2. “Good morning, and welcome to introduction to chemistry.” Not the real Mr. Cooper

3. Info Class: Chemistry Instructor: Mr. Cooper Office: A112 - I’m pretty much in my classroom before and after school E-mail: gcooper@lps.orggcooper@lps.org WebPage: lsw.lps.org click on teachers and find my name.

4. Class Format 1. Daily Quizzes (D.Q.): Each day will start with a quiz over the previous days material. Clear your desk, have out a piece of paper and be ready at the beginning of class. Recommendation: Write out the question and answer and keep a running list of the quizzes on the same sheet of paper. At the end of the term you will have created a review sheet. I will not be giving you a review sheet at the end of the term. Quizzes will not be picked up but your score will be recorded on your self- evaluation sheet.

5. Class Format 2. Lecture: Lecture notes are posted on the web. You can print them off before class. This allows you to listen and formulate questions and take your own notes rather than just copying down my outline. Printing them ahead of time is not required; but if all of the class has notes ahead of time, we can spend more time on lab work, book work and individual help rather than copying notes.

6. Class Format 3. Laboratory:. One typed lab report of your choice will be submitted per unit. You choose the lab you want to do the report on. It is suggested that you choose one at the beginning of the unit so that you can get it done ahead of time. It will be due day of test. Leave 5 minutes at the end of each period for clean up. Labs are not expected to be homework. If you work diligently in class, you should get them done. If not, the labs are on-line for you to finish at home. Lab may be replaced by worksheets, group work, a video, or demonstrations

7. Class Format 4. If there is time left at the end of class you are expected to be working on book problems, which are assigned on a daily basis. See unit outline. This is also time for you to get individual help.

8. Grading Lab Books and Lab Quizzes: A simple one-subject notebook is recommended. The day before each unit test will be a unit lab quiz. It is open lab book. See ppt notes for specifics on lab book expectations.

9. Grading Quizzes and tests will be announced in advance. Tests and quizzes will be closed note and closed book. Exception: lab quizzes are open lab book. Your textbook is your first resource; so read it!!! All materials for this course are based off of your text book. Also, calculators will be allowed on tests and quizzes. It is your responsibility to provide a calculator.

10. Grading Tests and quizzes are m.c., short answer, problem solving, make you think exams; not memorizing exams(although you will need to have some things memorized.)

11. Grading Test retakes will be offered during plc time and Sat. school in the media center. You must sign up for a retake so I can get it to the media center. It is your responsibility to rearrange your schedule if you wish to take advantage of retesting.

12. Grading The grading scale is as follows: A= 90.0-100 B+= 85.0-89.9 B= 80.0-84.9 C+=75.0-79.9 C= 70.0-74.9 D+= 65.0-69.9 D= 60.0-64.9 F= Below 60.0

13. Misc. Any assignments or test missed for truancy results in 60.0% of earned grade. This is district policy. Missed labs need to be done ASAP. Late work is accepted 1 day late for 1/2 credit.

14. Misc Tardies - Building policy is followed. Cell Phones - If I see it or hear it; I can take it for the rest of the day, turn it into security, or write a referral. iPods - iPods are not to be used during instructional time or lab time. You may use them during individual work time at your desk. I reserve the right to revoke privileges.

15. Student Expectations Do your job as a student which means: 1. Bring all needed materials to class. ex) books, notebooks, writing utensils, brain, good attitude, etc. 2. Respect each students right to learn and their property. 3. Listen carefully and follow instructions given. 4. READ, STUDY, PAY ATTENTION TO DETAIL 5. No food or drink. 6. Use class time to work. 7. ASK QUESTIONS in class or see me after school for help.

16. Misc. “I do not feel obliged to believe that the same God who has endowed us with sense, reason, and intellect has intended us to forgo their use.” - Galileo Galilei (astronomer and physicist) Remember, I am working hard for you. I expect that you will work hard for me. I find it offense when at the end of the term you are begging me to round or expecting me to do you some extra credit favor when you didn’t give me your best to begin with. Mr. Cooper

17. Equipment Use Review What lab equipment is used for handling a hot beaker? What lab equipment would be used to hold a piece of metal in a flame? What piece of lab equipment is used to measure volume? A BEAKER OR FLASK IS NEVER USED AS A MEASUREMENT DEVICE!!!

18. Tirrill (Bunsen) Burner

19. How the parts work. Turning the barrel Controls type of flame (orange or blue) by opening and closing the air vent. Always use blue flame (open vent); however, vent should not be wide open for initial igniting.

20. How the parts work. Gas Flow Control Controls the height of the flame through controlling the amount of gas flowing. Use appropriate flame height. NO TORCHES.

