H Chem Semester 1 Notes 2013.pdf - Troy High School

Name_____________________________________Period__________Honors ChemistrySemester 1Ms. RegliTroy High School

Table of ContentsTitlePagePeriodic Table.............................................................................................3Equations Sheet.........................................................................................4Ion Sheet.....................................................................................................5Activity Series of Metals / Solubility Rules..............................................7Unit 1: Measurements (Chapter 1)...........................................................8Unit 2: Density/Atomic Theory/Nuclear (Ch. 1, 2, 23)............................20Unit 3: Nomenclature (Chapter 2)............................................................44Unit 4: The Mole (Chapter 3)....................................................................52Unit 5: Stoichiometry (Chapter 3, 4)........................................................66Unit 6: Atomic Theory II/Quantum (Chapter 7, 8)...................................82Unit 7: Bonding (Chapter 9, 10).............................................................106©RegliGower 2

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General Conversion FactorsLength Mass Volume Area1 inch = 2.54 cm 1 lb. = 16 oz. = 453.6 g 1 L = 1.057 qt. 1 acre = 43,560 ft 21 meter = 39.37 in 1 kg = 2.205 lb. 1 gal. = 4 qt. = 8 pt.1 mile = 1.609 km = 5280 ft 1 oz. = 28.35 g 1 pt. = 2 cups = 16 fl. oz. Energy1 foot = 30.48 cm 1 mL = 1 cm 3 1.60 x 10 -19 J = 1 eV1 cal = 4.184 JPressure1 atm = 14.7 psi = 760 mmHg = 760 torr = 101.3 kPaGeneral InformationD = m/V Density of water = 1 g/mL m electron = 9.11 x 10 -31 kg1 mol = 6.022 x 10 23 E = mc 2 1 mol (gas) = 22.4 L @ STPE a =⏐⏐O-A⏐⏐ E r = (E a / A) x 100Geometry EquationsVol. of cylinder = πr 2 h Vol. of sphere = (4/3)πr 3 Vol. of cube = s 3 c = 2πr = πdAtomic Theory EquationsE = hf = hc/λΔE = E f - E iE n = -R H /n 2λ = h/mvr n = n 2 (5.3 x 10 -11 m)Atomic Theory Constantsh = 6.626 x 10 -34 Jsc = 3.00 x 10 8 m/sR H = 2.18 x 10 -18 J = 13.6 eVGas Law EquationsGas Law ConstantsThermochemistry EquationsThermochem Constants for H 2 OQ = mcΔT Q = C cal ΔT m H c H (T H -T f ) = m c c c (T f -T c ) c ice = 2.06 J/g°CΔH rxn = ΣΔH° f products - ΣΔH° f reactants ΔG = ΔH - TΔS c water = 4.184 J/g°C = 1 cal/g°CΔH rxn = Σ E reactants – Σ E productsc steam = 2.02 J/g°CQ = ml f Q = ml v c = kP l f = 334 J/gl v = 2260 J/gSolutions EquationsM = mol / L m = mol / kg solvent ΔT f = K f mN = Equivalents / L % by mass = mass solute / mass solution x 100 ΔT b = K b mM 1 V 1 = M 2 V 2 X = mol solute / mol solute + mol solvent π = MRTP 1 = X 1 P° 1ΔP = X 2 P° 1Acids / Bases EquationsAcids / Bases ConstantsK w = [H + ][OH - ] pH = - log [H + ] pOH = - log [OH - ] K w (25˚C) = 1.0 x 10 -14pH + pOH = 14 N a V a = N b V b K a ·K b = K wRedox EquationsRedox ConstantsΔG = -nFE cell E° cell = E° red + E° ox 1 e- = 1.6 x 10 -19 C1 mol e- = 96,487 C = 1 F©RegliGower 4

Inorganic NomenclatureIon SheetDirections: Make flashcards of the following ions. The name should be written onone side, the chemical formula on the other.Positive Ions with a fixed charge:Positive Ions (Cations)Group IA ions: Group IIA ions: Group IIIA ions: Misc. cations:Hydrogen H + Beryllium Be 2+ Aluminum Al 3+ Silver Ag +Magnesium Mg 2+ Boron B 3+Lithium Li + Calcium Ca 2+ +Ammonium NH 4Sodium Na + Strontium Sr 2+Potassium K + Barium Ba 2+ Zinc Zn 2+Rubidium Rb +Cesium Cs + Cadmium Cd 2+Positive Ions with multiple charges: (one card for each ion)IUPAC Name Common Name IonIron (II) Ferrous Fe 2+Iron (III) Ferric Fe 3+Copper (I) Cuprous Cu +Copper (II) Cupric Cu 2+Cobalt (II) Cobaltous Co 2+Cobalt (III) Cobaltic Co 3+Mercury (I)** Mercurous2+Hg 2Mercury (II) Mercuric Hg 2+Manganese (II) Manganous Mn 2+Manganese (III) Manganic Mn 3+Tin (II) Stannous Sn 2+Tin (IV) Stannic Sn 4+Lead (II) Plumbous Pb 2+Lead (IV) Plumbic Pb 4+**Mercury (I) exists as a pair or dimmer as shown. No single Hg+ ions are known to exist.©RegliGower 5

