A.P. Chemistry Unit 1 Chapters 1*, 2, 3 Why take Chemistry?

A.P. ChemistryUnit 1

Chapters 1*, 2, 3

Why take Chemistry?

Purpose of AP Chem at CG?

Introductory terms

• Table groups

Definition Review: Try to fill in

1. Chemistry2. Chemical

Property3. Physical

Property

4. Intensive property

5. Extensive property

Terms to recognize6. Element7. ATOM8. Compound9. Molecule10. Formula Unit

TERMS11. Pure substances12. Chemical reaction13. Physical change14. Mixture15. Homogeneous16. Heterogeneous17. Chemical symbol18. Chemical formula

More terms

19. Chemical Equation20. Reactants21. Products22. Coefficients23. Subscripts24. Matter25. Law of conservation of

energy26. Law of conservation of

matter

27. Name Steps in the “Scientific Method”

1. Asking a question2. Forming hypotheses3. Researching previously found information4. Designing experiments5. Conducting experiments / collecting data6. Determining variables: dependent and independent7. Organizing and analyzing data8. Stating conclusions9. Considering sources of error10. Communicating results11. Planning future experiments

Classification of Matter

Solutions

Homogeneous Heterogeneous

Easily Separated

Mixture

Always Homogeneous

AtomsPeriodic Table

Element

Molecules or CrystalsFormulaic

Compound

Pure Substance

ALL Matter

Atoms Molecules Elements Compounds?

1. Aluminum foil 6. oxygen gas2. Carbon Dioxide 7. sodium chloride3. Zinc 8. water4. Graphite 9. chlorophyll5. Helium 10. nitrogen

1. E, A 6. E, M2. C, M 7. C, “?M”3. E, A 8. C, M4. E, A 9. C, M5. E, A 10. E, M

Solid, Liquid, Gas(room temperature)1. Ammonia2. Gasoline3. Graphite4. H2O5. Shaving cream6. Aluminum7. Ice cream8. Helium9. Bromine10.Sugar

Solid, Liquid, Gas(room temperature)1. G2. L3. S4. L5. L & G …?6. S7. L8. G9. L10. S

Pure Substance or Mixture?

1. Water2. Hydrogen3. Salt4. Tea5. Sodium6. Sugar7. Iron oxide8. concrete

9. Raisin cookie10. Gatorade

Pure Substance or Mixture?

1. P2. P3. P4. M5. P6. P7. P8. M

9. M10. M

Chemical or Physical Change?

1. Dog is groomed2. Child gets taller3. Gas forms when

Baking soda is mixed with vinegar

4. Pencil is sharpened5. Paper burns6. Leaves turn color7. Ice melts

8. Sugar dissolves in water9. Cookie bakes in oven10. Cake mix is combined

with water.

Chemical or Physical Change?

1. P2. C3. C4. P5. C6. C7. P

8. P9. C10. P (?)

Metric fundamental units

• Kilogram• Meter• Second

• mole• Kelvin• Coulomb

Metric units for?....

1. Length 2. mass 3. area4. volume5. Density6. Weight and Force

7. Energy8. particles 9. Pressure10. Current11. Potential12. Power

fundamentals?....

1. yes 2. yes 3. yes4. yes5. yes6. Kg*m/s2

7. Kg*m2/s2

8. yes 9. Kg*m/s2

10. C/s11. Kg*m2/Cs2

12. Kg*m2/s3

13. yes

Metric units?

• Fundamental?

• Can you convert?

• Deci• Centi• Milli• Kilo• Nano• Giga• Micro• Mega

AP Chemistry

Chapter 2Atoms, Molecules and Ions

John Dalton

• English school teacher• 1766-1844• Author of the

Modern Atomic Theory

• loved studying the weather • saw the applications for

chemistry in his ideas about the atmosphere.

• Proposed Atomic Theory: 1803• Dalton's theory was presented

in New System of Chemical Philosophy (1808-1827).

John Dalton

• Was colorblind• Daltonism• His eyes were used to

prove it is a brain disorder

Postulates of Dalton’s atomic theory

Matter is made of atoms which stay the same during a chemical change.

An element is a substance made of one type of atom, each of which has the same properties.

A compound is matter made of two or more elements combined in fixed proportions.

A chemical reaction involves rearrangement of atoms into new substances, but no loss or gain of any atoms.

Law of definite proportions

• Molecules of the same compound are all the same.

The same elements can make many different compounds.

Atomic Structure

Modern Theory Says…

• Atoms are made of:

• Protons• neutrons • electrons • a small dense

nucleus of protons and neutrons and surrounding electrons.