21. Operation of the Tirrill (Bunsen) Burner Hook the hose to the gas inlet and gas jet Place spark ignitor next to top of barrel Turn on gas and ignite with sparker Make barrel and gas flow control adjustments for proper flame

22. Trouble Shooting You should be able to hear the gas flowing. If not: Check if gas flow control valve is open Check if jet valve is clogged. If so see your teacher.

23. Troubleshooting Gas attempts to light but goes out. Possible cause is: Air vent is too far open. Turn the barrel down.

24. Formatting a Lab Report Title: The word “title” is written and underlined; followed then by the name of the lab. Purpose: The word “purpose” is written and underlined; followed by the purpose of the lab. Procedure: Usually extremely detailed. You can summarize. Just a couple of sentences is fine. Procedure questions will be on quizzes. Data: The word “data” is written and underlined. For this section you will either be filling out charts or questions will be asked to help you gather data. Write out the question, underline it, leave a space, then answer the question.

25. Formatting a Lab Report Conclusion: The word “conclusion” is written and underlined. For this section you will be asked questions. Write out the question, underline it, leave a space, then answer the question. Application: Where is this concept used in the real world or in the scientific community? How does this affect your life or why is this important to have this knowledge for society or other real world application or future predicted use? –Minimum 3 sentences and Maximum 5 sentences –Must have one source to accompany this section. If you use a website please make sure you do not have a typo in the address. –No opinions. I am not your source nor are you a source. This is a research component. Do some research and quote your source or than your text. Do not use any “I” statements.

26. Sample lab report Title: Place title here. Purpose: Place the purpose here. Procedure: A couple of general sentences summarizing lab steps. Data: 1.I am writing out the question and underlining it. A space was left and question 1 is answered. 2. Another question is written out and underlined. A space was left and question 2 was answered.

27. Sample lab report Conclusion: 1.I am writing out the question and underlining it. A space was left and question 1 is answered. 2. Another question is written out and underlined. Do you see a pattern here? Application: Use or Application Source: WWW.SCIENCERULES.COM Keep answers clear and concise. Length is not important. I care about content and good communication.

28. Bad Student Example - Very Bad Application: After doing this lab, I sat and wondered how I would apply what I learned to something that expands outside of our classroom. Thinking about the candle burning sent me into deep contemplation. And then, out of complete randomness, I started thinking about our environment and the things that we burn which pollute it. I then thought of where all the statistics we hear about come from, and how the claims are substantiated. How do scientists know exactly what percent our ozone layer has deteriorated, and what percent of our atmosphere is made up harmful pollutants? Well when fossil fuels are burned, or maybe even things like wood or who knows, scientists most likely calculate the molecules given off so they can come up with these statistics. Well maybe they deal with moles or liters of gas at STP, who knows, but I ’ m sure somewhere in there scientists will have to convert from moles to molecules, or grams to moles, or grams to molecules, and in a sense that is what we have done in this lab. We found out how many moles of wax were burned over 3 minutes, and if we know what wax is made of then we can figure out what exactly was released into the atmosphere. Source: My brain.. no seriously.. my brain.

29. Acronyms KISS Keep It Simple Stupid SOP Standard Operating Procedure HUA Heard Understood Acknowledged WAG Wild Ass Guess

30. Metric Conversions grams (g) is used for mass (weight) liters (l) is used for volume meters (m) is used for distance kilo hecto deka SI deci centi milli k h dk g d c m l m

31. Metric Conversions kilo hecto deka SI deci centi milli We will use a problem solving process called dimensional analysis (tracks). Example 25.0 cg = _______ g 25.0 cg 1 g 100 cg = 0.250 g Example 0.351 hl = _______ ml 0.351 hl 100,000 ml 1 hl= 35,100 ml

32. Metric Conversions kilo hecto deka SI deci centi milli Your turns - convert the following: 15.72 g = ______ mg 15.72 g = _____ kg

33. Density = m  v D = m/v intensive property –density is the same no matter size 50 grams of gold has the same density as 150 grams of gold. Important density to remember –water is 1.0 g/ml at 4 o C –1 cm 3 = 1 ml

34. Density Problem A substance has a volume of 1.74 ml and a mass of 20.0 grams. What is the density? What is the substance? Use page 96 of your text.