Negative Ions (Anions)Monatomic anions:Group IVA ions: Group VA ions: Group VIA ions: Group VIIA ions:Carbide C 4- Nitride N 3- Oxide O 2- Fluoride F -Silicide Si 4- Phosphide P 3- Sulfide S 2- Chloride Cl -Bromide Br -Group IA ions: Iodide I -Hydride H -Polyatomic anions:Oxyhalogens:Hypochlorite ClO - Hypobromite BrO - Hypoiodite IO -Chlorite-ClO 2 Bromite-BrO 2 Iodite-IO 2Chlorate-ClO 3 Bromate-BrO 3 Iodate-IO 3Perchlorate-ClO 4 Perbromate-BrO 4 Periodate-IO 4Miscellaneous polyatomic ions: (one card for each ion)Peroxide O 22-Cyanide CN -Hydroxide OH - Thiocyanate SCN -Carbonate CO 32-Hydrogen carbonate / bicarbonate HCO 3-Monohydrogen phosphate HPO 42-Dihydrogen phosphate H 2 PO 4-Arsenite AsO 33-Arsenate AsO 43-Chromate CrO 42-Dichromate Cr 2 O 72-Phosphite PO 33-Phosphate PO 43-Acetate C 2 H 3 O 2-Binary Covalent Greek PrefixesPrefix Number Prefix NumberMono- 1 Hexa- 6Di- 2 Hepta- 7Tri- 3 Octa- 8Tetra- 4 Nona- 9Penta- 5 Deca- 10Cyanate OCN -Nitrite NO 2-Nitrate NO 3-Sulfite SO 32-Sulfate SO 42-Thiosulfate S 2 O 32-Permanganate MnO 4-Oxalate C 2 O 42-Phthalate C 8 H 4 O 42-Silicate SiO 44-Borate BO 33-©RegliGower 6

(Fig. 4.15) Activity Series of MetalsLiKBaCaNa React with cold water to produce H 2MgAlZnCrFeCd React with steam to produce H 2CoNiSnPb React with acids to produce H 2HCuAgHgPtAuDo NOT react with water or acidsSolubility Rules:1. ALL alkali metal compounds are soluble.2. ALL NH 4 + compounds are soluble.3. ALL compounds containing NO 3 - , ClO 3 - , and ClO 4 - are souble.4. Most OH - compounds are INSOLUBLE. Alkali metal hydroxides and Ba(OH) 2 areexceptions.5. Most Cl - , Br - , and I - compounds are soluble. The exceptions are: Ag + , Hg 2 2+ , Pb 2+ .6. All CO 3 2- , PO 4 3- , and S 2- compounds are INSOLUBLE. Exceptions are alkali metaland ammonium compounds.7. Most SO 4 2- compounds are soluble. Exceptions are Ba 2+ , Hg 2+ , and Pb 2+ .©RegliGower 7

I. Algebra In ChemistryUnit 1: MeasurementsA. Many relationships used in chemistry require working with algebraic equations. The equation orequations you start with are the _______________ equations while the equation that you solve for iscalled the ________________ equation. The working equation is formed by ____________ thedesired variable on one side of the equation. The basic rule for doing this is …. “What you do to oneside of the equation you must do to the _____________”Examples:1. Find x for the following:(a) x + y = z (b) x – y = z (c) 2x + y = z (d) 3x + 2y = z(e) x 2 + y = z (f) x 3 + 3y = z (g) 2x – 4 = 142. D = m/V is the given equation for the density of an object.(a) If the mass and density of an object were given, what would be the equation used tofind the volume?(b) If the density and volume of an object were known, what equation could be used to find itsmass?3. E = mc 2 ; Solve for (a) m (b) c4. PV = nRT; Solve for (a) V (b) T5. The formula for the volume of a cylinder is V = πr 2 h; Solve for (a) h & (b) r6. The formula for the volume of a sphere is V = 4/3 πr 3 . What equation could be used to solve forthe radius of a spherical container of known volume?7. Solve for V 2 :©RegliGower 8

