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“Good morning, and welcome to introduction to chemistry.”

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

2 Info Class: Chemistry Instructor: Mr. Cooper
Office: A112 - I’m pretty much in my classroom before and after school WebPage: click on teachers and find my name.

3 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.

4 Class Format 2. Lecture: I will provide you one set of Lecture notes. They are posted on the web. You can print them off if you lose your set. This allows you to listen and formulate questions and take your own notes rather than just copying down my outline. If all of the class has notes ahead of time, we can spend more time on answering questions, lab work, book work and individual help rather than copying notes.

5 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. The labs are on-line if you lose the one provided lab book. Labs may be replaced by worksheets, group work, a video, or demonstrations

6 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 an excellent time for you to get individual help.

7 Grading Lab Quizzes: The day before each unit test will be a unit lab quiz. It is open lab book.

8 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.

9 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.)

10 Grading There will be no test retakes!!!!
However, you will have a take home practice test. So, your unit test is actually your retake.

11 Grading The grading scale is as follows: A= 90-100 B+= 85-89
B= C+=75-79 C= D+= 65-69 D= F= Below 60.0

12 Misc. Any assignments or test missed for truancy results in 60.0% of earned grade. This is district policy. Late work policy – Once I have your assignment graded and handed back, I will no longer accept that assignment.

13 Misc Tardies - Building policy is followed.
Cell Phones – Building policy is followed. 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.

14 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.

15 All work will be shown at all times or no credit will be
given. I should never have to make this announcement again or even write in on a worksheet or test.

16 Do not give me late work without a roll of the dice.

17 Summary of Grades Lab Quiz will be day before test Due Day of test
Self-eval sheet One random homework pickup Lab Report Multiple Choice Test Short Answer Test

18 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) Air Force Core Values Integrity First Excellence in All Service Before Self 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

19 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!!!

20 Tirrill (Bunsen) Burner

21 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.

22 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.

23 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

24 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.

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

26 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.

27 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 other than your text. Do not use any “I” statements.

28 Sample lab report Title: Place title here.
Purpose: Place the purpose here. Procedure: A couple of general sentences summarizing lab steps. Data: 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.

29 Sample lab report Conclusion: 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: Keep answers clear and concise. Length is not important. I care about content and good communication.

30 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.

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

32 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

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

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

35 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 oC 1 cm3 = 1 ml

36 Density Problem A substance has a volume of 1.74 ml and a mass of 20.0 grams. What is the density?

37 Density Problem One more to test your algebra:
What is the volume of ice in a container if the density is g/ml and the mass is g?

38 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 - C12H22O11 Lemonade is not a pure substance

39 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

40 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

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

42 States of Matter Video

43 Properties of Matter Physical Property
An observed condition of the substance Physical properties help identify substances Examples include:

44 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 H2O Vapor - a substance that is in a gaseous state but liquid at room temp

45 Change of State - a physical change

46 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

47 Physical and Chemical Properties

48 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)

49 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.

50 A chemical change

51 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

52 Categorizing our environment

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

54 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

55 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


57 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

58 Categorizing our environment

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

60 Classifying Matter Compounds
Two or more elements combined through a chemical bond Can only be separated into simpler substances by chemical reactions Example Sugar - C12H22O11

61 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

62 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

63 Categorizing our environment

64 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

65 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




69 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? - Cl2O8 and OF2 Formula for? - dinitrogen tetroxide Answers -

70 Naming Molecular Compounds
Your turn. Try these. Name or write the formula for: Boron trichloride Dinitrogen tetrahydride N2O PF S4N CCl4 SO H2O Answers

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

72 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

73 Naming Cations Name the metal followed by the word ion Example
Na - sodium - neutral element Na1+- sodium ion - cation of the element Another example: Mg - magnesium Mg2+ - Magnesium ion

74 Naming Anions Ending changes are used for Anions
Elemental anions will end in -ide Example Cl2 - chlorine - neutral element Cl1- - chloride - anion of the element Another example O2 - oxygen O2- Oxide

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

76 Naming Your turn: Aluminum Phosphide Calcium Ion Iodine Ga3+ Nitrogen
Name or write the symbol for the following: Aluminum Phosphide Calcium Ion Iodine Ga3+ Nitrogen K Sulfide

77 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: MgCl2 K2S Magnesium Chloride Potassium Sulfide

78 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 X1+ and Y2-, what would be the formula? X2Y - Charges total 2 +’s and 2 -’s

79 Writing Formulas for Binary Ionic Compounds
If X3+ and Y2-, what would be the formula? X2Y3 - Charges total 6 +’s and 6 -’s Find the formula for the following pairs of ions: Na1+ , P3- Sr2+ , N3- Answers:

80 Now: Finish side 1 of worksheet
Work sections on back of worksheet Work homework problems

81 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?



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

85 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 CoI2 Answer -

86 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 Fe2S3 SnO2 Answers -

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

88 Polyatomic Ions A group of atoms that carry a charge Examples:
SO NO31- Names of polyatomic ions that contain oxygen will end in -ate or -ite -ite is one less oxygen then ate Example Sulfate is SO42- Sulfite is SO32- Chlorate is ClO31- Chlorite is ClO21- Other polyatomic ions NH41+ Ammonium CN1- cyanide OH1- Hydroxide

89 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

90 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 CuSO4

91 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)

92 Naming Acids Memorize HCl - Hydrochloric Acid H2SO4 - Sulfuric Acid
HNO3- Nitric Acid H3PO4 - Phosphoric Acid Note - Acids give off H1+ (Hydrogen ions) and bases give off OH1- ions What do you get when an acid and base combine?


94 Check for understanding
Name or write the formula for: Potassium Sulfate Chromium (III) Cyanide Fe(ClO3)3 CuCl Answers

95 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.