35. Density Problem One more to test your algebra: What is the volume of ice in a container if the density is 0.920 g/ml and the mass is 58.39 g? 0.920 g/ml = 58.39 g V V = 63.5 ml

36. Properties of Matter Definitions Matter - anything that has mass and takes up space Mass - amount of matter an object contains Substance (pure) - matter that has a uniform composition –Ex. Sugar - C 12 H 22 O 11 –Lemonade is not a pure substance

37. Properties of Matter States of Matter (Solid) Definite shape Definite volume Is incompressible (atom or molecules can not be pushed closer together) Examples - –coal, sugar, ice, etc

38. Properties of Matter States of Matter (Liquid) Matter that flows Has a fixed volume Takes the shape of its container Incompressible Examples –Water, milk, blood, etc

39. Properties of Matter States of Matter (Gas) Matter that takes the shape and volume of its container Easily compressed Examples –Oxygen, nitrogen, helium, etc

40. States of Matter Video

41. Properties of Matter Physical Property An observed condition of the substance Physical properties help identify substances Examples include: Color Solubility Odor Density Hardness Melting point (m.p.) Boiling point (b.p.) Malleability Ductility Luster

42. Properties of Matter Physical Change A change which alters a given material without changing its composition Nothing new is made Example –Ice melting - new state of matter but substance is still H 2 O –Vapor - a substance that is in a gaseous state but liquid at room temp

43. Change of State - a physical change

44. Properties of Matter Chemical Property The ability or inability of a substance to rearrange its atoms. Example –Gasoline has the ability to react violently with oxygen

45. Physical and Chemical Properties

46. Properties of Matter Chemical Change The actual rearrangement of atoms Example –The combustion of gasoline to make carbon monoxide, carbon dioxide, carbon, water (this produces a great amount of energy)

47. Classifying a physical or chemical change Ask yourself these questions: 1. Has something new been made? –If yes than a chemical change occurred –Indicators - color change, formation of precipitate, absorption or release of energy, formation of a gas 2. What does it take to get back to the original form? –If a physical process can revert it back than the change was physical.

48. A chemical change

49. Classify the following as a physical or chemical change A sidewalk cracking Blood clotting Getting a tan Making Kool-Aid Making a hard boiled egg Plastic melting in the sun Autumn leaf colors Digestion of food The ripening of a banana Making ice cubes Milk curdling Turning on the television Making toast Mowing the grass Paint fading Grey hair

50. Categorizing our environment

51. Classifying Matter Mixtures A physical blend of two or more substances. Examples: –Beef stew, air - mixture of gases

52. Classifying Matter Mixture (Heterogeneous) Not uniform in composition One portion of the mixture is different from the composition of another portion Example: –Soil - sand, silt, clay, decayed material

53. Classifying matter Mixture ( homogeneous) Completely uniform composition Components are evenly distributed throughout the sample Example –Alloys - mixture of metals (brass, steel) AKA - solution –Example - ammonia, alloys, kool-aid

55. Classifying Matter Mixtures can be separated by physical means Examples –A spoon can separate beef stew –Sulfur and iron can be separated with a magnet Tap water is a mixture that can be separated by distillation. Distillation - a separation techniques based on the physical property of boiling points. –Liquid is boiled to produce a vapor –Then condensed to a liquid leaving impurities behind

56. Categorizing our environment

57. Classifying Matter Elements Simplest form of matter Cannot be broken down into anything else Building blocks for all other substances Examples –Hydrogen, oxygen, carbon

58. Classifying Matter Compounds Two or more elements combined through a chemical bond Can only be separated into simpler substances by chemical reactions Example –Sugar - C 12 H 22 O 11

59. Chemical and physical properties of compounds are different from their constituent elements. Examples Sugar –Carbon is black –Hydrogen is a gas –Oxygen is a gas Salt - NaCl –Na (sodium) soft metal that explodes with water –Cl (Chlorine) pale yellow-green poisonous gas

60. Classifying matter review Remember Substance - all of one kind of matter –Examples: element or compound Mixture - has more than one kind of material –Examples - two or more compounds or elements that are mixed, not chemically combined

61. Categorizing our environment

62. The anatomy of the periodic table Get out your periodic tables Know where the following are on your periodic table (p.t) Group A (representative elements) Group B Metals Nonmetals Metalloids (Semimetals) –Note - aluminum is not considered a metalloid

63. The anatomy of the periodic table Know where the following are on your periodic table (p.t) continued Transition metals Inner transition metals Alkali metals Alkaline metals Halogens Noble gases

67. Naming Molecular Compounds Molecules are made up of nonmetals Prefixes are used to represent numbers of atoms. See your text for prefixes Binary compounds end in -ide Examples Name? - Cl 2 O 8 and OF 2 Formula for? - dinitrogen tetroxide Answers -

68. Naming Molecular Compounds Your turn. Try these. Name or write the formula for: –Boron trichloride –Dinitrogen tetrahydride –N 2 O 5 PF 5 S 4 N 2 CCl 4 SO 3 H 2 O Answers