B. It is often required to substitute one or more equations into another equation to form theworking equation.1. For example: If the radius, mass, and density of a cylinder are known, the equation for the heightof the cylinder can be formed by combining and rearranging the density and volume formulae.2. If the length, width, height, and density of a block of metal are known, show the equation that canbe used to calculate its mass.3. If the radius and mass of a metal ball are known, show the equation that can be used to determineits density.4. Derive an equation that can be used to determine the diameter of a crystal ball with a knownmass and density.5. If the radius and height of a cylinder are known, determine the formula that can be used to find theside of a cube that has twice the volume as the cylinder.C. Algebra can also be used to solve problems ……1. A chemist has 300 g of 60.0 % acid. How much water should be added so that it is only 45.0 %acid?©RegliGower 9

II. Useful Geometry FormulasArea of a square Area of a rectangle Area of a circleA = A = A =Circumference of a circlec = c =Density of an Object:D =Volume of a squareVolume of a rectangular solidV = V =Volume of a sphereV =Volume of a cylinderV =©RegliGower 10

III. Significant Figures:w = 6.87 cmA. Example:h = 0.05 cml = 17.9 cm1. Calculate the area of the dark rectangle.2. Calculate the volume of the object.3. Calculate the sum of the length, width, and height.4. What is the length of each segment?10 cm11 cmA = ____________________B = ____________________A B C DC = ____________________D = ____________________©RegliGower 11

B. Introduction:When making measurements or doing calculation, you should not keep more digits in a number thanis _________________. These rules of significant figures will show you how to determine the correctnumber of digits.C. What is a significant figure?Significant figures in a measurement are all values (digits) known precisely, plus ______ digit that isestimated.Example: Make the measurement with the correct significant figures.a. b. c. d. e. f. g. h.9 cm 10 cm 9 cm 10 cm 9 cm 10 cm 0 cm 1 cma. _____________c. _____________e. _____________g. _____________b. _____________d. _____________f. _____________h. _____________D. How do you determine sig figs in a measurement that has already been recorded?Sig Figs: The Rules1. Every nonzero digit in a recorded measurement is significant.Examples: 47,357 5 sig figs25 _______2. Zeroes appearing between nonzero digits are significant. (Sandwich rule)Examples: 1.007 4 sig figs305 _______3. Zeroes appearing in front of all nonzero digits are NOT significant. They are acting asplace holders (leading zeroes).Examples: 0.00238 3 sig figs0.98 _______0.000006 _______4. Zeroes at the end of a number AND to the right of a decimal point are significant.Examples: 426.00 5 sig figs2.060 _______0.8080 _______5. Zeroes at the end of a measurement where there is no decimal point are ambiguous. Toclearly show the correct number of sig figs, these measurements should be written in scientificnotation.Examples: 120 2-3 sig figs3000 1-4 sig figs1,000,000 _______Examples: Write the number 100,000 with (a) 1 sig fig, (b) 3 sig figs, and (c) 5 sig figs.©RegliGower 12

E. Practice:1. Determine the number of significant figures in each of the following measurements.(a) 54320.0 (b) 0.004550 (c) 151309 (d) 10.54(e) 5.20 x 10 5 (f) 15,000 (g) 10.04 (h) 0.07502. When completing calculations, it is often necessary to round the final answer to a particularnumber of significant figures (round up for 5 and above; keep digits the same for 4 and below).Round the above measurements to 2 significant figures.Example: 0.0753 = 0.075107.0 = ________________3. Determine the number of sig figs for each measurement. Round the measurements to 2 sigfigs. If original measurement only contains 1 or 2 sig figs, leave the second line blank.# sig figs Rounded Answer # sig figs Rounded Answer1. 0.0037 _______ ______________ 11. 14.04 _______ ______________2. 134.1 _______ ______________ 12. 5.401 _______ ______________3. 1,000,000 _______ ______________ 13. 1340 _______ ______________4. 5.730 x 10 2 ______ ______________ 14. 0.00566 _______ ______________5. 410.50 _______ ______________ 15. 0.8120 _______ ______________6. 79500 _______ ______________ 16. 18.009 _______ ______________7. 3071.04 _______ ______________ 17. 100.5 _______ ______________8. 4.08 x 10 -6 _______ ______________ 18. 3008 _______ ______________9. 0.998 _______ ______________ 19. 112040.0_______ ______________10. 1.570 _______ ______________ 20. 43.05 _______ ______________===========================================================================4. Rules for Significant Figures in Calculationsa. Multiplication or Division: The number of sig figs in the result is the same as the numberin the least precise (least sig figs) measurement.Example: 4.56 m x 1.4 m = 6.38 m 2 (Rounds to TWO sig figs) = 6.4 m 2(a) 17.24 x 0.52 (b) 118.24 x 3.5 (c) 1.007 x 14.407.58b. Addition and Subtraction: The answer should be rounded off to retain digits only as faras the first column containing estimated digits (remember that the last digit is estimated).Example: 12.11 m + 8.0 m + 1.013 m = 21.123 m (Rounds to ONE place after the decimal) = 21.1 m(a) 21 cm – 18.3 cm = (b) 10000.00 mm + 25.116 mm =©RegliGower 13