Famous experiments leading to this view of the atom:

• Joseph John Thomson:

• The Cathode Ray Tube Experiment (Cambridge; 1897);

• discoverer of the electron, Nobel Prize Winner.

The Cathode Ray Tube and Thomson’s Plum Pudding atom

Robert Millikan

• The Oil Drop Experiment (USA, 1909)

• measured charge on an electron

• calculated the mass of electrons.

Millikan’s Experiment

• Used calculation to determine the charge on each suspended droplet

• All were multiples of 1.6x10-19 Coulomb

Ernest Rutherford

• Studied Gold Foil Experiments by Geiger and Marsden (1911)

• atoms are mainly empty space with a small, massive, dense, positively charged nucleus.

• Also discovered the proton.

Gold Foil Experiment Video?

• Led to the idea of a nuclear atom

• Led to the idea that atoms are mainly made of empty space.

Modern Atomic theory

• the atom consists of 40 fundamental particles.

• The electron is a quark, but the proton and the neutron are not.

ISOTOPES• Atoms of the

same element that have different masses (different # of neutrons in the nucleus).

ISOTOPES TO KNOW• The three isotopes

of hydrogen:• Protium = H-1• Deuterium = H-2

Water made with this is called heavy water.

• Tritium = H-3 and is radioactive!

AZ symbols

• A: mass #, nuclear particles, P+N

• Z: atomic #, nuclear charge, P

• Neutrons =?• X = element symbol

Write an “AZ” nuclide symbol for

• strontium 90• silicon 30• radon 226

Atomic number vs mass number

Protons : nuclear charge vs protons + neutrons : number of nuclear particles

• Atomic weight vs atomic number

~Weight of protons, neutrons, electronsVs

Proton number: nuclear charge

Atomic particles

Atomic Weights

• All relative to the Carbon-12 isotope• Carbon-12 is the mass standard

• One mole of carbon 12 = 12 grams

Atomic Weights

• Represent: average mass of isotopes and their percent composition in nature.

Measuring atomic weight

• Units are “amu”• Atomic mass units

• One mole of amu = 1 gram

Measuring the weight of atoms

By Mass spectrometry. • An unknown is

compared to a known sample ( the standard).

• Particles are accelerated through a gas and bent by a magnetic field.

• The curvature of their pathway is measured and mass is calculated.

• F=Bvq.

Mass Spec and examples

• See worksheet

Example • Given data for Chromium, determine its average atomic weight.

• Isotope Mass Frac. Abundance

• Cr-50 49.9461 0.0435• Cr-52 51.9405 0.8379• Cr-53 52.9407 0.0950• Cr-54 53.9389 0.0236

51.9959 u

Copper’s two isotopes are mass numbers 63 and 65: What percent abundance is each if the average atomic mass is 63.5?

No calculator!

• What is the mass in grams of a mole of titanium atoms?

• What is the mass of one atom of Ca-40?

• Answer in g and amu

History of the periodic table

Dmitri Mendeleev (Russia)

• Wrote periodic law

• Chart based on atomic weights.

John Newlands, Great Britain

• law of octaves. • ridiculed b/c of

inconsistencies

Julius Lothar Meyer, Germany• periodic law around the

same time as Mendeleev.• not credited as

Mendeleev b/c ?? Predictions undiscovered…..

Henry Mosely• 1887-1915• Studied with

Rutherford, • measured nuclear

charge / atomic number of elements.

Henry Mosely• Reordered periodic

table by at.# and it is better

Henry Mosely• Volunteered for

service in WWI. Was a signal officer for the British Army and killed in action at Gallipoli in 1915.

The Periodic Table

• Periods: • 7 horizontal Rows

• Families / Groups: 18 vertical columns

Sketch and label

Binary

• Made of two elements.

• Examples?

Molecular

• Made of molecules• Nonmetal atoms

making compounds• Covalent: shared

electrons.

Ionic

• Made with a metal ion or ammonium and an anion…..

Ionic or Molecular?

• Na2CO3

• C4H10

• MgSO4

• Al2 (SO4)3

• KF

• CuBr2

• H2O

• Li2O

• NH4I

• RbClO

Making ions

• Cations vs anions

• Only changes in electrons

Formula unit vs molecule

• Ionic vs not ionic• Low ratio of ions vs

formula of particle

Ionic charges

• Metals vs nonmetals

• Metals make cations

• Nonmetals make anions

Ion names

• Metals vs nonmetals

• The –ide ending for nonmetals

Nomenclature practice time

• Use the ion sheets

• Use the flowchart

• Follow the rules

• Practice!