69. Take ten minutes and work a few problems on the “Naming covalent compounds” side of your worksheet.

70. Ions An atom that carries a charge The charge on the ion is called the Oxidation state or Oxidation Number Cation - positively charged atom –Metals form cations –CATions are PAWsitive Anion - negatively charged atom –Nonmetals form anions

71. Naming Cations Name the metal followed by the word ion Example –Na - sodium - neutral element –Na 1+ - sodium ion - cation of the element Another example: –Mg - magnesium Mg 2+ - Magnesium ion

72. Naming Anions Ending changes are used for Anions Elemental anions will end in -ide Example –Cl 2 - chlorine - neutral element –Cl 1- - chloride - anion of the element Another example –O 2 - oxygenO 2- Oxide

73. Writing Formulas for Binary Ionic Compounds The periodic table tells you the charge for group A (aka - the representative elements) Group 1A - 1+ Group 2A - 2+ Group 3A - 3+ Group 4 - depends Group 5A - 3- Group 6A - 2- Group 7A - 1- Group 8A or (0) - does not form ions

74. Naming Your turn: –Name or write the symbol for the following: AluminumPhosphide Calcium IonIodine Ga 3+ Nitrogen KSulfide

75. Naming Binary Ionic Compounds Name the metal then the nonmetal with the ending changing to -ide –The -ide tells the person it is a binary compound and the anion portion. Examples:MgCl 2 K 2 S Magnesium Chloride Potassium Sulfide

76. Writing Formulas for Binary Ionic Compounds All compounds are electrically neutral To write the formula, figure out how many cations and anions are needed so that the number of positives and negatives are equal. Find the least common multiple to figure out the total number of +’s and -’s. Then divide by the charge to find out how many of each atom is needed! If X 1+ and Y 2-, what would be the formula? X 2 Y - Charges total 2 +’s and 2 -’s

77. Writing Formulas for Binary Ionic Compounds If X 3+ and Y 2-, what would be the formula? X 2 Y 3 - Charges total 6 +’s and 6 -’s Find the formula for the following pairs of ions: –Na 1+, P 3- Sr 2+, N 3- Answers:

78. Now: –Finish side 1 of worksheet –Work sections 1 - 4 on back of worksheet –Work homework problems

79. Writing formulas for multivalent ionic compounds Transition metals have the ability to form more than one cation Therefore, a roman numeral is placed in the name to signify the charge on the cation Example: –Iron (III) Chloride Write the formula?

82. Writing formulas for mulitvalent ionic compounds Write formulas for the following: Copper (I) Oxide Copper (II) Oxide Answers -

83. Naming compounds with multivalent metals If the metal is in group B it requires a roman numeral in the name. You will have to deduce the roman numeral based on the formula. Example –Name CoI 2 Answer -

84. Naming compounds with multivalent metals Deducing the roman numeral Multiply the charge on the anion by the number of anions and then divide by the number of cations to get the roman numeral. Write the names for Fe 2 S 3 SnO 2 Answers -

85. Take ten minutes and work on sections 5 and 6 on the back side of your worksheet.

86. Polyatomic Ions A group of atoms that carry a charge Examples: –SO 4 2- NO 3 1- Names of polyatomic ions that contain oxygen will end in -ate or -ite -ite is one less oxygen then ate Example –Sulfate is SO 4 2- Sulfite is SO 3 2- –Chlorate is ClO 3 1- Chlorite is ClO 2 1- Other polyatomic ions –NH 4 1+ Ammonium CN 1- cyanide –OH 1- Hydroxide

87. Writing formulas using polyatomic ions The polyatomic ion is treated as one unit. Balance the charges Place parenthesis around the polyatomic ion if there is more than one Example –Write the formula for Iron (II) Nitrate

88. Naming using Polyatomic ions Name the metal then name the polyatomic ion. If you need a roman numeral; include it. Treat the polyatomic ion as one unit (as if it were one atom) Example - Name CuSO 4

89. Exceptions for roman numerals Silver, Cadmium and Zinc do not get roman numerals. Ag is always +1, Cadmium and Zinc are always +2 Tin and Lead need roman numerals. They are multivalent (multiple oxidation states)

90. Naming Acids Memorize HCl - Hydrochloric Acid H 2 SO 4 - Sulfuric Acid HNO 3 - Nitric Acid H 3 PO 4 - Phosphoric Acid Note - Acids give off H 1+ (Hydrogen ions) and bases give off OH 1- ions What do you get when an acid and base combine?

92. Check for understanding Name or write the formula for: –Potassium Sulfate –Chromium (III) Cyanide –Fe(ClO 3 ) 3 –CuCl Answers

93. Now finish your worksheet and work on your homework. Get help Make sure and check your answers. You will be writing formulas all year and doing math based on these formulas. You get the formula wrong you get the math wrong.