IV. Scientific Notation (Exponential Notation)A. Chemistry examples:1. Avogadro’s Number2. Mass of an electronB. Technique to change from positional notation to scientific notation:1. Leave ________ number to the ____________ of the decimal.2. When the decimal is moved to the _________, the exponent is_______________.3. When the decimal is moved to the _________, the exponent is ______________.4. Number must contain the same number of __________________ as the originalvalue.C. Convert the following to scientific notation:1. 135000 (3 sig figs) ____________________2. 0.005500 ____________________3. 120,000,000,000 (2 sig figs) ____________________4. 0.00000004441 ____________________D. Use of calculator with scientific notation:Example: 1.61 x 10 -19 What you see on your calculatorStep 1: Enter the numberStep 2: Press the Exponent buttonStep 3: Enter the exponentStep 4: If negative exponent, useE. Exponent problems (Use correct sig figs!)1. (3.5 x 10 -2 )(4.44 x 10 12 )(1.280 x 10 -22 )orkey2. (1.76 x 10 -2 )(4.2 x 10 -4 )(6.99 x 10 6 )(8 x 10 14 )Raising to a powerStep 1: Enter numberStep 2: PressStep 3: Enter powerStep 4: PressTaking a rootStep 1: Enter numberStep 2: PressStep 3: Enter rootStep 4: PressExample: (a) (14.5) 6 Example: (a)(b) (1.72 x 10 5 ) 4(b)©RegliGower 14

V. Metric SystemA. Based on powers of 10Ex. 1 m = _______ dm = _______ cm = _______ mmB. Uses “________________” and “____________________.”1.2.3.4.5.MeasurementMetric Base UnitC. Metric Prefixes: MEMORIZE this table!1.2.3.4.5.6.7.Base Unit8.9.10.11.12.13.Prefix Symbol Multiplier/FactorD. Examples: Multiplier ALWAYS goes with the _______________________.1 Mm = _______ m1 µg = _______ g1 Ts = _______ s1 pm = _______m©RegliGower 15

E. Converting within metric system using dimensional analysis:1. Convert to base unit by canceling units (Top unit cancels with __________ unit).2. Place the multiplier with the __________________.3. Place a _______ in front of the unit with ________________.*** 4. To enter multiplier into the calculator, use a _____ before the exponent key.Example: 10 -6F. Metric dimensional analysis examples:1. Convert 3.6 nm to m.2. Convert 55.6 g to Tg3. Convert 575 cm to Mm.4. Convert 0.456 dag to pg.5. Convert 78.5 km to µm.6. Convert 0.000590 mL to GL.Metric / English Conversion Factors (given on test):Length Mass Volume Time1 inch = 2.54 cm 1 lb. = 16 oz. = 256 drams 1 L = 1.057 qt. 1 fortnight = 2 weeks1 meter = 39.37 in 1 kg = 2.205 lb. 1 gal. = 4 qt. = 8 pt.1 mile = 1.609 km 1 lb = 453.6 g 1 pt. = 2 cups Density of water1 furlong = 220 yd. 1 oz. = 28.35 g 1 mL = 1 cm 3 1 g = 1 mL1 pt. = 16 fl. oz. 1 g = 1 cm 3VI. Conversion Factors:A. Whenever two measurements are equal, or __________________, a ratio of these twomeasurements will equal ________.Example: ______ ft. = ______ in. Can be written as the following ratios:©RegliGower 16

B. Conversion factor: ratio of ____________________ measurements.C. Write conversion factors for the following pairs of units:1. miles and feet2. days and year3. yard and feetD. Assume all conversion factors are _________________ significant. (Use initial numberto determine sig figs).VII. Dimensional Analysis IUnits (______________________) are used to solve a problem.Examples:A. The average human brain weighs 8.13 lb. What is the mass in ng?B. How many microseconds in 8.37 years? Write answer in scientific notation.C. A container contains 15 kL. Convert this to cm 3 .D. Apollo 13 re-entered the Earth’s atmosphere at a speed of 32,805 ft/s. What was thespeed in miles per hour (mph)?E. An arrow moves towards you at 235 m/s. How many miles could the arrow move in oneday? (Assume the arrow never falls to the Earth).©RegliGower 17

F. (a) Determine the number of cm 3 in a 20.0 fl. oz. bottle of Coke. (b) What is the mass ofthe Coke in pounds, assuming that it is the density of water (1 g / mL)?G. The speed of light is 3.00 x 10 8 m/s. How many miles does light travel per year?H. Carl Lewis set the world record for the 100.0 m dash on August 25, 1991 in the finals of theWorld Track Championships with a time of 9.86 seconds. What was his average speed inmiles per hour?VIII. Dimensional Analysis II: Square and cubic unitsA. Convert 3.7 ft 2 to in 2 .B. The engine in a Jeep Cherokee is 4.0 L. Calculate the engine volume in (a) in 3 and(b) ft 3 .C. The density of gold is 19.3 g/mL. Calculate the density of gold in (a) lb/ft 3 , (b) kg/m 3 .©RegliGower 18