Name these Compounds

1. HClO2. HF3. H2O2

4. PbCrO4

5. LiC2H3O2

6. CO

7. CdI2

8. N2O5

9. CuSO4

10. SrBr2

11. H3PO4

12. Ca(NO3)2

Name these Compounds

1. Hypochlorous acid2. Hydrofluoric acid3. Dihydrogen dioxide4. Lead II chromate5. Lithium acetate6. Carbon monoxide

7. Cadmium iodide 8. Dinitrogen pentoxide9. Copper II sulfate10. Strontium bromide11. Phosphoric acid12. Calcium nitrate

Write Formulas for these Compounds

1. Cobalt III hydroxide2. Barium phosphate3. Magnesium chloride4. Aluminum iodide5. Sodium oxide6. Perchloric acid

7. Nitrous acid8. Oxalic acid9. Hydrobromic acid10.Diphosphorous

pentoxide11.Dinitrogen

monosulfide12.Silver carbonate

Write Formulas for these Compounds

1. Co(OH)3

2. Ba3(PO4)2

3. MgCl2

4. AlI3

5. Na2O

6. HClO4

7. HNO2

8. H2C2O4

9. HBr10. P2O5

11. N2S

12. Ag2CO3

Organic Chemistry Intro• carbon based with hydrogen

and oxygen mainly.

• Hydrocarbons: C and H only

• Alkanes: Hydrocarbons with only singly bonded carbon atoms.

• 1st rule: C bonds 4 times

• Alcohols: Have an -OH functional group

• Isomers: Molecules with the same formula but different structures

• R: general symbol for a carbon chain

ORGANIC NOMENCLATURE

• If R = __Carbon, it is called:

• 1 Carbon = meth(yl)• 2 Carbons = eth(yl)• 3 Carbons =prop(yl)• 4 Carbons =but(yl)• 5 Carbons =pent(yl)• 6 Carbons =hex(yl)• 7 Carbons =hept(yl)• 8 Carbons =oct(yl)• 9 Carbons =non(yl)• 10 Carbons =dec(yl)

Draw… and decide if isomers exist.

• Propane

• 1-butanol

• pentane

Name each. If they have any isomers, name one of them too.• CH3CH2CH2CHOHCH2CH2CH2CH2CH3

• CH3CH2CH2CH2CH2CH3

• CH3OH

STOICHIOMETRY

• Using a balanced equation to make theoretical predictions.

• Beqs show COM!

Simple reactions

• Combination (synthesis)

• Smaller reactants make a more complex product.

Simple Reactions

• Decomposition

• A reactant forms simpler products

Classify a-e

• Combination• a c

• Decompostion• b d e

Combustion

• Hydrocarbon and oxygen gas reactants

• CO2 and H2O products

Combustion: write reactions

• Octane

• 2-Hexanol

• 3-Heptanol

“weight” Units• Molecular or

formula weight: u or amu

• Molar mass: g/mol

Formula weight ormolecular weight of:

• A. carbon dioxide• 44.0 amu• B. water• 18.0 amu• C. oxygen gas• 32.0 amu• D. Table salt• 58.5 amu

Molar Mass of:

• A. carbon dioxide• 44.0 g/mol• B. water• 18.0 g/mol• C. oxygen gas• 32.0 g/mol• D. Table salt• 58.5 g/mol

Mole

• 6.02x1023

• Used to count particles

• A mole of miniature marshmallows would cover the USA to a depth of 600 miles.

Mole Relationships• 1 mole = 6.02x1023 at or mc or fu

• 1 gram = 6.02x1023 amu

• 1 mole = at wt of any formula (g)

• 1 mole gas (STP) = 22.4 L

Practice Mole Calculations

1. 0.0365 g 2. 1.0x1024 atoms3. 23.6 mol4. 1.1x10-21 amu5.250. Liters6.42.1%C, 6.4%H, 51.5%O7.51.9 %N

Example Problem• Determine the empirical and molecular formula of a compound found by combustion to contain 39.9% Carbon, 6.7% hydrogen and 53.4% oxygen. The molecular weight of the compound is 120 amu.

• Empirical: CH2O• Molecular: C4H8O4

Problem Solving• Pretend to have 100 grams …. Or….

• If grams are given, use them!

• Change grams to moles for each element.

• Look at mole ratios to work out lowest whole number subscripts.

• Use known molar mass to find molecular formula with integer multiplier.

• A sample of a compound is found to contain 17.5 g Na, 39.7 g Cr and 42.8 g O. What is its empirical formula?

• Na2Cr2O7

• Sorbic acid is added to food as a mold inhibitor. Its composition is 64.3% C, 7.20% H , and the rest oxygen. Its molecular weight is 112 u. What is the molecular formula for sorbic acid?