95. Helpful hints for balancing chemical equations Balance hydrogens second to last Balance oxygens last Check for lowest ratio Coefficients must be whole numbers Don’t break up your compounds with coefficients –NaCl cannot become Na6Cl Do not change your subscripts Balance the polyatomic ions as one unit (if it didn’t break apart) Perform a final check

96. Balance the following C 2 H 6 + O 2 --> CO 2 + H 2 O Na 3 PO 4 + Mg(NO 3 ) 2 --> NaNO 3 + Mg 3 (PO 4 ) 2

97. Types of Reactions Including reaction prediction

98. Generals about writing Equations Reactants on the left and products on the right Symbols - see text for symbols that are included in equations. –Ex: g for gas, l for liquid, s for solid –Downward arrow for precipitate, aq for aqueous Catalyst goes above the arrow KI –Ex H 2 O 2(aq) ---> H 2 O (l) + O 2(g) Diatomic Molecules - BrINClHOF –Elemental state - Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2

99. 1. Synthesis (Combination) Two or more substances react to form a single substance R + S --> RS Ex) SO 3(g) + H 2 O (l) --> H 2 SO 4(aq) Usually gives off energy when forming bonds Example: Write the balanced equation for: magnesium ribbon reacting with oxygen Mg (s) + O 2(g) ---> MgO (s) 2 Mg (s) + O 2(g) --> 2MgO (s)

100. 1. Synthesis (Combination) Your turn. Write balanced equations for the following: –Aluminum (s) reacts with oxygen (g) –Hydrogen (g) reacts with oxygen (g) Answers:

101. 2. Decomposition A single compound is broken down into simpler products RS --> R + S Ex) BCl 3 --> B + Cl 2 Requires energy to break chemical bonds (heat, light, electricity) Example - Write the balanced equation for mercury (II) oxide decomposing; HgO --> Hg + O 2 2HgO --> 2Hg + O 2

102. 2. Decomposition Your turn. Write balanced equations for the following: The decomposition of water The decomposition of lead (IV) oxide Answers

103. 3. Single Replacement Reactions An element replaces an element of a compound T + RS --> TS + R Ex) Zn (s) + H 2 SO 4(aq) --> ZnSO 4(aq) + H 2(g) A metal may replace a metal or a nonmetal may replace a nonmetal Activity Series - list of metal in order of decreasing activity Nonmetals reactivity decreases as you go down the periodic table This is limited to the halogens -group 7A

104. 3. Single replacement reactions Ex) Write the balanced equation when aluminum reacts with sulfuric acid Al (s) + H 2 SO 4(aq) --> Al 2 (SO 4 ) 3(s) + H 2(g) 2Al (s) + 3H 2 SO 4(aq) --> Al 2 (SO 4 ) 3(s) + 3H 2(g)

105. 3. Single replacement reactions Your turn. Write balanced equations for the following: When chlorine reacts with potassium iodide When copper (assume Cu 2+ ) is added to Iron (II) Sulfate Answers –

106. 4. Double Replacement Exchange of positive ions between two compounds. Just swap the positive ions and write the new formula. R + S - + T + U - --> R + U - + T + S - Ex) FeS (s) + 2HCl (aq) --> H 2 S (g) + FeCl 2(aq) Ex) Write the balanced equation for barium chloride added to potassium carbonate BaCl 2(aq) + K 2 CO 3(aq) --> BaCO 3(s) + KCl (aq) BaCl 2(aq) + K 2 CO 3(aq) --> BaCO 3(s) + 2 KCl (aq)

107. 4. Double Replacement Your turn. Write balanced equations for the following. Iron (III) Sulfide reacting with hydrochloric acid Answer

108. 5. Combustion Reactions Oxygen reacts with another substance, often producing heat and light Often involve hydrocarbons –Compounds of hydrogen and carbon Combustion of hydrocarbons produces a lot of energy, therefore, hydrocarbons are used as fuels. Examples: methane, propane, butane, octane

109. 5. Combustion Reactions Two types of combustion 1. Complete combustion C x H y + O 2(g) --> CO 2(g) + H 2 O (g) + energy 2. Incomplete combustion Two more products: CO and C C x H y + O 2(g) --> CO 2(g) + H 2 O (g) + CO (g) + C (s) + energy

110. 5. Combustion Reactions Ex) Write a balanced equation for the complete combustion of C 3 H 8. C 3 H 8(g) + O 2(g) --> CO 2(g) + H 2 O (g) + energy C 3 H 8(g) + 5 O 2(g) --> 3 CO 2(g) + 4H 2 O (g) + energy Your turn: Write a balanced equation for the complete combustion of C 8 H 18. Answer

113. Precipitation Reactions Most ionic compounds dissociate into cations and anions when dissolved in water. A complete ionic equation (basically a double replacement reaction) shows ionic compounds as free ions. In other words, write in the charges.