97 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

98 Balance the following Na3PO4 + Mg(NO3)2 --> NaNO3 + Mg3(PO4)2
C2H O2 --> CO H2O Na3PO Mg(NO3)2 --> NaNO Mg3(PO4)2

99 Including reaction prediction
Types of Reactions Including reaction prediction

100 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 aq for aqueous Catalyst goes above the arrow KI Ex H2O2(aq) ---> H2O(l) + O2(g) Diatomic Molecules - BrINClHOF Elemental state - Br2I2N2Cl2H2O2F2

101 1. Synthesis (Combination)
Two or more substances react to form a single substance R + S --> RS Ex) SO3(g) + H2O(l) --> H2SO4(aq) Usually gives off energy when forming bonds Example: Write the balanced equation for: magnesium ribbon reacting with oxygen Mg(s) O2(g) ---> MgO(s) 2 Mg(s) + O2(g) --> 2MgO(s)

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

103 2. Decomposition A single compound is broken down into simpler products RS --> R + S Ex) BCl3 --> B + Cl2 Requires energy to break chemical bonds (heat, light, electricity) Example - Write the balanced equation for mercury (II) oxide decomposing; HgO --> Hg + O2 2HgO --> 2Hg + O2

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

105 3. Single Replacement Reactions
An element replaces an element of a compound T + RS --> TS + R Ex) Zn(s) + H2SO4(aq) --> ZnSO4(aq) + H2(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

106 3. Single replacement reactions
Ex) Write the balanced equation when aluminum reacts with sulfuric acid Al(s) + H2SO4(aq) --> Al2(SO4)3(s) + H2(g) 2Al(s)+ 3H2SO4(aq) --> Al2(SO4)3(s) + 3H2(g)

107 3. Single replacement reactions
Your turn. Write balanced equations for the following: When chlorine reacts with potassium iodide When copper (assume Cu2+) is added to Iron (II) Sulfate Answers

108 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) --> H2S(g) + FeCl2(aq) Ex) Write the balanced equation for barium chloride added to potassium carbonate BaCl2(aq) + K2CO3(aq) --> BaCO3(s) + KCl(aq) BaCl2(aq) + K2CO3(aq) --> BaCO3(s) + 2 KCl(aq)

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

110 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

111 5. Combustion Reactions Two types of combustion 1. Complete combustion
CxHy + O2(g) --> CO2(g) + H2O(g) + energy 2. Incomplete combustion Two more products: CO and C CxHy + O2(g) --> CO2(g) + H2O(g) + CO(g) + C(s) + energy

112 5. Combustion Reactions Ex) Write a balanced equation for the complete combustion of C3H8. C3H8(g) + O2(g) --> CO2(g) + H2O(g) + energy C3H8(g) + 5 O2(g) --> 3 CO2(g) + 4H2O(g) + energy Your turn: Write a balanced equation for the complete combustion of C8H18. Answer



115 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.

116 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

117 Precipitation Reactions Complete Ionic Equations
In an aqueous solution, substances exist as free ions. The equation shows this. Example for AgNO3(aq) + NaCl(aq) Ag+1(aq) + NO31-(aq) + Na1+(aq) + Cl1-(aq) --> AgCl(s) + Na+1(aq) + NO31-(aq)

118 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: Ag1+(aq) + Cl1-(aq) --> AgCl(s)

119 Precipitation Reaction
Example: Write a complete and net ionic equation for the reaction of aqueous solutions of iron (III) nitrate and sodium hydroxide. Fe3+(aq) + NO31-(aq) + Na+1(aq) + OH-(aq) --> Fe(OH)3(s) + Na+1(aq) + NO31-(aq) Fe3+(aq) + OH1-(aq) --> Fe(OH)3(s)

120 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

121 Molar Conversions The Mole

122 HOW LARGE IS IT??? 1 mole of hockey pucks would equal the mass of the moon! 1 mole of basketballs would fill a bag the size of the earth! 1 mole of pennies would cover the Earth 1/4 mile deep!

123 Molar Conversions Converting from moles to grams to representative particles and vice versa. Use the following conversion factor: 1 mole = 6.02 x 1023 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 atoms Cu 6.02 x molecules O2 6.02 x units NaCl

124 Molar Conversion Examples
How many moles of carbon are in 26.0 g of carbon?

125 Molar Conversion Examples
How many molecules are in 2.50 moles of C12H22O11?

126 Molar Conversion Examples
Find the mass of 2.1  1024 formula units of NaHCO3.

127 Molar Conversion Examples
Find the number of units of Iron (III) Chlorate in 98.6 g of Iron (III) Chlorate.

128 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 = 0oC Pressure = 1 atm (atmosphere) 1 atmosphere is defined as the amount of pressure the earth’s atmosphere places on you at sea level

129 Calculations w/ molar volume
Determine the volume, in liters, of 0.60 mol SO2 gas at STP. Answer - 0.60 mol SO L SO2 1 mol SO2 = 13 L SO2

130 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 -

131 M = moles of solute Molarity Liters of solution 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.

132 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?

133 Molarity Problems You work them.
How many moles of solute are present 1.5 L of 0.24 M Na2SO4?

134 Preparing a solution How would you make ml of a 0.25 M solution of copper (II) chloride? 0.25 M = mol/ L - change ml to liters and solve for moles. You need 0.13 moles of CuCl2. Converting to grams equals 17 grams. Final answer Take 17 grams of CuCl2 and dissolve in enough water to make ml of solution. Dissolve the 17 grams in say 400 ml of water. Once the CuCl2 is dissolved add water up to ml.

135 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.

136 Making Dilutions M1 x V1 = M2 x V2
Making dilutions from known concentrations: M1 x V1 = M2 x V2 Volume can be in liters or mL as long as the same units are used.

137 Dilution Problems How would you prepare 1.00x102 mL of 0.40 M MgSO4 from a stock solution of 2.0 M MgSO4? 0.40 M x 100 mL= 2.0 M x V2 V2 = 20 mL Answer - Take 20 mL of 2.0 M MgSO4 and dilute with enough water to make 100 mL of solution.

