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

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

Not the real Mr. Cooper

Info

• Class: Chemistry

• Instructor: Mr. Cooper

• Office: A112 - I’m pretty much in my classroom before and after school

• E-mail: [email protected]

• WebPage: lsw.lps.org click on teachers and find my name.

mailto:[email protected]

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.

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.

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

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.

Grading

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

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.

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.)

Grading

• There will be no test retakes!!!!

• However, you will have a take home practice test. So, your unit test is actually your retake.

Grading

• The grading scale is as follows:

• A= 90-100 B+= 85-89

• B= 80-84 C+=75-79

• C= 70-74 D+= 65-69

• D= 60-64 F= Below 60.0

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.

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.

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.

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.

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

• 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

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

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!!!

Tirrill (Bunsen) Burner

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.

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.

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

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.

Troubleshooting

• Gas attempts to light but goes out. Possible cause is:

• Air vent is too far open. Turn the barrel down.

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.

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.

Sample lab reportTitle: 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.

Sample lab reportConclusion: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.

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.

Acronyms

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

Metric Conversionsgrams (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

Metric Conversionskilo 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

Metric Conversionskilo hecto deka SI deci centi milli

• Your turns - convert the following:

• 15.72 g = ______ mg 15.72 g = _____ kg

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

Density Problem

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

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?

Properties of MatterDefinitions

• 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

Properties of MatterStates of Matter (Solid)

• Definite shape

• Definite volume

• Is incompressible (atom or molecules can not be pushed closer together)

• Examples -– coal, sugar, ice, etc

Properties of MatterStates of Matter (Liquid)

• Matter that flows

• Has a fixed volume

• Takes the shape of its container

• Incompressible

• Examples– Water, milk, blood, etc

Properties of MatterStates of Matter (Gas)

• Matter that takes the shape and volume of its container

• Easily compressed

• Examples– Oxygen, nitrogen, helium, etc

States of Matter Video

Properties of MatterPhysical Property

• An observed condition of the substance• Physical properties help identify substances• Examples include:

Properties of MatterPhysical 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

Change of State - a physical change

Properties of MatterChemical Property

• The ability or inability of a substance to rearrange its atoms.

• Example– Gasoline has the ability to react violently

with oxygen

Physical and Chemical Properties

Properties of MatterChemical 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)

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.

A chemical change

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

Categorizing our environment

Element Compound

Pure Substances

Heterogeneous

aka - solution

Homogeneous

Mixtures

Classifying Matter

Classifying MatterMixtures

• A physical blend of two or more substances.

• Examples:– Beef stew, air - mixture of gases

Classifying MatterMixture (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

Classifying matterMixture ( 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

Classifying MatterMixtures 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


Element Compound

Pure Substances

Examplesriver water

milkbeef stew

Heterogeneous

aka - solution

Examplespopsteel

kool-aid

Homogeneous

Mixtures

Classifying Matter

Classifying MatterElements

• Simplest form of matter

• Cannot be broken down into anything else

• Building blocks for all other substances

• Examples– Hydrogen, oxygen, carbon

Classifying MatterCompounds

• Two or more elements combined through a chemical bond

• Can only be separated into simpler substances by chemical reactions

• Example– Sugar - C12H22O11

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

Classifying matter reviewRemember

• 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


ExamplesCarbon (C)Gold (Au)Neon (Ne)

Element

ExamplesSodium Chloride (NaCl)

Sugar (C12H22O11)Dihydrogen Monoxide (H20)

Compound

Pure Substances

Examplesriver water

milkbeef stew

Heterogeneous

aka - solution

Examplespopsteel

kool-aid

Homogeneous

Mixtures

Classifying Matter

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


• Know where the following are on your periodic table (p.t) continued

• Transition metals• Inner transition metals• Alkali metals• Alkaline metals• Halogens• Noble gases

Ionic Moleculari.e. covalent

Naming Compounds

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 -

Naming Molecular Compounds

• Your turn. Try these.• Name or write the formula for:

– Boron trichloride– Dinitrogen tetrahydride– N2O5 PF5 S4N2 CCl4 SO3 H2O

• Answers

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

your worksheet.

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

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

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

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

Naming

• Your turn:– Name or write the symbol for the following:

• Aluminum Phosphide

• Calcium Ion Iodine

• Ga3+ Nitrogen

• K Sulfide

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


• 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


• 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:

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

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?