• C6H8O2

•Challenge Problem, AP level

An organic acid contains only C, H and O. A 12.72 mg sample of the acid is completely burned in oxygen. It yields 18.63 mg of carbon dioxide and 7.62 mg of water. What is the mass percentage of each element in the organic acid? What’s the empirical formula?

• Review: Write an equation for the formation of carbon dioxide from its elements.

• C + O2 ----> CO2

• How many grams carbon are needed to produce 150. grams of carbon dioxide?

• 40.9 grams

Problem Solving• Stoichiometry uses the balanced equation ratios.

• Balanced equation coefficients are about particle to particle ratios.

• Coefficients mean moles or atoms/molecules/formula units

• Or (by thinking of Avogadro’s Hypothesis) volumes(any units) for gaseous substances.

AMEDEO AVOGADRO

Stoichiometry

• Write an equation for the reaction of sodium in water.• 2Na + 2H2O ---> 2Na+ + 2OH- + H2

• How many molecules of hydrogen gas are produced by the reaction of 0.25 grams of sodium METAL?

• 3.3x1021 molecules

Stoichiometry

• Write the equation for the dehydration of ethyl alcohol and butanoic acid into ethyl butyrate, an ester.

• C2H5OH + C3H7COOH ---> C3H7COOC2H5 + H2O• How many grams of water can be made from 8.22x1023

molecules of ethyl alcohol (ethanol)?• 24.6 grams

Ethyl butyrate is the odor of pineapples.

Limiting Reagents and Percent Yield

• Write an equation for the synthesis of aluminum chloride.

• 2Al + 3Cl2 ---> 2AlCl3

• If 3.00 g Al react with 13.0 g Cl2, how much AlCl3 can be produced?

Use an IRF box

13 g =

0.183 moles

0

moles

2 Al + 3Cl2 ------> 2AlCl3

I Initial moles

R reacted moles

F

final moles

3 g =

0.111 moles

• Determine the limiting reagent.

• Use initial moles compared to how many are required for each reaction.

• Low number limits the process.

Use an IRF box

13 g =

0.183 moles

0

moles

* Limiting reagent!!

2 Al + 3Cl2 ------> 2AlCl3

I Initial moles

R reacted moles

F

final moles

3 g =

0.111 moles

Use an IRF box

13 g =

0.183 moles

0

moles

**

All 0.111 moles used

0 moles left over

2 Al + 3Cl2 ------> 2AlCl3

I Initial moles

R reacted moles

F

final moles

3 g =

0.111 moles

Fill in “R” Row

• Mole ratios in “R” row must match the reaction coefficient ratios.

• The next coefficient divided by LR coefficient, multiplied by limiting moles ---> “R” moles

Use an IRF box

13 g =

0.183 moles

0

moles

0.111 moles 0.167 moles 0.111 moles

2 Al + 3Cl2 ------> 2AlCl3

I Initial moles

R reacted moles

F

final moles

3 g =

0.111 moles

Fill in “F” row

• Subtract for reactants, add for products

• Once the box is filled in with moles, any question can be answered.

Use an IRF box

13 g =

0.183 moles

0

moles

0.111 moles 0.167 moles 0.111 moles

0 moles 0.016 moles 0.111 moles

2 Al + 3Cl2 ------> 2AlCl3

I Initial moles

R reacted moles

F

final moles

3 g =

0.111 moles

Now solve the problem!

• If 3.00 g Al react with 13.0 g Cl2, how much AlCl3 can be produced?

• 14.8 grams • If 12.0 grams is

recovered, what is the percent yield?

• 81.0%

Problem Solving• Thinking of the balanced

equation as a recipe might help.

• Determine how many times the reaction “recipe” can be carried out with each amount of moles.

• Reactant that can make the fewest “batches” is the limiting reagent.

Limiting Reagent and Percent Yield

• Write an equation for the synthesis of lithium hydroxide from lithium oxide and water.

• Li2O + H2O ---> 2 LiOH• If 42.0 grams lithium

oxide react with 20.0 grams water, how much LiOH can be produced?

• 53.1 grams• If 45.0 grams are

obtained, what is the percent yield of the experiment?

• 84.7%Lithium metal

Limiting Reagents and %Yield

• Write an equation for the preparation of hydrocyanic acid and water from ammonia, methane and oxygen.

• 2NH3 + 3O2 + 2CH4 --> 2HCN + 6H2O• How many grams of HCN can be obtained from the reaction of

25.0 grams ammonia, 75.0 grams oxygen and 25.0 grams methane?

• 39.7 grams• If 11.0 grams HCN is obtained, what is the percent yield of the

reaction process?• 27.7 %

That’s the end of Ch. 3!