138 Dilution Problems Your Turn
How would you prepare 90.0 mL of 2.0 M H2SO4 from 18 M stock solution? Answer

139 Dilution Problems Your Turn - 1 more
If 250 mL of a 12.0 M HNO3 is diluted to 1 L, what is the molarity of the final solution? Answer -

140 Percent Composition


142 % composition your turn
Hydroxide makes up what percent of Calcium Hydroxide? Answer

143 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. FeSO4 . 7H2O

144 Empirical vs. Molecular Formula

145 Calculating Empirical Formulas
Lowest whole-number ratio of the atoms of the elements in a compound C6H12O6 (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).

146 Next - Calculating empirical formulas
The empirical formula for glucose is CH2O 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

147 What is the empirical formula of a compound that is 25
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.

148 This would mean that the ratio of nitrogen to oxygen
is N1.85O4.63. We can divide each number in the ratio of N1.85O4.63 by 1.85 to get N1O2.50. Since we cannot have 2.50 atoms of oxygen, we must multiply through each number by 2 to even it out, getting N2O5 as our empirical formula.

149 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.

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

151 Empirical vs. Molecular Formula

152 Calculating Molecular Formulas
Although sometimes a molecular formula may be the same as a molecule’s empirical formula, like in carbon dioxide (CO2), 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

153 Calculate the molecular formula of the compound whose molar mass is g and empirical formula is CH2O. We know that the molecular formula will have a molar mass of g. We also know, by calculating the gmm of CH2O, that CH2O has an empirical formula mass (efm) = g CH2O. 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 to ,

154 Therefore, we must multiply each number of atoms in CH2O by 6 to get the molecular formula of C6H12O6. You can double-check your answer by recalculating the molar mass of C6H12O6. gmm C6H12O6 = 6 x g + 12 x g + 6 x g = g C6H12O6 This agrees with the molar mass we were given, so the molecular formula we calculated is correct.

155 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.

156 Example (toughy) 1.00 g of menthol on combustion yields g of H2O and g of CO2. What is the empirical formula? Solution:

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

158 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.

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

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

161 Stoichiometry Example Problem #2

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

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

164 Stoichiometry Example Problem #4

165 % 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

166 % Yield Definitions Actual Yield
The amount that is actually formed when the reaction is carried out in the laboratory.

167 % Yield = actual yield X 100 theoretical yield
% yield will never be over 100% Most likely it will never even be 100%

168 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

169 % 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?

170 % yield example problem answer
BaCl2 + K2SO4 --> BaSO KCl 3. 75 g BaCl2 1 mol BaCl mol BaSO g BaSO4 208.3 g BaCl2 1 mol BaCl mol BaSO4 = 4.20 g BaSO4 3.89 g BaSO4 x = % 4.20 g BaSO4

171 % yield example problem your turn
13.35 grams of magnesium hydroxide is produced when grams of magnesium nitrate reacts with an excess of aluminum hydroxide. What is the percent yield?

172 % yield example problem answer

173 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.

174 Limiting Reagent One track to determine limiting reagent
A second track to determine product

175 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 --> Cu2S Pick a reactant and calculate how much of the other reactant is needed. 80.0g Cu 1mol Cu 1mol S g 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 Cu2S g Cu2S 63.5g Cu 2mol Cu mol Cu2S = 1.00x102 g Cu2S

176 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?

177 Limiting Reagent Example problem- Your Turn
Acetylene (C2H2) 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?

178 The Development of Atomic Models
Democritus was a pre-Socratic Greek philosopher (born around 460 BC). Democritus was originator of the belief that all matter is made up of various imperishable, indivisible elements which he called "atomos", from which we get the English word atom. According to legend, Democritus was supposed to be mad because he laughed at everything, and so he was sent to Hippocrates to be cured. Hippocrates pointed out that he was not mad, but, instead, had a happy disposition. That is why Democritus is sometimes called the laughing philosopher. BB - Model Dalton’s Model More to come

179 Plum Pudding Model Thomson's model was compared (though not by Thomson) to a British treat called plum pudding, hence the name. It has also been called the chocolate chip cookie model, but only by those who have not read Thomson's original paper Proposed by J. J. Thomson ( ), the discoverer of the electron in 1897. The plum pudding model was proposed in March, 1904 before the discovery of the atomic nucleus. In this model, the atom is composed of electrons surrounded by a soup of positive charge to balance the electron's negative charge, like plums surrounded by pudding. The electrons were thought to be positioned throughout the atom. Electrons could move like letters in alphabet soup Instead of a soup, the atom was also sometimes said to have had a cloud of positive charge.

180 Nuclear Model The Gold foil experiment or the Rutherford experiment was an experiment done by Ernest Rutherford ( ) in This experiment discovered the nucleus. Led to the downfall of the plum pudding model of the atom. Alpha particles (positive particles--Helium Nuclei) were shot at gold foil. Particles passed through the gold foil. A few shot back. Conclusions: 1. Atom is mostly empty space 2. Dense center called the nucleus 3. Electrons were stuck surrounding the nucleus.


182 Planetary Model Introduced by Niels Bohr, a Danish physicist ( ), in 1913. Because of its simplicity, the Bohr model is still commonly taught to introduce students to quantum mechanics. The Bohr model depicts the atom as a small, positively charged nucleus surrounded by waves of electrons in orbit — similar in structure to the solar system, but with electrostatic forces providing attraction, rather than gravity. "The opposite of a correct statement is a false statement. But the opposite of a profound truth may well be another profound truth." Niels Bohr

183 Quantum Mechanical Model
Erwin Schrödinger (August 12, 1887 – January 4, 1961) An Austrian physicist, achieved fame for his contributions to quantum mechanics, especially the Schrödinger equation, for which he received the Nobel Prize in 1933. This model is based on probability Where are you going to find and electron 90% of the time. Atom is viewed as a fuzzy cloud. Schrödinger equations create electron clouds (orbitals) with specific shapes.

184 Main Points For “Atoms” Video
What is the key to understanding atomic structure? The discovery of what particle is associated with the Crook’s Tube? What did Rutherford expect to happen in the gold foil experiment? What was Rutherford’s genius? What conclusions did Rutherford draw from the Gold foil experiment? How much smaller is the nucleus than the electron cloud? What determines the shape of the electron cloud?

