(s) + F2(g)  NaF(s) + Cl2(g) Note that fluorine replaces chlorine in the compound • Aluminum metal reacts with aqueous copper (II) nitrate Al(s)+ Cu(NO3)2(aq) 2 2

  50. 4. Double Replacement Reactions • Double Replacement Reactions occur when a metal replaces a metal in a compound and a nonmetal replaces a nonmetal in a compound • Compound + compound  product + product • AB + CD  AD + CB

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