Writing formulas for mulitvalent ionic compounds

• Write formulas for the following:

• Copper (I) Oxide

• Copper (II) Oxide

Answers -

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 -

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 -

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

Polyatomic Ions• A group of atoms that carry a charge• Examples:

– SO42- NO3

1-

• 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 SO3

2-

– Chlorate is ClO31- Chlorite is ClO2

1-

• Other polyatomic ions– NH4

1+ Ammonium CN1- cyanide– OH1- Hydroxide

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

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

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)

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?

Name the cation firstthen name the anion

Example: Lithium FluorideMagnesium Carbonate

No

Name the cation firstPlace a roman numeral

Name the anionExample: Iron (II) Sulfate

Yes

Does the compound contain amultivalent ion?

aka - transition metal orgroup B element

Ionic

Yes

Use prefixes to representthe number of atoms.

Example: H2O Dihydrogen MonoxideCO2 Cabon Dioxide

Molecular

No

Is there a metal?

Naming Compounds

Check for understanding

• Name or write the formula for:– Potassium Sulfate– Chromium (III) Cyanide

– Fe(ClO3)3

– CuCl

• Answers

• 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.

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

Balance the following

• C2H6 + O2 --> CO2 + H2O

• Na3PO4 + Mg(NO3)2 --> NaNO3 + Mg3(PO4)2

Types of Reactions

Including reaction prediction

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

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)

1. Synthesis (Combination)

• Your turn. Write balanced equations for the following:– Aluminum (s) reacts with oxygen (g)– Hydrogen (g) reacts with oxygen (g)

• Answers:

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

2. Decomposition

• Your turn. Write balanced equations for the following:

• The decomposition of water

• The decomposition of lead (IV) oxide

• Answers

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

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)

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

–

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)

4. Double Replacement

• Your turn. Write balanced equations for the following.

• Iron (III) Sulfide reacting with hydrochloric acid

• Answer

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


• 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


• 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

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.

Precipitation reactionsPredicting 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

Precipitation ReactionsComplete Ionic Equations

• In an aqueous solution, substances exist as free ions. The equation shows this.

• Example for AgNO3(aq) + NaCl(aq)

Ag+1(aq) + NO3

1-(aq) + Na1+

(aq) + Cl1-

(aq) --> AgCl(s) + Na+1(aq) + NO3

1-(aq)

Precipitation ReactionsNet 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)

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) + NO3

1-(aq) + Na+1

(aq) + OH-(aq) --> Fe(OH)3(s) + Na+1

(aq) + NO31-

(aq)

• Fe3+(aq) + OH1-

(aq) --> Fe(OH)3(s)

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

The Mole

I. Molar Conversions

• 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!

Molar ConversionsConverting 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 unitsb. molecular compounds are called moleculesc. atoms are called atoms.

Example of representative units - 6.02 x 1023 atoms Cu 6.02 x 1023 molecules O2

6.02 x 1023 units NaCl

Molar Conversion Examples• How many moles of carbon are in

26.0 g of carbon?

Molar Conversion Examples• How many molecules are in 2.50

moles of C12H22O11?

Molar Conversion Examples• Find the mass of 2.1 1024

formula units of NaHCO3.

Molar Conversion Examples• Find the number of units of Iron

(III) Chlorate in 98.6 g of Iron (III) Chlorate.

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

Calculations w/ molar volume

• Determine the volume, in liters, of 0.60 mol SO2 gas at STP.

• Answer -– 0.60 mol SO2 22.4 L SO2

1 mol SO2 = 13 L SO2

Calculations w/ molar volumeYour 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 -

Molarity

• Unit of Concentration – There are many units of concentration

• Molarity is most useful to the chemist

M = moles of solute

Liters of solutionLiters 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.

Molarity ProblemsYou work them.

• A saline solution contains 0.90 g NaCl in exactly 100 ml of solution. What is the molarity of the solution?

Molarity ProblemsYou work them.

• How many moles of solute are present 1.5 L of 0.24 M Na2SO4?

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 CuCl2. Converting to grams

equals 17 grams.• Final answer

– Take 17 grams of CuCl2 and dissolve in enough water to make 500.0 ml of solution.

• Dissolve the 17 grams in say 400 ml of water. Once the CuCl2 is dissolved add water up to 500.0 ml.

Preparing a solutionyour turn

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

Making Dilutions

• 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.

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.

Dilution ProblemsYour Turn

• How would you prepare 90.0 mL of 2.0 M H2SO4 from 18 M stock solution?