185 Dalton’s Atomic Theory
Democritus - Greek philosopher who first suggested atoms John Dalton ( ) Studied ratios in which elements combine Dalton put together the first atomic theory

186 Dalton’s Atomic Theory
All elements are composed of tiny indivisible particles called atoms Atoms of the same element are identical. The atoms of any one element are different from those of any other element Atoms of different elements can physically mix together or can chemically combine with one another in simple-whole number ratios to form compounds Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction

187 Finding how many subatomic particles for each atom
Atomic Number - whole number on p.t. Gives the number of protons Atoms are electrically neutral; there, positives equal negatives. Atomic number also equals number of electrons

188 Finding how many subatomic particles for each atom
Mass number = protons plus neutrons Mass number = p no Mass number is not found on the periodic table So, nO = mass number - p+ If carbon has a mass number of 14, how many e-’s, p+’s, and no’s does it have?

189 Symbols Mass number is in top left and atomic number is in bottom left
16O 9Be 8 4 How many subatomic particles in each? Oxygen - Beryllium -

190 Isotopes Atoms with the same number of protons but different numbers of neutrons Ex) Carbon 12 vs. Carbon 14 These atoms have a different mass Chemically alike because still have the same number of protons

191 Isotopes of Hydrogen Hydrogen -1 simply called hydrogen
Hydrogen called deuterium Hydrogen called tritium

192 Development of AMUs Atomic Mass Units (AMUs)
Protons have a mass of 1 amu 1.67 x g Neutrons have a mass of 1 amu Electrons have a mass of 0 amu 9.11 x g

193 Atomic Mass The weighted average mass of the isotopes in a naturally occurring sample of the element Don’t confuse with “mass number” To calculate atomic mass you need 3 pieces of information 1. The number of stable isotopes 2.The mass of each isotope 3.The natural percent abundance of each isotope

194 Atomic Mass Example Problem - Calculate the atomic mass for element X. One isotope has a mass of 10 amus (10X) and is 20% abundant. The other has a mass number of 11 amus (11X) and an abundance of 80%. To solve: Multiply the mass number times the abundance than add them together.

195 Atomic Mass 10 x 0.20 = 2.0 11 x 0.80 = 8.8 Add = 10.8 The atomic mass of element X is 10.8 amus

196 Atomic Mass Your turn. Solve:
What is the atomic mass of Element Z? The isotopes are 16Z, 17Z, 18Z; with percent abundances of , 0.037,

197 Atomic Mass Answer

198 Atomic (Hotels) Theory
Floors of the hotel are known as energy levels The period corresponds to the floor. The number of the floor is called the Principle Quantum Number Energy levels are divided into sublevels. These are the rooms of our atomic hotel. There are different types of rooms. Atomic Orbital - Region of high probability of finding an electron There are 4 types of rooms (orbitals). s, p, d, f rooms s - Spherical p - dumbbell d- clover leaf f - too complex to describe. s - superior p - preferred d- desirable f- fair Sharp principal diffuse fundamental

199 Atomic (Hotels) Theory
Three Managers each with their own rule: 1. Aufbau principle Rooms closest to the basement are checked out first Electrons enter orbitals of lowest energy first 2. Pauli Exclusion Principle - no more than 2 per room An atomic orbital contains a maximum of 2 electrons Roommates must have opposite spins to share a room

200 Atomic (Hotels) Theory
3. Hund’s Rule Similar rooms must have an occupant before pairing up as roommates When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins

201 S-orbitals of the 1st and 2nd energy levels.

202 P-orbital 2nd energy level







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

210 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 Atomic Number Atomic Mass



213 Focus Questions to upcoming video 8 Questions
What is the periodic law? Who is Dmitri Mendeleev and what did he do? How is today’s periodic table different from Mendeleev’s? What characteristic is common among the noble gases What are the names for vertical columns and horizontal rows What characteristics are common amongst group 1A? What are the 2 most important things about the periodic table? Why is fluorine the Tyrannosaurus rex of the periodic table?

214 Atoms

215 Periodic Table and Electron Configurations
Build-up order given by position on periodic table; row by row. Elements in same column will have the same outer shell electron configuration.

216 The relation between orbital filling and the periodic table

217 Electron Configuration
Orbitals have definite shapes and orientations in space (insert Fig 2.11 of text) (if it will not all fit on one screen, put part (a) on one screen and part (b) on the next)

218 Orbital occupancy for the first 10 elements, H through Ne.

219 Atomic radii of the main-group and transition elements.
Trends in the Periodic Table Atomic radii of the main-group and transition elements.

220 Trend for atomic radii Left to right atoms get smaller Why?
Increase in nuclear charge More protons and more electrons means greater electrostatic attractions (stronger magnet) Top to bottom atoms get larger Increase in energy levels (You are adding floors to your hotel). Electrons are further from the nucleus

221 Atomic Radius Atomic Radii for Main Group Elements
Atomic radii actually decrease across a row in the periodic table. Due to an increase in the effective nuclear charge. Within each group (vertical column), the atomic radius tends to increase with the period number.

222 Trend for Ion Size Ion is a charged atom.
Metals lose electrons and nonmetals gain electrons to create ions. Cations are pawsitive (positive) and Anions are negative. Cations are smaller than their corresponding atom. Why? Loss of electrons means the positive nucleus has a greater attraction on the remaining electrons Anions are larger than their corresponding atom. Why? Gain of electrons means the nucleus has less attraction for the electrons as well as the electrons are repulsing each other causing an increase in the size of the electron clouds

223 Radii of ions This is a “self-consistent” scale based on O-2 = 1.40 (or 1.38) Å. Ionic radii depend on the magnitude of the charge of the ion and its environment. (more later) Positively charged ions are smaller than their neutral analogues because of increased Z*. Negatively charged ions are larger than their neutral analogues because of decreased Z*. Same periodic trends as atomic radii for a given charge

224 Trend for ion size Decrease across a period then jumps in size at nonmetals and continues to decrease Increases on the way down a group as you are adding energy levels (electrons are farther from the nucleus)

225 Ionization energy The energy required to remove an electron

226 First ionization energies of the main-group elements
Trends in the Periodic Table


228 Ionization Energy Ionization energy is a periodic property
In general, it increases across a row. Why? increasing attraction as the number of protons in the nucleus increases (stronger magnet) it decreases going down a group. Why? Outer shell electrons are further from the nucleus so less electrostatic attraction. Nucleus has less pull on them. Shielding also plays a factor.