• Mole day is coming SOON!• Make the moleata!• Talk to all other classes about donations for mole

day?

Mole Day

• CG Chemistry

T-Shirt Ideas

T-shirt Ideas

Build the Moleata!

A.P. Chemistry

Chapter 4Chemical Reactions: An

Introduction

I. Ionic Solutions

• Deionized water vs. tap water, bath water, lake water, ocean water?

• Ions!

Svante Arrhenius

• Svante August Arrhenius b. Uppsala, Sweden, February 19, 1859

• son of Svante Gustaf and Carolina

• educated at the Cathedral school; Showed an aptitude in mathematics and physics.

• 1876: University of Uppsala: mathematics, chemistry and physics.

• 1881: Stockholm’s Academy of Sciences.

• twice married - in 1894 to Sofia Rudbeck, (one son) and 1905 to Maria Johansson (one son and two daughters)

• died at Stockholm, October 2, 1927, and is buried at Uppsala.

Ionic solutions are .....

• electrolytic• Capable of

conducting electricity• Many ionic solids are

electrolytic in water.• the ions =

electrolytes (particles that conduct electricity)

What ions are found in a solution of...

• KOH?• CaCl2?• (NH4) 2SO4?• Write equations

for what these compounds do in water.

• KOH (s) --> K+ (aq) + OH- (aq)

• CaCl2 (s) --> Ca2+ (aq) + 2Cl- (aq)

• (NH4) 2SO4 (s) ---> 2NH4+ (aq) + SO4

2- (aq)

Nonelectrolytic substances

• Make no ions (electrolytes) in solution.

• Examples:• sucrose, C12H22O11 • methanol, CH3OH• urea, NH2CONH2

• antifreeze, HOC2H4OH• All are molecular

Dissolve but do not make ions• C12H22O11 (s)---->C12H22O11 (aq)• CH3OH(l) ----->CH3OH (aq)• NH2CONH2 (s)----->NH2CONH2 (aq)• HOC2H4OH (l)----->HOC2H4OH (aq)

Strong and Weak Electrolytes• refers to degree (%)

of ionization of solute.

• Acids and bases are described as strong or weak.

Strong vs. Weak Acids

• Strong acids ionize 100%

• Weak acids ionize only partially

Students must Know the 6 strong acids

• nitric• perchloric• sulfuric• hydrochloric• hydrobromic• Hydroiodic• Write equations

showing what each strong acid does in water.

The Strong Acids in Water

• Nitric HNO3 ---> H+ + NO3-

• Perchloric HClO4 ---> H+ + ClO4-

• Sulfuric H2SO4 ---> 2H+ + SO42-

• Hydrochloric HCl ---> H+ + Cl-

• Hydrobromic HBr ---> H+ + Br-

• Hydroiodic HI ----> H+ + I-

Dilute vs. Concentrated Solutions • refers to the amount

dissolved per volume of solution.

• dilute solutions: small amount dissolved.

• Concentrated: more dissolved

Writing Chemical Equations• can be done:

• A. molecularly: show the whole mixture.

• B. Ionically – complete ionic: indicates any electrolytes in

mix– Net ionic: only shows species that changed.

• AP CHEM: requires net ionic equation writing.

I. Precipitation Reactions• two aqueous solutions

are mixed and one of the products is insoluble.

• Aqueous: Dissolved in Water.

• Precipitate: insoluble species.

• Spectator Ions: in the mixture but do not take part in a reaction.

• SOLUBLE OR NOT?

• Know: Solubility rules of ionic compounds..... page 136.

• Knowing the solubility song / chart helps.

Examples: Write the net ionic reactions!• 1. Potassium Chloride and Silver

nitrate react in aqueous solution.

• Net ionic: Cl- + Ag+ ----> AgCl

• 2. Ammonium sulfate and calcium chloride react in aqueous solution.

• Net ionic: Ca2+ + SO42- ----> CaSO4

• 3. Sodium carbonate and copper II sulfate react in aqueous solution.

• CO32- + Cu2+ -----> CuCO3

More Practice. What will happen when the following mix?

• NiCl2 and Na3PO4

• NaCl and Fe(NO3)2

• Al2(SO4)3 and KOH

• Pb(C2H3O2)2 and NH4Cl

Combustion: burns in O2

• hydrocarbons

• Products always CO2 and H2O(l)

• exothermic

• Heat makes H2O vaporize.

Equations to Balance• Recognize some

organic alkanes and alkenes, alkynes and alcohols.