D. A spherical container with a diameter of 2.85 dam is filled with water. (a) Determine thevolume of the sphere in cm 3 . (b) Determine the mass of the water in kilograms.E. The dimensions of a swimming pool are 13.5 ft. x 22 m x 225 cm. (a) Determine thevolume of the pool in m 3 . (b) Determine the mass of the water in pounds.F. A 12.0 fl. oz. soda spilled onto the floor into a cylindrical puddle with a 15.4 inch diameter.Calculate the depth (height) of the puddle in µm.G. The volume of a red blood cell is 90.0 µm 3 . What is its diameter in mm? Assume it isspherical.H. The lid of a soup can is 5.40 cm across and the can is 12.2 cm high. What is the volume ofthe can in fluid ounces?©RegliGower 19

I. DensityUnit 2: Density, Atomic Theory & NuclearA. Density is defined as _______________ per unit _______________ and has the symbol______ (or the Greek letter “rho” = _______).B. The formula for density is D = _______ Density = _________________C. Density can have many different units but the most commonly used units are as follows:English systemfor solids & liquidsfor gasesMetric systemD. Converting from between units is easy if you use ____________________ analysis.Example1: Gold has a density of 19.3 g/cm 3 . What is its density in….(a) kg/m 3(b) g/mLII. Specific Gravity(c) lb/ft 3A. Specific gravity is a measure of the mass of an object compared to the mass of an equalvolume of ___________________.B. It has the formula: sp. gr. =_____________ = mass of objectmass of equal volume of waterC. Since the density of water is 1.00 g/mL, we see that specific gravity is equal numerically tothe ______________ but has ______ __________!!!Example 2: Gold has a density of 19.3 g/mL. Calculate the specific gravity of gold.D. The specific gravity has the same value in any _____________ of units, since it expressesthe quotient of the mass of the substance divided by the mass of an equal volume of_____________. It is expressed by a pure number ___________________________.Specific gravity is usually used with _________________.E. The water standard is sometimes taken at the temperature where it has maximum density.At typical laboratory values (0 o C to 30 o C) the density of water does not vary much and therounded value ________________________ can be used.©RegliGower 20

F. Steps to calculate the density of objects:1. Find the mass (balance).2. Find the volume (dimensions of displacement).3. Use the equation D = m / VExample 3: Geometric shapesStep 1: Measure the mass of the object.Step 2: Measure the dimensions of the object.Step 3: Calculate density (D = m / V).mass =_____________ gl =______cmw =______cmh =______cmOther formulae: V cyl = πr 2 h V sphere = 4/3 πr 3 V cube = s 3Example 4: Irregular shapesStep 1: Measure mass of object“Water displacement”Step 2: Put water in a graduated cylinder andrecord initial volume.Step 3: Carefully put object in and recordthe final volume.Step 4: Calculate density.ObjectInitialmeasurementFinalmeasurementmass =________ gV i = ________mLV f = ________mL©RegliGower 21

III. Density Calculations:1. An object has a mass of 1.00 kg and a density of 4.00 g/mL. What is its volume in liters?2. What is the mass of a 500 cm 3 object that has a specific gravity of 0.800?3. If an object has a density of 7.2 g/mL, what is its sp. gr.?4. What is the density (g/cm 3 ) of solid 75.84 gram cylinder that is 4.00 cm tall and 2.00 cmwide?5. Calculate the diameter of a 16.0 pound shot put. A shot put is a solid metal ball thrown bytrack athletes and it has a sp. gr. of 7.78. Calculate the diameter in inches.©RegliGower 22

IV. Accuracy and Precision:Methods of expressing laboratory error:In many of the laboratories done in a science class, you ultimately end up measuring ordetermining something that is already ______________. For example, a student may be givenan aluminum cylinder and be asked to measure it so that they can determine its density. Whydon’t they just look up the “accepted” value of the density of aluminum in the book? Well, thepurpose of most introductory laboratories is to teach you the proper experimental techniques,not to discover new information.Once a student has obtained specific values, how do they know how close their measurementis to the true (accepted) value that they are trying to obtain? This is done most simply bydetermining ________________ and ________________ (percent) error.A. Introduction: Chemistry is a quantitative science…it uses lots of numbers!! Thesenumbers come from experimental measurements and each measurement has some degree ofuncertainty in it.1. Reasons for UNCERTAINTY in measurements:(a) ___________________________ = Construction of the device.(b) ___________________________ = Incorrect usage of the device.(c) ___________________________ = Temperature, pressure, etc.2. Methods of expressing uncertainty:(a) +/- notation(b) Accuracy and Precision: “The Dartboard Analogy”accurate but precise but accurate not accuratenot precise not accurate and precise not precise3. Explain how a student can get precise inaccurate measurements.B. Accuracy: The closeness of a measurement (calculation) to the ________ (True orAccepted) value. It is expressed in terms of _____________.O = _________________, experimental valueA = _________________ (True) value1. Absolute error (E a )a. E a =2. Relative error (E r ) (Percent % error)a. E r =©RegliGower 23