229 6) The trend across from left to right is accounted for by a) the increasing nuclear charge.

230 Electronegativity - This is the most important trend to understand for this class.
The tendency for an atom to attract electrons when chemically bonded. Same trend as ionization energy. In general, it increases across a row. Why? increasing attraction as the number of protons in the nucleus increases (stronger magnet) it decreases going down a group. Why? Outer shell electrons are further from the nucleus so less electrostatic attraction. Nucleus has less pull on them. Shielding also plays a factor.

231 See chart in book for summary
Trends in three atomic properties See chart in book for summary

232 Check for understanding
Which of the following atoms has the largest atomic radii, ion size, electronegativity, and ionization energy Na, Mg, K, Ca, S, Cl, Se, Br

233 Ionic Metallic Covalent
Bonding Ionic Metallic Covalent

234 Valence Electrons Atoms in a group behave similarly because they have the same number of valence electrons. Valence electrons - electrons in the highest occupied energy level To find the number of valence electrons just look at the group number

235 Lewis Electron Dot Structures
Symbol of element with dots around it representing valence electrons Example: C  2 electrons per side totaling 8  No pairs until each room has an electron

236 Lewis Electron Dot Structures
Example: O Pairs of electrons are adjacent not across from one another. This will help with identifying shape. Unshared Pair or Lone Pair

237 Octet Rule In forming compounds, atoms tend to achieve the noble gas electron configuration. They lose, gain, or share electrons with another atom to achieve 8. Metals lose electrons leaving a complete octet in the lower energy level Nonmetals gain electrons to fill the energy level to achieve 8.

238 Electron Configurations of Ions
Na - 1s22s22p63s1 Na1+ - 1s22s22p6 O - 1s22s22p O2- - 1s22s22p6 Write the following on your periodic table. Group 1A loses 1 electron Group 2A loses 2 electrons Group 3A loses 3 electrons Group 4A Depends on the atom Group 5A gains 3 electrons Group 6A gains 2 electrons Group 7A gains 1 electron Group 8A (0) - does not form ions



241 VSEPR Theory Valence Shell Electron Pair Repulsion Theory
“Electron pairs around atoms tend to be as far apart as possible.” Similar charges (I.e., negative charges from electrons) tend to repel each other and want to be spaced apart at maximum angles. Used to predict molecular geometries Bond angles Angles between bonds Spacing apart as far as possible Lone pairs of electrons will repel bonded atoms a bit more than expected toward each other around the central atom



244 Drawing Lewis Dot Structures
Lewis Dot Structure - A symbolic description of the distribution of valence electrons in a molecule. Dots are used to represent individual electrons and lines are used to represent covalent bonds. Lines are not drawn for ionic bonds!!! Drawing Lewis Dot Structures 1. Add up the total number of valence electrons in the molecule by totaling the valence electrons on each atom in the molecule or polyatomic ion. For example let’s work the following together: PO43-, O3 , BrF5

245 2. Draw the skeleton structure of the molecule or polyatomic ion in which the covalent bonds between the atoms are drawn as single lines. Each bond equals two valence electrons. If the molecule has more than two atoms, the atom with the lowest electronegativity is generally the central atom and is written in the middle.

246 3. Distribute valence electrons around the outer atoms as nonbonding electrons until each atom has a complete outer shell (i.e. 8 electrons except for H which has only 2 valence electrons). 4. Add the remaining valence electrons to the central atom.

247 5. Check the central atom. * If the central atom has eight electrons surrounding it, the Lewis Structure is complete. * If the central atom has less than eight electrons, remove a nonbonding electron pair from one of the outer atoms and form a double bond between that atom and the central atom. If needed, continue to remove nonbonding electron pairs from the outer atoms until the central atom has a complete octet. * If the central atom has more than eight electrons, then this means that the central atom has expanded its valence shell to hold more than eight electrons. This is allowed for atoms with valence shells in the third energy level or higher (i.e. in or beyond the third period of the periodic table).

248 Other notes Length of bond Energies of bonds
Single bonds are longer than double which are longer than triple. C-C =154 pm, C=C =134 pm, C≡C =120 pm) Energies of bonds Triple bonds have more energy than double and double have more energy single. C-C =348 kj/mole, C=C =614 kj/mole, C≡C =839 kj/mole)

249 When the electron number is odd the octet rule is broken. Example NO:

250 Polar Bonds and Molecules
Covalent Bonding Polar Bonds and Molecules

251 Covalent Bonding -- Polar Bonds and Molecules --
Bond Polarity “The Tug of War” The pairs of electrons that are bonds between atoms are pulled between the nuclei of the atoms in a bond. The electronegativities of the atoms determines who is winning Yet; there is no winner. The tug of war never ends. Classifications for Bonds Nonpolar covalent When atoms pull the bond equally Happens with two atoms of equal electronegativity, most often using the same atoms Examples: H2, O2, N2 Polar covalent When atoms pull the bond unequally Happens with two atoms of different electronegativities Example: HCl, HF, NH

252 Covalent Bonding -- Polar Bonds and Molecules --
Bond Polarity In a polar molecule, one end of the molecule is slightly more electronegative than the other atom, resulting in one atom being slightly negative (-) because of higher electronegativitiy, and the other atom being slightly positive (+) because of lower electronegativity.  is known as a partial charge since it is much less than 1+ or 1- charge.