114. Precipitation reactions Predicting the precipitate Use the chart on the back of your periodic table. Which of the following compounds are not soluble –Calcium Sulfate –Sodium Acetate –Silver Chloride –Aluminum Hydroxide –Potassium Phosphate

115. Precipitation Reactions Complete Ionic Equations In an aqueous solution, substances exist as free ions. The equation shows this. Example for AgNO 3(aq) + NaCl (aq) Ag +1 (aq) + NO 3 1- (aq) + Na 1+ (aq) + Cl 1- (aq) --> AgCl (s) + Na +1 (aq) + NO 3 1- (aq)

116. Precipitation Reactions Net Ionic Equation A net ionic equation indicates those ions that took part in the reaction. Net ionic equation for the reaction from the previous slide is: Ag 1+ (aq) + Cl 1- (aq) --> AgCl (s)

117. Precipitation Reaction Example: Write a complete and net ionic equation for the reaction of aqueous solutions of iron (III) nitrate and sodium hydroxide. Fe 3+ (aq) + NO 3 1- (aq) + Na +1 (aq) + OH - (aq) --> Fe(OH) 3(s) + Na +1 (aq) + NO 3 1- (aq) Fe 3+ (aq) + OH 1- (aq) --> Fe(OH) 3(s)

118. Precipitation Reactions Your turn. Write a complete ionic equation and a net ionic equation for the reaction of aqueous solutions of silver nitrate and potassium sulfate. Answer

119. The Mole I. Molar Conversions

120. 1 mole of hockey pucks would equal the mass of the moon! 1 mole of pennies would cover the Earth 1/4 mile deep! 1 mole of basketballs would fill a bag the size of the earth!

121. Molar Conversions Converting from moles to grams to representative particles and vice versa. Use the following conversion factor: 1 mole = 6.02 x 10 23 representative units = molar mass (g) or formula weight Representative units - a. ionic compounds are called formula units b. molecular compounds are called molecules c. atoms are called atoms. Example of representative units - 6.02 x 10 23 atoms Cu 6.02 x 10 23 molecules O 2 6.02 x 10 23 units NaCl

122. Molar Conversion Examples How many moles of carbon are in 26.0 g of carbon? 26.0 g C 1 mol C 12.0 g C = 2.17 mol C

123. Molar Conversion Examples How many molecules are in 2.50 moles of C 12 H 22 O 11 ? 2.50 mol C 12 H 22 O 11 6.02  10 23 molecules 1 mol C 12 H 22 O 11 = 1.51  10 24 molecules C 12 H 22 O 11

124. Molar Conversion Examples Find the mass of 2.1  10 24 formula units of NaHCO 3. 2.1  10 24 units NaHCO 3 84.0 g NaHCO 3 6.02  10 23 units NaHCO 3 = 290 g NaHCO 3

125. Molar Conversion Examples Find the number of units of Iron (III) Chlorate in 98.6 g of Iron (III) Chlorate. 98.6 g Fe(ClO 3 ) 3 6.02 x 10 23 units Fe(ClO 3 ) 3 306.3 g Fe(ClO 3 ) 3 = 1.94 x 10 23 units Fe(ClO 3 ) 3

126. Moles in a Gas 1 mole of gas takes up 22.4 L of space at standard temperature and pressure. Conversion factor - 1 mole = 22.4 L –Remember this is for a gas only Standard Temperature and Pressure (STP) –Temp = 0 o C –Pressure = 1 atm (atmosphere) –1 atmosphere is defined as the amount of pressure the earth’s atmosphere places on you at sea level

127. Calculations w/ molar volume Determine the volume, in liters, of 0.60 mol SO 2 gas at STP. Answer - –0.60 mol SO 2 22.4 L SO 2 1 mol SO 2 = 13 L SO 2

128. Calculations w/ molar volume Your Turn How many atoms of He are contained in your party balloon if the balloon takes up 4.2 L of space? Of course, this is one cold party, as it would be held at STP. Answer -

129. Molarity Unit of Concentration –There are many units of concentration Molarity is most useful to the chemist M = moles of solute Liters of solution Liters of solution means the total volume of water and solute. If I want a liter of solution I will not use a liter of water.

130. Molarity Problems You work them. A saline solution contains 0.90 g NaCl in exactly 100 ml of solution. What is the molarity of the solution?