• Octane

• 2-Hexene

• 1-butyne

• 3-Heptanol

Acid-Base Basics

• Taste, feel ?• pH ranges ?• Ions they make in

solution ?• Household

examples of each• Page 139

The Strong Acids in Water

• Nitric HNO3 ---> H+ + NO3-

• Perchloric HClO4 ---> H+ + ClO4-

• Sulfuric H2SO4 ---> 2H+ + SO42-

• Hydrochloric HCl ---> H+ + Cl-

• Hydrobromic HBr ---> H+ + Br-

• Hydroiodic HI ----> H+ + I-

The Strong Bases in water:

• LiOH LiOH ---> Li+ + OH-

• NaOH NaOH ---> Na+ + OH-

• KOH KOH ---> K+ + OH-

• Ca(OH)2 Ca(OH)2 ---> Ca2+ + 2OH-

• Sr(OH)2 Sr(OH)2 ---> Sr2+ + 2OH- Ba(OH)2 Ba(OH)2 ---> Ba2+ + 2OH-

Acid and Base Definitions

Arrhenius

Acids contain hydrogen and make hydrogen ions in water. Bases contain hydroxide and make hydroxide ions in water.

Bronsted and Lowry acids and bases: Proton donors and proton acceptors.

Indicators to know• Indicators are

molecules that change colors at different pH levels.

• phenolphthalein• Litmus• Methyl red• Others?

Reactions of Acids/Bases

• Learn to Write Net ionic equations for many examples!

Neutralization reactions• acid + base water + salt.• salt: metal cation combined

with an anion (often soluble in water)

• General Equation

• HA + MOH -----> H2O + MA

• make examples…to net ionic

Carbonates and Acids

Some Salts + acids gases. Carbonates + Acids ----> carbon dioxide, water and a

salt. General Equation MCO33 + HA ----> H22O + CO22 + MA Write examples….. To net ionic. Demo and test for gas: Acetic acid and sodium

(bi)carbonate

Sulfites and Acids

• Sulfites + Acids -----> sulfur dioxide, water and a salt.

• General:

• MSO33 + HA -------> H22O + SO22 + MA

• examples: Write net reactions.

• SO2 gas…. Stinky, irritating

Sulfides and Acids

• A sulfide reacts with an acid to produce hydrogen sulfide gas and a salt.

• General: • MS +HA --> H22S + MA • examples• test for gas?

One Base Reaction

• Decompostition of Hydroxides

• MOH MO + H2O

• Examples…..

Oxidation and Reduction• Aka single displacement, synthesis or

decomposition reactions.

• electrons exchanged / atoms change oxidation states (charges).

Define: Oxidation and Reduction

• Oxidation: the loss of electrons

• Reduction: The gain of electrons

• LEO says “GER”

Rules for deciding Oxidation State:• Elements alone have an ox. state of zero.• H is always 1+, unless it is with a metal as a

hydride. • O is always 2- unless it is in a (rare) peroxide:

H22O22, Li22O22, K22O22, or Na22O22.

• All common /main group metals keep their periodic pattern charge.

• Other semi and nonmetal elements’ oxidation states are determined last.

• Sum of ox. States = charge on species.

Tell each elements’ Oxidation State:

1. H33PO44

2. KNO33

3. Ca(NO22)22

4. BrO2 -

5. BrO33-

6. BrO4 4 -

7. CH44

8. NH44Cl

9. Cl22O

10.N22O11.NO12.NO22

13.P22O55

14.KMnO44

15.Fe22(SO44)33

16.Na2C2O4

Types of “redox” reactions

• Synthesis / Combination reactions

• Decomposition Reactions

• “single replacement” reactions

• Combustion reactions• Other complex

reactions

What is oxidized, what’s reduced?

• Copper nitrate solution reacts with zinc metal to make aqueous zinc nitrate and metallic copper.

• Lithium metal reacts with a cobalt II chloride solution to make metallic cobalt and aqueous lithium chloride.

What is oxidized, what’s reduced?

• Ca + O2 ---> CaO

• HgCl2 ----> Hg + Cl2

Will redox occur? Using the “Activity Series”

• See the AP pages for the reduction potential list.

What is reduced and what is oxidized? What are Ox. And Red.

Agents? Write 1/2 reax.

• Examples:• iron nail in copper sulfate.• Aluminum foil in tin II chloride

solution.• Copper wire is placed in silver

nitrate solution.

CH. 19, section 1

• Balancing complex redox reactions in acid/base environments

Last Topic: Solution Chem/Stoich

“Volumetric Analysis”

Measuring Concentration of Solutions

• can be done in several different ways, including....

• Molarity• moles dissolved

per liter of solution.

Other Concentration

Definitions

• Molality: moles dissolved per kilogram of solvent

• Mass percentage: mass of solute compared to mass of solution.

• Mole fraction: moles dissolved compared to moles of total solution particles.