253 Covalent Bonding -- Polar Bonds and Molecules --
Bond Polarity Electronegativities and Bond Types H: Cl: Since hydrogen is less, it will have the positive partial charge while chlorine has the negative partial charge. 3.0 – 2.1 = 0.9 HCl is polar covalent. 0.0 – 0.1 difference Nonpolar covalent bond H – H (0.0 difference) 0.1 – 1.7 difference Polar covalent bond H – Cl (0.9 difference) 1.7 + difference Ionic bond Na+Cl- (2.1 difference)

254 Covalent Bonding -- Polar Bonds and Molecules --
Polar Molecules Dipole Molecule that has two poles Example: HCl from the previous page Polar vs. Nonpolar H2O and CO2 Both have 3 atoms; yet, One is polar and one is nonpolar. Why? Structure (with bond polarity) determines the molecules polarity.

255 3 video clips coming up




259 Get out your bonding sheet of chemistry
Yes, we are going to discuss more on bonding so try and hold your enthusiasm as rioting is not tolerated and things need to be accomplished. Bonding appreciates your cooperation and will sign autographs at the end of the period. Thank you.

260 Intermolecular Attractions
Attractions Between Molecules van der Waals forces Two types: dispersion forces and dipole interactions Dispersion forces Weakest of all molecular interactions Caused by movement of electrons Occurs in the BrINClHOF’s


262 Intermolecular Attractions Attractions Between Molecules
2nd van der Waals force Dipole interactions Occurs when polar molecules are attracted to one another Partial charge (+) of one polar molecule is attracted to the opposite partial charge (-) of another molecule


264 Intermolecular Attractions attractions between molecules
Hydrogen bonding Hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom Example: water


266 Short Lab - No lab report
In the back in a tray is a micropipet and a penny. Place 1 drop of water on the penny and then take a guess as to how many drops of water you can fit on a penny. Write your guess on the board at the front of the room. Place your name next to your guess. Then count the drops of water you fit on the penny before it overflows. Record this count on the board next to your guess. As your water drop grows, watch it from the side. Clean up and have a seat at your desk.


268 Gases There are four variables that affect a gas. 1. Pressure
2. Volume 3.Temperature 4. Number of molecules

269 The variables Pressure units - there are many units for pressure.
kPa - kilopascal (101.3) Atm - atmospheres (1 ) mm Hg - millimeters of Hg (760) - torr Volume is measured in Liters Temperature is in Kelvin K = oC or oC = K - 273 If the Kelvin temperature doubles the K.E. doubles.

270 The pressure-volume relationship Boyle’s Law
Pressure and volume are inversely related. One goes up the other goes down P1 x V1 = P2 x V2


272 Boyle’s Law sample problem
A high-altitude balloon contains 30.0 L of helium gas at 103 kPa. What is the volume when the balloon rises to an altitude where the pressure is only 25.0 kPa? 103 kPa x 30.0 L = 25.0 kPa x V2 V2 = 124 L

273 Boyle’s Law sample problem
Your turn Your birthday balloon travels with you from Lincoln to Denver. The balloons volume is 4.0 L with an atmospheric pressure of kPa. You arrive in Denver where the atmospheric pressure is 90.0 kPa. What is the new volume of your balloon?

274 Boyle’s Law sample problem

275 The temperature-volume relationship Charles’s Law
Volume and temperature have a direct relationship. One goes up the other goes up One goes down the other goes down V1 = V2 T T2


277 Charles’s Law sample problem
A balloon inflated in a room at 24oC has a volume of 4.00 L. The balloon is then heated to a temperature of 58oC. What is the new volume if the pressure remains constant? 4.00L = V2 297 K K V2 = 4.46 L

278 Charles’s Law sample problem
Your turn A balloon is inflated in a room at 24oC and has a volume of 4.00 L. The balloon is placed in a freezer and then removed the volume is now 3.25 L. What was the temperature of the freezer in Celsius?

279 Charles’s Law sample problem

280 The Temperature-Pressure Relationship Gay-Lussac’s Law
The pressure of a gas is directly proportional to the temperature of a gas Temperature goes up; pressure goes up P1 = P2 T1 T2

281 Gay-Lussac’s Law example problems
The gas left in a used aerosol can is at a pressure of 103 kPa at 25oC. If the can is thrown into a fire, what is the pressure of the gas when it reaches 928oC? 103 kPa = P2 298 K K P2 = 415 kPa

282 Gay-Lussac’s Law example problems
Your turn A container of propane has a pressure of kPa at a morning temperature 15oC. By mid afternoon the temperature has reached 32oC. What is the pressure inside the propane tank?

283 Gay-Lussac’s Law example problems

284 The combined gas law P1 x V1 = P2 x V2 T T2

285 The combined gas law example problem
The volume of a gas-filled balloon is 30.0 L at 40oC and 153 kPa. What volume will the balloon have at STP? 153 kPa x 30.0 L = kPa x V2 313 K K V2 = 39.5 L

286 The combined gas law example problem
Your turn A gas-filled balloon is 25.0 L at 35oC and 145 kPa. What is the temperature if the volume increases to 28.0 L and a pressure of 152 kPa?

287 The combined gas law example problem

288 The Ideal Gas Law PV=nRT P = Pressure V = Volume n = Number of Moles
R = ideal gas constant T = temperature in Kelvin


290 R -The ideal gas constant
Depends on unit of pressure L . Atm / K . mol 62.4 L . mmHg / K . mol (torr is mm Hg) 8.31 L . kPa / K . mol

291 Ideal Gas Law example problem
Calculate the pressure of 1.65 g of helium gas at 16.0oC and occupying a volume of 3.25 L? You will need g to moles and Celsius to Kelvin: 1.65 g He 1 mole He 4.0 g He = mol He K = oC ; = 289 K For this problem you will need to pick an R value. For this problem I will choose to use the R value containing kPa.