• Answer

Dilution ProblemsYour 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 -

Percent Composition

% composition your turn

• Hydroxide makes up what percent of Calcium Hydroxide?

• Answer

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

Empirical vs. Molecular Formula

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).

• 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

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.

€

mol N = 25.9 g N

1•

1 mol N

14.0067 g N= 1.85 mol N

€

mol O = 74.1 g O

1•

1 mol O

15.9994 g O= 4.63 mol O

This would mean that the ratio of nitrogen to oxygenis 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 ofoxygen, we must multiply through each number by 2 to even it out, getting N2O5 as our empirical formula.

• 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.

• Give it a try• Determine the empirical formula for a

compound containing 7.8% carbon and 92.2% chlorine.

Empirical vs. Molecular Formula

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

Calculate the molecular formula of the compound whose molar mass is 180.1583 g and empirical formula

is CH2O.

€

180.1583

30.0264= 6

Now, in order to figure out what we must multiply each number in the empirical formula by, we must figure out bywhat 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 molarmass of 180.1583 g. We also know, by calculating the gmm of CH2O, that CH2O has an empirical formula mass (efm) = 30.0264 g CH2O.

• 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 12.0111 g + 12 x 1.00794 g + 6 x 15.9994 g = 180.1583 g C6H12O6

• This agrees with the molar mass we were given, so the molecular formula we calculated is correct.

• 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.

Example (toughy)

• 1.00 g of menthol on combustion yields 1.161 g of H2O and 2.818 g of CO2. What is the empirical formula?

• Solution:

Stoichiometry

Calculations of quantities in chemical reactions.

The use of ratios to calculate quantities

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.

Stoichiometry Example Problem #1

• How many moles of ammonia are produced when 0.60 mol of hydrogen reacts with nitrogen?


• Your Turn

• How many moles of aluminum sulfide are produced when 1.2 moles of aluminum reacts with sulfur?


• Answer


• How many grams of ammonia will be produced by reacting 5.40 g of hydrogen with nitrogen?


• Your Turn

• How many grams of aluminum are needed to react with 2.45 g of copper (II) chloride?


• Answer

% 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

% Yield

• Definitions

• Actual Yield

• The amount that is actually formed when the reaction is carried out in the laboratory.

% Yield =

actual yield X 100

theoretical yield

% yield will never be over 100%

Most likely it will never even be 100%

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

% 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?

% yield example problemanswer

• BaCl2 + K2SO4 --> BaSO4 + 2 KCl3. 75 g BaCl2 1 mol BaCl2 1 mol BaSO4 233.4 g BaSO4

208.3 g BaCl2 1 mol BaCl2 1 mol BaSO4

= 4.20 g BaSO4

3.89 g BaSO4 x 100 = 92.6 %

4.20 g BaSO4

% yield example problemyour 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?

% yield example problemanswer

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.

Limiting Reagent

• One track to determine limiting reagent

• A second track to determine product

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 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 Cu2S 159.1g Cu2S

63.5g Cu 2mol Cu 1mol Cu2S

= 1.00x102 g Cu2S

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?

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?

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 - ModelDalton’s ModelMore to come

Plum Pudding Model

• Proposed by J. J. Thomson (1856 - 1940), 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.

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

Nuclear Model• The Gold foil experiment or

the Rutherford experiment was an experiment done by Ernest Rutherford (1871 - 1937) in 1909. 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 space2. Dense center called the nucleus3. Electrons were stuck surrounding the nucleus.

Planetary Model• Introduced by Niels Bohr,

a Danish physicist (1885 - 1962), 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

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.

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?

Dalton’s Atomic Theory

• Democritus - Greek philosopher who first suggested atoms

• John Dalton (1766-1844)

• Studied ratios in which elements combine

• Dalton put together the first atomic theory

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

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

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?

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 -

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

Isotopes of Hydrogen

• Hydrogen -1 simply called hydrogen

• Hydrogen - 2 called deuterium

• Hydrogen - 3 called tritium

Development of AMUs

• Atomic Mass Units (AMUs)

• Protons have a mass of 1 amu

• 1.67 x 10-24 g

• Neutrons have a mass of 1 amu

• Electrons have a mass of 0 amu

• 9.11 x 10-28 g

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

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.