131. Molarity Problems You work them. How many moles of solute are present 1.5 L of 0.24 M Na 2 SO 4 ?

132. Preparing a solution How would you make 500.0 ml of a 0.25 M solution of copper (II) chloride? 0.25 M = mol/0.5000 L - change ml to liters and solve for moles. You need 0.13 moles of CuCl 2. Converting to grams equals 17 grams. Final answer –Take 17 grams of CuCl 2 and dissolve in enough water to make 500.0 ml of solution. Dissolve the 17 grams in say 400 ml of water. Once the CuCl 2 is dissolved add water up to 500.0 ml.

133. Preparing a solution your turn How would you prepare a solution of 0.40 M KCl? If a volume is not given assume 1 L.

134. Making Dilutions Making dilutions from known concentrations: M 1 x V 1 = M 2 x V 2 Volume can be in liters or mL as long as the same units are used.

135. Dilution Problems How would you prepare 1.00x10 2 mL of 0.40 M MgSO 4 from a stock solution of 2.0 M MgSO 4 ? 0.40 M x 100 mL= 2.0 M x V 2 V 2 = 20 mL Answer - Take 20 mL of 2.0 M MgSO 4 and dilute with enough water to make 100 mL of solution.

136. Dilution Problems Your Turn How would you prepare 90.0 mL of 2.0 M H 2 SO 4 from 18 M stock solution? Answer

137. Dilution Problems Your Turn - 1 more If 250 mL of a 12.0 M HNO 3 is diluted to 1 L, what is the molarity of the final solution? Answer -

138. Percent Composition

140. % composition your turn Hydroxide makes up what percent of Calcium Hydroxide? Answer

141. Hydrates Hydrates are substances that contain water within the crystalline structure of the compound. The water is not chemically bound; it is trapped within the crystal. Ex. FeSO 4. 7H 2 O

142. Empirical vs. Molecular Formula

143. Calculating Empirical Formulas –Lowest whole-number ratio of the atoms of the elements in a compound C 6 H 12 O 6 (glucose) The ratio that glucose normally has for carbon:hydrogen:oxygen is 6:12:6. The lowest ratio that glucose has for carbon:hydrogen:oxygen is 1:2:1 (each number can be divided by the smallest number in the ratio which is 6).

144. The empirical formula for glucose is CH 2 O since this is the lowest whole- number ratio of atoms for that compound. –May or may not be the same as the normal molecular formula of a compound Next - Calculating empirical formulas

145. What is the empirical formula of a compound that is 25.9% nitrogen and 74.1% oxygen? If 25.9% of the compound is nitrogen and 74.1% of the compound is oxygen, then a compound with a mass of 100 g has 25.9 g of nitrogen and 74.1 g of oxygen. To calculate the empirical formula, we need to relate the moles of each atom in the compound, so we need to convert the masses of the elements to moles.

146. This would mean that the ratio of nitrogen to oxygen is N 1.85 O 4.63. We can divide each number in the ratio of N 1.85 O 4.63 by 1.85 to get N 1 O 2.50. Since we cannot have 2.50 atoms of oxygen, we must multiply through each number by 2 to even it out, getting N 2 O 5 as our empirical formula.

147. In calculating empirical formulas, remember that the number of atoms is a whole number. If the number of atoms for an element is close to a whole number (i.e., 2.1 or 2.2 or 2.8, or 2.9), you can usually round up or down to get a whole number. If you should get a number of atoms closer to 2.33 or 2.5, multiply each number in the formula by a number that gets that to a whole number. For example, if you calculated 2.33, you would multiply this by 3 to get a value of 7 for that number.

148. Give it a try Determine the empirical formula for a compound containing 7.8% carbon and 92.2% chlorine.

149. Empirical vs. Molecular Formula

150. Calculating Molecular Formulas Although sometimes a molecular formula may be the same as a molecule’s empirical formula, like in carbon dioxide (CO 2 ), we have seen that the empirical formula for glucose is not the same as its molecular formula. One can determine the molecular formula of a compound by knowing its empirical formula and its mass. Next - Example

151. Calculate the molecular formula of the compound whose molar mass is 180.1583 g and empirical formula is CH 2 O. Now, in order to figure out what we must multiply each number in the empirical formula by, we must figure out by what number we must multiply the empirical formula mass to get the molecular formula mass. To get from 30.0264 to 180.1583, We know that the molecular formula will have a molar mass of 180.1583 g. We also know, by calculating the gmm of CH 2 O, that CH 2 O has an empirical formula mass (efm) = 30.0264 g CH 2 O.

152. Therefore, we must multiply each number of atoms in CH 2 O by 6 to get the molecular formula of C 6 H 12 O 6. You can double-check your answer by recalculating the molar mass of C 6 H 12 O 6. gmm C 6 H 12 O 6 = 6 x 12.0111 g + 12 x 1.00794 g + 6 x 15.9994 g = 180.1583 g C 6 H 12 O 6 This agrees with the molar mass we were given, so the molecular formula we calculated is correct.