Molarity Equation

• M = n V

M = Molarityn = moles dissolved soluteV = volume of solution in

liters

Molality Equation

• m = n• kg

• m = molality• n = moles dissolved

solute• kg = mass of solvent

in kilograms.

Mass Percent Equation

• Mass of solute___• Mass of solution

• Express concentration as a percentage.

• Any units for mass will do.

Mole Fraction Equation

• Moles of solute ________________ • Moles of solute + moles of solvent

• Express result as a decimal number

Practice: Molarity Problems

• What is the molarity of a solution containing 34.2 grams of sucrose in 2.00 liters of solution?

• 0.0500 M• What volume of 0.65

M HCl holds 3.0 grams of HCl?

• 0.126 L or 126 mL

Molarity and Dilution Problems• What mass of silver

nitrate must be added to a flask to make 500. mL of 0.025 M solution?

• 2.1 g• How many mL of 4.00

M acetic acid are needed to make 500. mL of 1.00 M solution?

• 125 mL

More Problems! • How many mL of 8.4 M KNO3 are needed to make 3.00 Liters of 2.5 M solution?

• 890 mL• Suppose 200.mL of

water are added to 400.mL of 1.20 M HNO3. What is the molarity of the resulting mixture?

• 0.800 M

Stoichiometry with Molarity = Quantitative Analysis

• A 1.000 L sample of polluted water was analyzed for lead II ion by adding excess sodium sulfate to it. The mass of lead II sulfate precipitating was 220.0 mg What is the mass of lead in the water?

• 150.3 mg • What would be the

concentration of lead in the water?

• 0.000725 M

Solution Stoich • A flask contains water mixed with some HCl. The solution is titrated with 0.225 M KOH until a pH of 7 is reached. 15.20 mL of the KOH solution are needed. What is the mass of the HCl in the flask?

• 0.125 grams

Solution concentrations

• If 35.0 grams of potassium nitrate are dissolved in 55.0 grams of water, the solution has a density of 1.108 g/mL. Determine the molarity, molality, mass%, and mole fraction concentration values of the solution.

Learn to balance complex redox reactions: separate note page

Include Ch. 19 section 1 problems with the chapter 4 problem set.

Booknotes: not required for 19.1

Chemical or Physical Property? #1-7

1. It’s a liquid2. The pH is 123. It burns in air4. It tastes sweet5. It is green6. It weighs 5 lbs.7. It bubbles in acids

Chemical or Physical Property? #1-7

1. P2. P3. C4. P (?)5. P6. P7. C

Intensive vs. Extensive Properties?

1. Its mass is 50 g.2. It dissolves in oil.3. Its density is 1.5g/ml4. It is 6 inches long5. It conducts electricity6. It is acidic7. It is at room temperature.

Intensive vs. Extensive Properties?

1. E2. I3. I4. E5. I6. I7. E

Homogeneous or Heterogeneous?

1. Concrete2. Jello3. Muddy water4. Diamond5. Hair6. Children in a class7. Tossed salad8. milk

Homogeneous or Heterogeneous?

1. He2. Ho3. He4. Ho5. He6. He7. He8. Ho

Element, Compound or Mixture?

1. C2. E3. C4. M5. E6. C7. C8. M

9. M10. M

Name these elements

1. Sb2. As3. Ni4. Fe5. Zr6. Ra7. Au

8. Na9. Sr10. Ag11. Ba12. P13. F14. Mg

Name these elements

1. antimony2. arsenic3. nickel4. iron5. zirconium6. radium7. gold

8. sodium9. strontium10. silver11. barium12. phosphorus13. fluorine14. magnesium

Write symbols for these elements

1. Aluminum2. Tin3. Rubidium4. Argon5. Helium6. Neon7. uranium

8. lead9. potassium10. calcium11. zinc12. chlorine13. copper14. tungsten

Write symbols for these elements

8. Pb9. K10. Ca11. Zn12. Cl13. Cu14. W

1.Al2. Sn3. Rb4. Ar5. He6. Ne7. U

Solubility Song

• Sing• Make a ChartWhat is the chemistry

of a soluble ionic compound?

Solubility Quiz: Soluble or not?

1. Iron II hydroxide2. Potassium

phosphate3. Barium nitrate4. Strontium sulfate5. Calcium chloride6. Silver acetate

Solubility Quiz: Soluble or not?

1. not2. sol3. sol4. not5. sol6. not

Name two solutions with soluble salts that would combine to form the precipitates in #1, 4 and 6

Ions to Know

• Thoughts on patterns

What is the common ionic charge for each element?

1. calcium2. argon3. potassium4. nitrogen5. chlorine6. aluminum7. oxygen

What is the common ionic charge for each element?