Example # 1: A student measures the length of an object to be 7.45 cm. If the actual length is8.000 cm, calculate the absolute and relative error of his measurement.Question: Explain why relative error is more useful than absolute error? Give an example.Example # 2: Calculate the density of the aluminum block.l =______cmh =______cmmass =_____________w =______cmIf the accepted value for the density of aluminum is 2.70 g/cm 3 , calculate the absolute andrelative error of this value.Example # 3: Calculate the density of this cylinder:m = _______________ gh = _______________ cmd = _______________ cmIf the accepted value of this cylinder is __________ g/cm 3 , calculate the absolute and relativeerror of this value.Example # 4: Calculate the density of this sphere:m = _______________ gd = _______________ cmIf the accepted value of this cylinder is __________ g/cm 3 , calculate the absolute and relativeerror of this value.©RegliGower 24

V. Elements:A. A substance in which all of the ____________ have the same number of protons in thenucleus.B. ______________________ (a Swedish chemist): generally given credit for creating themodern symbols of elements.C. Atomic number: Number of _________________.D. Approximately __________.E. 90 elements occur _________________ (elements 43 & 61 are man-made)F. 93 and beyond are ___________________ transuranium elements.Examples:G. Names & symbols:1. Carbon – _____; Calcium – _____; Chlorine – ______2. 104 and beyond Unnilquadium (Unq) Un = _____ nil = _____ quad = _____3. Elements named after _____________ (old).a. Sodium __________ _______________________b. Gold __________ _______________________c. Silver __________ _______________________d. Potassium __________ _______________________e. Lead __________ _______________________f. Antimony __________ _______________________g. Iron __________ _______________________h. Tungsten __________ _______________________i. Tin __________ _______________________j. Copper __________ _______________________k. Mercury __________ _______________________H. States of matter1. solid (s)2. ________________3. gas (g)I. ___________________: 2 or more elements or compounds physically joined.1. Alloy: Brass: _______________Bronze:_______________J. ___________________: 2 or more elements chemically bonded to one another.Ex. : O :HH©RegliGower 25

K. Elements: ______________________________________ into simpler substances bychemical or physical means (Excluding nuclear processes).L. Periodic Table:1. Vertical columns – ____________________________________ (Chemically similar)2. Rows - ______________________3. Important Group Names:a. Group IA(1): ___________________________ (except Hydrogen)b. Group IIA(2): alkaline earth metalsc. Group VIIA(17): ____________________________d. Group VIIIA(18): noble gasese. _______________ ; Metalloids; ________________f. Most elements are _________________ ( sulfur, carbon, and sodium)g. Some are gases, & a few are liquids at room temperature (i.e. Br & Hg)VI. Properties of MatterA. Substances:1. Substances can be either pure ______________ or _______________.2. Substances are ________________ by enumerating their physical and chemicalproperties.3. All specimens of a given substance will have the ____________ chemical andphysical properties.4. Examples of substances:elementscompoundsB. Mixtures:1. Mixtures: Two or more pure _________________.a. Homogeneous mixtures __________________________________________b. Heterogeneous mixtures__________________________________________2. Examples of mixtures:HomogeneousHeterogeneous©RegliGower 26

C. Physical Properties of substances: Include those features, which definitely distinguishone substance from another.1. Density- ____________________________2. Specific gravity- ______________________________3. Hardness- Ability to resist scratching. MOH hardness scale is used as a basis ofcomparing the hardness of substances.a. Low numbers = relatively ________ substances.b. High numbers = relatively ________ substances.MOH scale: Talc 1Gypsum 2Calcite 3Fluorite 4Apatite 5Feldspar 6Quartz 7Topaz 8Corundum 9___________ 104. Odora.Good smells: ___________________________________________________b. Bad smells: ____________________________________________________5. Color- ______________________________________________________________6. Taste- __________________________________-___________________________7. Solubility in solventsExample:a. ____________ soluble vitaminsb. ____________ soluble vitamins8. Physical state: _______________________________________________________M.P. _________________ F.P. ___________________B.P. ___________________9. Properties of metalsa.Malleability ____________________________________________________b. Ductility _______________________________________________________c. Conduction of heat ______________________________________________10. Accidental physical properties: Not used to _________________ a substance.Examples: _________________________________©RegliGower 27