153. Give it a try Determine the molecular formula of a compound that is 40.0% C, 6.6% H, and 53.4% O and the molar mass is 120.0g.

154. Example (toughy) 1.00 g of menthol on combustion yields 1.161 g of H 2 O and 2.818 g of CO 2. What is the empirical formula? Solution:

155. Stoichiometry Calculations of quantities in chemical reactions. The use of ratios to calculate quantities

156. The five step process 1. Start with the balanced equation 2. Set up the problem - put down the tracks 3. Convert to moles if needed. This means you would be given grams, representative units or liters. 4. Convert to moles of what you want. You will use the mole ratio from the balanced equation. 5. Convert to what you are trying to find (grams, liters, representative units) if needed.

157. Stoichiometry Example Problem #1 How many moles of ammonia are produced when 0.60 mol of hydrogen reacts with nitrogen?

158. Stoichiometry Example Problem #2 Your Turn How many moles of aluminum sulfide are produced when 1.2 moles of aluminum reacts with sulfur?

159. Stoichiometry Example Problem #2 Answer

160. Stoichiometry Example Problem #3 How many grams of ammonia will be produced by reacting 5.40 g of hydrogen with nitrogen?

161. Stoichiometry Example Problem #4 Your Turn How many grams of aluminum are needed to react with 2.45 g of copper (II) chloride?

162. Stoichiometry Example Problem #4 Answer

163. % Yield Definitions 1. Theoretical Yield The maximum amount of product that can be formed from a given amount of reactants In other words, the calculated amount predicted through stoichiometry

164. % Yield Definitions Actual Yield The amount that is actually formed when the reaction is carried out in the laboratory.

165. % Yield = actual yield X 100 theoretical yield % yield will never be over 100% Most likely it will never even be 100%

166. Why will % yield never be 100% Advantageous to add an excess of an inexpensive reagent to ensure that all of the more expensive reagents reacts Reactant may not be 100% pure Materials are lost during the reaction –If a reactions takes place in a solution it may be impossible to get all of the reactants or products out of the solution If the reactions takes place at a high temperature, materials may be vaporized and escape into the air Side reactions may occur –Example Mg burned in air. Some of Mg reacts with nitrogen reducing the amount of MgO produced. Loss of product when filtering or transferring If reactants are not carefully measured

167. % yield example problem In a reaction between barium chloride and potassium sulfate, 3.89 g of barium sulfate is produced from 3.75 g of barium chloride. What is the percent yield?

168. % yield example problem answer BaCl 2 + K 2 SO 4 --> BaSO 4 + 2 KCl 3. 75 g BaCl 2 1 mol BaCl 2 1 mol BaSO 4 233.4 g BaSO 4 208.3 g BaCl 2 1 mol BaCl 2 1 mol BaSO 4 = 4.20 g BaSO 4 3.89 g BaSO 4 x 100 = 92.6 % 4.20 g BaSO 4

169. % yield example problem your turn 13.35 grams of magnesium hydroxide is produced when 42.50 grams of magnesium nitrate reacts with an excess of aluminum hydroxide. What is the percent yield?

170. % yield example problem answer

171. Limiting Reagent 1. Limits or determines the amount of product that can be formed 2. The reagent that is not used up is therefore the excess reagent These types of problems require 2 sets of tracks. Quantities of both reagents will be given. Therefore, you need to find out which one is the limiting reagent.

172. Limiting Reagent One track to determine limiting reagent A second track to determine product

173. Limiting Reagent Example problem How many grams of copper (I) Sulfide can be produced when 80.0 grams of Cu reacts with 25.0 grams of sulfur? 2Cu + S --> Cu 2 S Pick a reactant and calculate how much of the other reactant is needed. 80.0g Cu 1mol Cu 1mol S 32.1g S 63.5g Cu 2mol Cu 1mol S = 20.2g S So, 20.2 g of S is needed; 25.0g is supplied Plenty of S; therefore, Cu is limiting reagent. Use Cu to solve the problem 80.0g Cu 1mol Cu 1mol Cu 2 S 159.1g Cu 2 S 63.5g Cu 2mol Cu 1mol Cu 2 S = 1.00x10 2 g Cu 2 S

174. Limiting Reagent Example Problem - Your Turn How many grams of hydrogen can be produced when 5.00g of Mg is added to 6.00 g of HCl?

175. Limiting Reagent Example problem- Your Turn Acetylene (C 2 H 2 ) will burn in the presence of oxygen. How many grams of water can be produced by the reaction of 2.40 mol of acetylene with 7.4 mol of oxygen?