292 Ideal Gas Law example problem
P x 3.25 L = mol x 8.31 kPa . L x 289 K mol . K Do the algebra and solve; if you do it right, guess what? You get the answer right. Neat concept, huh? Maybe your mommy will give you a cookie. = 305 kPa Your turn How many moles of gas are present in a sample of Argon at 58oC with a volume of 275 mL and a pressure of atm.

293 Ideal Gas Law example problem

294 Surface Tension

295 Intermolecular Forces Bulk and Surface

296 heat solid  melt  heat liquid  boil  heat gas
Phase Changes Energy Changes Accompanying Phase Changes All phase changes are possible under the right conditions (e.g. water sublimes when snow disappears without forming puddles). The sequence heat solid  melt  heat liquid  boil  heat gas is endothermic. cool gas  condense  cool liquid  freeze  cool solid is exothermic.

297 Phase Changes Heating Curves

298 Heating Curve Illustrated

299 Vapor Pressure Explaining Vapor Pressure on the Molecular Level
Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface. Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium.

300 Vapor Pressure Volatility, Vapor Pressure, and Temperature
If equilibrium is never established then the liquid evaporates. Volatile substances evaporate rapidly. The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.

301 Liquid Evaporates when no Equilibrium is Established

302 Vapor Pressure Vapor Pressure and Boiling Point
Liquids boil when the external pressure equals the vapor pressure. Temperature of boiling point increases as pressure increases. Two ways to get a liquid to boil: increase temperature or decrease pressure. Pressure cookers operate at high pressure. At high pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required. Normal boiling point is the boiling point at 760 mmHg (1 atm).

303 Water

304 Ice Liquid water’s density is greatest 4oC.
Ice has a 10% greater volume; therefore, lower density.

305 Ice

306 Why does ice behave so differently?
As kinetic energy (speed of the molecules) decreases, hydrogen bonds hold water molecules in place. The water molecules are held in an open framework creating a hexagonal symmetry of molecules. This increases the volume.

307 Aqueous Solutions Water samples that contain dissolved substances.
Solvent - dissolving medium Solute - dissolved particles Example - salt water - NaCl is the solute and the water is the solvent Characteristics: Homogeneous Stable They do not settle out Both solvent and solute will pass through a filter

308 Describe the process of solvation.
Watch the next 2 video clips and then answer the above question.



311 Other solution questions:
What substances don’t dissolve in water? Why? Why does grease dissolve in gasoline and not water? How does soap work? How is the relationship summed up?

312 All ionic compounds are electrolytes
Compounds that conduct an electric current in aqueous solution or molten state. All ionic compounds are electrolytes

313 nonelectrolytes Do not conduct an electric current in aqueous solution or molten Many molecular compounds are nonelectrolytes Ex sugar, rubbing alcohol Some very polar molecules become electrolytes when dissolved in water. Why? Because they ionize in solution HCl + H2O --> H3O+ + Cl- Weak electrolyte - only a fraction of solute exists as ions Strong electrolyte - large portion of the solute exists as ions

314 Henry’s Law - The solubility of a gas in a liquid is directly proportional to the pressure of the gas above the liquid. - pressure goes up, solubility goes up Mathematically stated S1 = S2 P P2

315 Example problem – you try
If the solubility of a gas in water is 0.77 g/L at 3.5 atm of pressure, what is its solubility (in g/L) at 1.0 atm of pressure? Answer How does temperature affect the solubility of a gas?

316 Thermochemistry -- The Flow of Energy: Heat --
Thermochemistry: the study of heat changes in chemical reactions Chemical potential energy: energy stored within the structural units of chemical substances

317 Thermochemistry -- The Flow of Energy: Heat --
Chemical System Types System type Description q (change in heat) Endothermic System absorbing heat from the surroundings q > 0 Exothermic System releasing heat to the surroundings q < 0

318 Thermochemistry -- The Flow of Energy: Heat --
Law of Conservation of Energy: In any chemical or physical process, energy is neither created nor destroyed

319 Thermochemistry -- The Flow of Energy: Heat --
The calorie Expressed as a c (lower case) Quantity of heat needed to raise the temperature of 1 g of pure water 1C Calorie Expressed as a C (upper case) Dietary Calorie 1 Calorie = 1 kilocalorie = 1000 calories

320 Thermochemistry -- An Intro Video --

321 Thermochemistry -- The Flow of Energy: Heat --
Specific Heat - a physical property (intensive property) - describes how much heat a substance can hold Water vs Metal - how does the temperature of a swimming pool compare from 3 p.m. to 3 a.m.? - how does the surface of your car hood compare from 3 p.m. to 3 a.m.?

322 Thermochemistry -- The Flow of Energy: Heat --
Joule SI unit of heat and energy Raises the temperature of 1 g of pure water C 4.184 J = 1 cal Heat Capacity Amount of heat needed to increase the temperature of an object exactly 1C Will change depending on the mass and chemical composition Specific Heat Quantity of heat needed to raise the temperature of 1g of substance 1oC

323 Thermochemistry -- The Flow of Energy: Heat --
Specific Heat Capacity Heat (q) specific heat capacity (C) Mass (m) change in temperature (T) q = mC T

324 Thermochemistry -- The Flow of Energy: Heat --
Example: How many kilojoules of heat are absorbed when 1.00 L of water is heated from 18C to 85C? Solution: q = mCT q = 1000g x 4.18 J x 67oC g o C q = 2.8E5 J 1 KJ 1000 J = 280 KJ

325 Thermochemistry -- The Flow of Energy: Heat --
Example: A chunk of silver has a heat capacity of 42.8 J/C. If the silver has a mass of 181 g, calculate the specific heat of silver. Solution: q = mCT 42.8 J = 181g x C x 1OC C = J/goC

326 Thermochemistry -- Measuring and Expressing Heat Changes --
Your Turn: The temperature of a piece of copper with a mass of 95.4 g increases from 20.0oC to 43.0oC when the metal absorbs 849 J of heat. What is the specific heat of copper?