D. Chemical Properties: Describe the ability of a substance to form ___________________under given conditions.1. Chemical change: A change from one substance into another.2. Chemical properties may be considered to be a listing of all the chemical____________ of a substance.E. Changes in Matter: Matter undergoes ______________ and ______________ changes.1. Energy may be defined as the ability to ___________________.2. Forms of energy: _____________________________________________________3. Matter always possesses ______________ in one form or another.F. Physical Change: The composition of a __________________ is not ________________and the substance _________________its own identity.Examples:(1) _______________________________(2) _______________________________(3) _______________________________(4) _______________________________(5) _______________________________(6) _______________________________G. Chemical Change: The substance loses its _______________, and the new substanceformed has new _______________ and _______________ properties.Example reaction:Evidence of chemical changes:(1) ___________________ (2) ___________________(3) ___________________ (4) ___________________Examples:(1) ______________________________________(2) ______________________________________(3) ______________________________________(4) ______________________________________(5) ______________________________________(6) ______________________________________©RegliGower 28

5. Thomson Model of the Atom Theorya. _______ Thomson discovered that atoms are made of _______________, inother words they are made of ___________ things.b. J.J. did experiments with __________________________ (CRT) and hefound that:i. Cathode rays are _____ particles that he called _______________.ii. This showed that atoms are not _____________________!iii. He determined the _______________________ ratio of the electron.(_____________________)iv. He knew atoms were _______________ so he proposed a model ofthe atom called ____________________. The plums were _________.v. Robert A. _______________ determined the charge of the _____ withhis oil _____________ experiment.6. Rutherford Model of the Atom Theorya. The discovery of _____________ led to further advances in the atomic theory.b. New Zealand physicist Ernest B. __________________ and his associates(Geiger & Marsden) used radioactive ______ particles to probe the _________.c. He discovered the ________________________ with the __________scattering experiment.d. Diagram of the experiment:(a) Experimental setupfor Rutherford’sexperiment: α particlesemitted by radiumstrike metallic foil andsome reboundbackward;(b) backward reboundof α particles explainedas repulsion fromheavy positivelycharged nucleus.©RegliGower 30

II. Atomic Structure:e. His calculations regarding the deflected particles indicated that atoms have a_________, ______________, _______________________.f. If the atom were a ________ in diameter, the nucleus would be the size of a___________…yet the nucleus contains virtually all of the atom’s __________.In other words most of the atom is made up of _________ _________.g. He proposed the still popular (yet wrong) planetary model:i. electrons _________ the positive _____________.ii. like the ______________ around the _________.h. Problems with a planetary atomic model; It could not explain:i. Electron ___________: classical physics theory says that a chargedparticle (like an electron) moving in a circular orbit would _________energy and slow down, eventually collapsing out of its _________ and___________ into the_________.ii. Periodic _____________________ behavior.iii. Atomic ___________ spectra.7. The ___________ model of the atom: proposed by Danish physicist,_________________, would first explain some of these problems.A. 3 particles1. Proton (p)a. located in the _____________b. unit charge = _______c. mass = 1.67265 x 10 -24 gd. relative mass = _____ (relative to the other particles)e. charge = + 1.6022 x 10 -19 C (Coulomb)f. the mass is 1,836 x the mass of a _____________.g. discovered by ___________________h. made of _____ quarks:u ud(Proton)2. Electron (e)a. located outside the nucleus in ____________ levels (shells).b. unit charge = _____c. mass = 9.11 x 10 -28 g (9.10953 x 10 -28 g)d. relative mass = ______ (tiny compared to n & p)e. charge = −1.6002 x 10 -19 C (Coulomb)f. the mass is 1/1,836 the mass of a ___________.g. discovered by ______________.©RegliGower 31

3. Neutron (n)a. located in the _____________.b. unit charge = _______ “neutral”c. mass = 1.67495 x 10 -24 g slightly more massive than a proton.d. relative mass = ____e. discovered by ________________ in 1932 (Why so late? ________________)f. made of ____ quarks:u d(Neutron)dB. Nucleus1. The central ________ of the atom.2. Contains the neutrons and protons:____________ = a particle in the nucleus (n or p)3. _________ = the nucleus of an atom.4. Contains almost all the _________ of the atom.5. Has a __________ volume compared to the _________ atomic volume.Ping pong ball in the _______________.6. Density of nucleus = _________________ g/cm 3 .7. Radionuclide: an unstable __________. Why are some nuclei stable and othersare unstable? ________8. Nuclear stability is due to:a. Nuclear binding _____________________∗Holds protons and neutrons together.b. n/p ratio = stable atoms have a favorable n/p ratio.Examples:__________________________________________________________________________________C. Atomic Size1. Atomic diameters = ____________________________________1 Angstrom = ___________; 1 picometer = __________2. Diameter of atomic nuclei = _________________________3. How many carbon atoms are there in a pencil line 1.00 inch long?(radius = 0.77 Angstroms; assume 1 atom wide)©RegliGower 32