1. 2+2. 03. 1+4. 3-5. 1-6. 3+7. 2-

Name these ions

• SO42-

• NO3-

• PO33-

• C2H3O2-

• NH4+

• S2O32-

• C2O42-

Name these ions

• Sulfate• Nitrate• Phosphate• acetate• ammonium• thiosulfate• oxalate

Write formulas for these ions

• Carbonate• Nitrite• sodium• iodite• sulfite• sulfide• bromate

Write formulas for these ions

• CO32-

• NO2-

• Na+

• IO2-

• SO32-

• S2-

• BrO3-

What is the charge on these ions?

• PO3

• N• AsO4

• ClO2

• S2O8

• NH4

• IO4

Something New…in Chapter 1• Antoine LaVoisier1743-1794

French Chemist

father of modern chemistry

At age 28 married 13-year-old Marie-Anne who translated from English for him and illustrated his books; she was well educated in chemistry herself. •burned P and S in air, and proved the products weighed more than the reactants but the weight gained was lost from the air. •Thus established the Law of Conservation of Mass.

•1778: demonstrated the "air" responsible for combustion; named this portion of air oxygen and the other part of air “azote” (Greek for no life).

•discovered that hydrogen combined with oxygen to produce water.

Antoinne

LaVoisier

•1787: invented the system of nomenclature still used today. •1789: published the first modern chemical textbook, with his theories:•a clear Law of Conservation of Mass•There is no such thing as phlogiston•a list of elements, including oxygen, nitrogen, hydrogen, phosphorus, mercury, zinc, and sulfur….. but also included light and caloric

Lavoisier: "I have tried...to arrive at the truth by linking up facts; to suppress as much as possible the use of reasoning, which is often an unreliable instrument which deceives us, in order to follow as much as possible the torch of observation and of experiment."

Phlogiston Theory

• Ancient Greeks thought there were four substances in the world: E,A, F, W

• In the 1600’s Johann Becher added to the list: Phlogiston is a 5th “element”… it’s in any substance that burns!

•LaVoisier

•worked as a tax collector

•beheaded during the French revolution for using public money to fund his research.

LaVoisier beheaded in France.

A Word about ENERGY

• What is it?• Law of Conservation

of Energy

Energy:part of chemical reactions

• Energy is required to break bonds.

• Energy is released as bonds form.

• This is true in all physical and chemical changes.

Reactions and Energy• Endothermic• Exothermic• More about NRG in later

chapters.

9. What is the % Ag in your alloy? Compare the percent (what is your error%) you got to the theoretical value for the % Ag in dimes made before 1950.Suggest an error to account for the difference. Choose one with the correct direction.

Results of the atomic theory• It yields

definitions of:• Elements• Compounds• Chemical

Reactions

Dalton’s Postulates lead to two laws:• Conservation of Mass• Multiple Proportions:If two

elements form more than one compound, the mass ratios of one of the elements in one compound to the same element in the other compound is always in a small whole number ratio.

• Think of benzene and methane

Mendeleev

• Born in Siberia 1834• Youngest of 14 children.• Hated everything in

school except science.• Father died when he was

2, mother favored him as a student and child, she died after he got admitted to university at age 15.

Mendeleev

• In 1855 was told he had two years to live, probably had tuberculosis.

• Worked as a professor of chemistry at St. Petersburg, Russia

• Organized known elements according to their properties and thus discovered the periodic law.

Mendeleev• Meyer also discovered a

periodic law, but Mendeleev published first.

• Was a talented and popular public speaker

• Married Feozva, had two children… they did not get along. Divorced her and married Anna, with whom he had four children.

• The Czar looked the other way on his “bigamy”

Mendeleev

• His periodic law was most accepted after it was shown his predictions of the existence of other elements were correct.

• Eka-silicon and two others were discovered.

• Died in 1907 at the age of 73.

Molecular vs. Empirical formula

• Molecular formula is the real formulafor a compound

• Empirical formula is the lowest ratio of elements in the compound.

• Example: • ethylene glycol is

C2H6O2

(molecular) • empirical formula is

CH3O (lowest ratio)

Structural Formula

• Arrangement of the atoms in a formula to show what shape, function or type of molecule it is.

• H-O-HCH3COOH

• Many organic molecules are frequently written structurally.

Organic Carbon Chain Classes

• Alkanes• Alkenes• Alkynes

The End of Chapter 1&2 notes

• Time really matters.• T-shirt? Submit ideas

Asap• Mole Day volunteers

needed: make a moleata, run a contest.

A.P. Chemistry: Chapter 3

• Calculations with Chemical Formulas and Equations

What is a Mole?