327 Thermochemistry -- Measuring and Expressing Heat Changes --

328 Heterogeneous Aqueous Systems
1. Difference in types of heterogeneous aqueous systems is particle size Suspensions - particles settle out of solution can be filtered ex muddy water Colloids - 1. particles stay suspended in dispersion medium 2. many are cloudy - may look clear when diluted 3. exhibit Tyndall effect - scattering of light ex - paints, glues, gelatin desserts, etc Solutions - we already went over solutions

329 Colloidal Systems Emulsions - Colloidal dispersion of liquids in liquids needs an emulsifying agent ex - soap

330 Properties of acids and Bases
Taste Acids taste sour ex - lemons Bases taste bitter ex - soap Feel Acids feel like water; but have you ever gotten fruit juice on a canker sore or cut Bases feel slippery ex - soap and water

331 Properties of acids and bases
Reaction with metals Acids - Hydrogen gas is produced when reacted with certain metals Bases - typically don’t react with metals Both are electrolytes React to form salt and water Milk of Magnesia (Magnesium Hydroxide) is a base used to treat excess stomach acid problems.


333 Definitions of Acids and Bases
Arrhenius Acid/Base - focused on products HCl + H2O --> H3O+ + Cl- NaOH + H2O --> Na+(aq) + OH-(aq) Acids form H+’s and Bases form OH-’s

334 Definitions of Acids and Bases
Bronsted-Lowery Acid/Base - focused on what happens during formation HCl + H2O --> H3O+(aq) + Cl-(aq) Acid - substance that donates a proton Base - substance accepts a proton In the above example, what is the base and what is the acid? How about this example? NH3 + H2O --> NH4+(aq) + OH-(aq)

335 Definitions of Acids and Bases
Lewis acid/base H+ + OH- ---> H2O Acid - a substance that can accept a pair of electrons to form a covalent bond Base - a substance that donates a pair of electrons to form a covalent bond What is the Lewis acid and base? AlCl3 + Cl- --> AlCl4-

336 Problem Write an equation for the ionization of nitric acid and explain how it fits each definition?

337 Hydrogen Ions from Water This is the “bases” to start understanding pH
Water is considered neutral Collision between water molecules can cause a hydrogen ion to transfer from one molecule to another. H2O H2O H3O OH- hydronium ion hydroxide ion

338 Self-Ionization of Water
Water self ionizes to the concentration of 1.0 x 10-7 mol/L. When the concentration of each ion equals 1.0 x 10-7 mol/L the solution is said to be neutral Therefore, since water is considered neutral the concentrations of the ions can be calculated through the Ion product constant.


340 The Ion-Product Constant Kw
Notation [H+] - concentration of hydrogen ions Or hydronium ions [OH-] - concentration of hydroxide ions When [H+] and [OH-] are multiplied we get the ion-product constant. Kw = [H+] x [OH-] = 1.0 x (mol/L)2 or M2 This is an inverse relationship. One goes up, the other goes down.

341 The Ion-Product Constant Kw example problem
If [H+] = 1.0 x 10-5 mol/L, is the solution acidic, basic, or neutral? What is the [OH-] of this solution? Answer Acidic - the [H+] is greater than 1.0 x 10-7 mol/L 1.0 x 10-5 mol/L x [OH-] = 1.0 x M2 [OH-] = 1.0 x 10-9 mol/L

342 The Ion-Product Constant Kw example problem
If [OH-] = 2.8 x 10-8 mol/L, is the solution acidic, basic, or neutral? What is the [H+] of this solution? Answer

343 The pH concept [H+] is cumbersome so the pH scale was created.
pH is the negative logarithm of the hydrogen-ion concentration. pH = -log[H+]

344 Sample pH problems 1 of 3 The hydrogen-ion concentration of a solution is 2.7 x mol/L. What is the pH of the solution? Answer

345 Sample pH problems 2 of 3 The pH of a solution is 6.8. What is the [H+]? Answer

346 Sample pH problems 3 of 3 What is the pH of a solution if the [OH-] = 4.0 x mol/L Answer

347 pOH pH + pOH = 14 -log[H+] + -log[OH-] = 14
Example – If the H+ is 7.2 x 10-9 mol/L what is the pOH? 5.9

348 Naming Acids And writing formulas

349 General Form - HX (X is an anion or polyatomic ion)
Rules - 3 of them 1. When anion ends in -ide, acid name begins with hydro and -ide is changed to -ic with the word acid Ex - HCl is hydrochloric acid Try - H2S

350 Rule 2 - When anion ends in -ite, ending changes to -ous with the word acid. (No hydro) Name these - H2SO3 , HNO2 Rule 3 - When anion ends in -ate, ending changes to -ic with the word acid Name these - HNO3 , HC2H3O2

351 Work backwards to get the formula
Write the formula for the following: Chloric acid Hydrobromic acid Phosphorous Acid Don’t confuse phosphorous and phosphorus Answers

352 Neutralization Reaction
Base and acid react to produce a salt and water

353 Titration A lab technique where a neutralization reaction is performed to determine the concentration of an unknown.

354 The anatomy of a titration:
Standard solution - solution of known concentration. End Point - the point at which neutralization is achieved Indicator - chemical that changes color with a change in pH. We will be using phenolphthalein. Clear in an acid pink in a base You want a light pink Buret - Measurement device


356 Titration Calculations 4 steps
Start with the balanced equation Find the moles in the standard solution Set up ratio to find moles of unknown Find molarity(mol/L) or volume

357 Titration Calculations example problem
A ml solution of H2SO4 is neutralized by ml of 1.0 M NaOH. What is the concentration of H2SO4? H2SO NaOH --> Na2SO H2O L NaOH 1.0 mol NaOH 1 mol H2SO4 1 L NaOH mol NaOH L H2SO4 = 0.35 mol/L H2SO4

358 Titration Problems your turn
What is the molarity of phosphoric acid if 15.0 mL of the solution is neutralized by 8.5 mL of 0.15 M NaOH?

359 Titration Problems 1 more your turn
How many milliliters of 0.45 M hydrochloric acid must be added to 25.0 mL of 0.15 M NaOH?

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