Atomic Mass

• 10 x 0.20 = 2.0

• 11 x 0.80 = 8.8

• Add 2.0 + 8.8 = 10.8– The atomic mass of element X is 10.8

amus

Atomic Mass

• Your turn. Solve:– What is the atomic mass of Element Z?

The isotopes are 16Z, 17Z, 18Z; with percent abundances of 99.759, 0.037, 0.204.

Atomic Mass

• Answer

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

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

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

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

P-orbital

2nd energy level


• 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


• 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

Focus Questions to upcoming video8 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?

Atoms

10-15 m

Neutron

Proton

Nucleus (protons and neutrons)

Space occupied by electrons

10-10 m

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.

The relation between orbital filling and the periodic table

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)

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

Atomic radii of the main-group and transition

elements.

Trends in the Periodic Table

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• Why?

– Increase in energy levels (You are adding floors to your hotel). Electrons are further from the nucleus

Atomic Radius• 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.

Atomic Radii for Main Group Elements

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

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*.

Radii of ions

Same periodic trends as atomic radii for a given charge

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)

Ionization energy

• The energy required to remove an electron

First ionization energies of the

main-group elements

Trends in the Periodic Table

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.

Ionization energy

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

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.

Trends in three atomic properties

See chart in book for summary

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

Bonding

Ionic

Metallic

Covalent

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

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

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

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.

Electron Configurations of Ions

• Na - 1s22s22p63s1 Na1+ - 1s22s22p6

• O - 1s22s22p4 O2- - 1s22s22p6

• Write the following on your periodic table.• Group 1A - 1+ - loses 1 electron• Group 2A - 2+ - loses 2 electrons• Group 3A - 3+ - loses 3 electrons• Group 4A - - Depends on the atom• Group 5A - 3- - gains 3 electrons• Group 6A - 2- - gains 2 electrons• Group 7A - 1- - gains 1 electron• Group 8A (0) - does not form ions

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

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

• 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.

• 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.

• 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).

• Other notes– Length of bond

• 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)

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

Covalent Bonding

Polar Bonds and Molecules

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 winningYet; 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


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.


Bond Polarity• Electronegativities and Bond Types

– H: 2.1 Cl: 3.0 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)


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.

3 video clips coming up

• 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.

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

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

Intermolecular Attractionsattractions 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

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.

Gases

• There are four variables that affect a gas.

• 1. Pressure

• 2. Volume

• 3.Temperature

• 4. Number of molecules

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 + 273 or oC = K - 273• If the Kelvin temperature doubles the K.E. doubles.

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

Boyle’s Lawsample 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


• Your turn

• Your birthday balloon travels with you from Lincoln to Denver. The balloons volume is 4.0 L with an atmospheric pressure of 101.3 kPa. You arrive in Denver where the atmospheric pressure is 90.0 kPa. What is the new volume of your balloon?


• Answer

The temperature-volume relationshipCharles’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

T1 T2

Charles’s Lawsample 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 331 K

V2 = 4.46 L


• 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?


• Answer

The Temperature-Pressure RelationshipGay-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

Gay-Lussac’s Lawexample 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 1201 K

P2 = 415 kPa


• Your turn

• A container of propane has a pressure of 108.6 kPa at a morning temperature 15oC. By mid afternoon the temperature has reached 32oC. What is the pressure inside the propane tank?


• Answer

The combined gas law

P1 x V1 = P2 x V2

T1 T2

The combined gas lawexample 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 = 101.3 kPa x V2

313 K 273 K

V2 = 39.5 L


• 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?


• Answer

The Ideal Gas Law

PV=nRT• P = Pressure

• V = Volume

• n = Number of Moles

• R = ideal gas constant

• T = temperature in Kelvin

R -The ideal gas constant

• Depends on unit of pressure

• 0.0821 L . Atm / K . mol

• 62.4 L . mmHg / K . mol (torr is mm Hg)

• 8.31 L . kPa / K . mol

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 = 0.413 mol He• K = oC + 273 ; 16. 0 + 273 = 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.


• P x 3.25 L = 0.413 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 0.987 atm.


• Answer

Surface Tension

Intermolecular Forces Bulk and Surface

Phase ChangesPhase ChangesEnergy Changes Accompanying Phase ChangesEnergy 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.• The sequence • cool gas condense cool liquid freeze cool solid

is exothermic.

Phase ChangesPhase Changes

Heating CurvesHeating Curves

Heating Curve Illustrated

Vapor PressureVapor Pressure

Explaining Vapor Pressure on Explaining Vapor Pressure on the Molecular Levelthe 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.

Vapor PressureVapor Pressure

Volatility, Vapor Pressure, and Volatility, Vapor Pressure, and TemperatureTemperature

• 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.

Liquid Evaporates when no Equilibrium is Established

Vapor PressureVapor PressureVapor Pressure and Boiling PointVapor 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).

Ice

• Liquid water’s density is greatest 4oC.

• Ice has a 10% greater volume; therefore, lower density.

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.

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

Describe the process of solvation.

• Watch the next 2 video clips and then answer the above question.

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?

Electrolytes

• Compounds that conduct an electric current in aqueous solution or molten state.

• All ionic compounds are electrolytes

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

• 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

• P1 P2

• 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?

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


Chemical System Types

System type Descriptionq (change

in heat)

EndothermicSystem absorbing heat from the surroundings

q > 0

ExothermicSystem releasing heat to

the surroundingsq < 0


Law of Conservation of Energy:

In any chemical or physical process, energy is neither created nor destroyed


The calorie• Expressed as a c (lower case)• Quantity of heat needed to raise the temperature

of 1 g of pure water 1CCalorie• Expressed as a C (upper case)• Dietary Calorie• 1 Calorie = 1 kilocalorie = 1000 calories

Thermochemistry-- An Intro Video --


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.?


Joule• SI unit of heat and energy• Raises the temperature of 1 g of pure water 0.2390C• 4.184 J = 1 calHeat Capacity• Amount of heat needed to increase the temperature of an object

exactly 1C• Will change depending on the mass and chemical compositionSpecific Heat

Quantity of heat needed to raise the temperature of 1g of substance 1oC


Specific Heat Capacity

Heat (q) specific heat capacity (C)

Mass (m) change in temperature (T)

q = mC T


Example:How many kilojoules of heat are absorbed when 1.00 L of water is heated from 18C to 85C?Solution:

q = mCTq = 1000g x 4.18 J x 67oC

g o C

q = 2.8E5 J 1 KJ 1000 J

= 280 KJ


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 = mCT

42.8 J = 181g x C x 1OC

C = 0.236 J/goC

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?

Thermochemistry-- Measuring and Expressing Heat Changes --

Calorimeter

Heterogeneous Aqueous Systems1. Difference in types of heterogeneous aqueous

systems is particle size

Suspensions - 1. particles settle out of solution2. can be filtered

ex muddy water

Colloids - 1. particles stay suspended in dispersion medium2. many are cloudy - may look clear when diluted3. exhibit Tyndall effect - scattering of light

ex - paints, glues, gelatin desserts, etc

Solutions - we already went over solutions

Colloidal Systems

Emulsions - Colloidal dispersion of liquids in liquids

needs an emulsifying agent

ex - soap

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

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.

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

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)

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-

Problem

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

Hydrogen Ions from WaterThis 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

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/Lthe 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.

The Ion-Product ConstantKw

• 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 10-14 (mol/L)2 or M2

• This is an inverse relationship.– One goes up, the other goes down.

The Ion-Product ConstantKw 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 10-14 M2

– [OH-] = 1.0 x 10-9 mol/L

The Ion-Product ConstantKw 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

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+]

Sample pH problems1 of 3

• The hydrogen-ion concentration of a solution is 2.7 x 10-10 mol/L. What is the pH of the solution?

• Answer


• The pH of a solution is 6.8. What is the [H+]?

• Answer


• What is the pH of a solution if the [OH-] = 4.0 x 10-11 mol/L

• Answer

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

Naming Acids

And writing formulas

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

• 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

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

Neutralization Reaction

• Base and acid react to produce a salt and water

Titration

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

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 baseYou want a light pink

• Buret - Measurement device

Titration Calculations4 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

Titration Calculations example problem

• A 25.75 ml solution of H2SO4 is neutralized by 18.23 ml of 1.0 M NaOH. What is the concentration of H2SO4?

• H2SO4 + 2 NaOH --> Na2SO4 + 2 H2O

• 0.01823 L NaOH 1.0 mol NaOH 1 mol H2SO4

1 L NaOH 2 mol NaOH 0.02575L H2SO4

• = 0.35 mol/L H2SO4

Titration Problemsyour 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?

Titration Problems1 more your turn

• How many milliliters of 0.45 M hydrochloric acid must be added to 25.0 mL of 0.15 M NaOH?

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

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