III. Mass Relationships in atomsA. Atomic Number (Z)1. Equal to the number of ___________.a. Equals the number of ________________ in a ________________________.i. Example: (a) Oxygen (b) Oxide ion (ions: charged particles)b. Each type of ______________ has a specific number of protons…thisdetermines the element’s _________.B. Mass Number (A)1. Equals the number of _____________ + the number of ______________.a. It is the number of ____________.2. # of neutrons = ________________________________________________3. Correct notation:C. Isotopes: Atoms with the same _________ number but different ___________numbers.1. Same number of ___________2. Different number of _____________.3. Isotopes are named by their ________________.a. Carbon-14 b. Uranium-2384. Example: The three isotopes of Hydrogen5. Example: The three oxygen isotopes6. Ions: _____________ atoms (or groups of atoms) that have lost or gained electrons.Examples:IV. Nuclear Chemistry: Chemistry of the nucleus (______ + ______).A. Radioactivity: The ___________________ emission of _______________ or EMR(_____________________) from the nucleus.1. Henri Becquerel: Discovered radioactivity (1896) using _____________________and uranium ore.2. Types of Radioactivity:a. Alpha (_____): Nucleus of a helium atom. (_________)b. Beta (_____): High speed electron emission from the __________. (______)c. Gamma ray (______): Photon of high energy light. (_________)©RegliGower 33

3. Penetrating power of radiation:a. Alpha can go through _____________.b. Beta particles can go through _______________________.c. Gamma rays can go through ________________________.________>________>________B. Nuclear equations: Must obey Law of Conservation of __________ (top line) and Law ofConservation of _________ (bottom line).1. Alpha emission (__________)2. Beta emission (________or________)a. Electron is formed in the _______________.b. Nucleus has too many __________________.c. Neutron spontaneously becomes a ______________, which causes a highenergy electron to be ejected from the nucleus.3. Positron emission: (________or________)a. Positron is formed in the _______________.b. Nucleus has too many ___________________.c. Proton spontaneously becomes a ________________, which causes a highenergy positron to be ejected from the nucleus.©RegliGower 34

4. Gamma emission (_________)a. Nucleus is in an _______________ energy state (excess energy from anotherdecay).b. As nucleus loses energy, a _________________ is emitted.c. Asterisk (*) is used to symbolize __________________________.5. Electron Capturea. An electron from the innermost energy level “falls” into the nucleus.b. Electron capture is more common with __________ nuclei.c. The product is the same as that of ______________ emission.6. General examples:a. β, β decayb. Uranium decay series: Radon-222 - α,α,β,β,α decayC. Decay Series: A radioactive decay often results in a ________________ nucleus that isalso radioactive. A radioactive _____________ _____________ refers to successivedecays, which starts with one parent isotope and proceeds through a number of daughterisotopes. The series ends when a stable, _____________________ isotope is produced.©RegliGower 35

D. Nuclear Stability1. Belt of Stability (FIGURE 1)Figure 1: Stableand unstablenuclides2. Binding Energy per Nucleon versus Atomic Mass (FIGURE 2)Figure 2: A plot ofnuclear bindingenergy per nucleonversus massnumber. Nuclei withlarge bindingenergies per nucleiare the most stable.3. Why are some nuclei stable and others are not?Rule 1: The greater the binding energy per nucleon (energy holding the nucleustogether), the greater the stability. (See Figure 2).Which isotope is the most stable, according to Fig. 2? ___________©RegliGower 36

Rule 2: Nuclei of low atomic numbers with a 1:1 ratio of neutrons to protons arevery stable. (See Fig. 1) Example: Carbon-12 Helium-4a. Isotopes decay with ________.b. Isotopes decay with ________.c. Isotopes decay with ________.Refer to Figure 1** Radioactive isotopes decay until they reach the “Belt of Stability.”Rule 3: The most stable nuclei tend to contain an ___________ number of bothprotons and neutrons.Example: Iron-56 Oxygen-16E. How is Binding Energy determined?1. Nuclear binding energy: The energy required to break up a nucleus into itscomponent ____________ and ___________ (nucleons).2. Binding energy comes from the mass defect of the nucleus.3. Mass Defect: The total mass of the stable nucleus _________ the sum of themasses of the nucleons. The “missing” mass has converted into energy!(________________)Note: Whencalculating masses,include the mass ofthe electrons becausethe mass of the wholeatom includeselectrons.Masses of subatomic particlese -0.00054858 amup +1.007276 amun1.008665 amu(p + + e - ) 1.007825 amuNote: The whole atom isNOT the sum of its parts.MISSING MASSMass defect = Atomic mass of the isotope - Σ mass of subatomic particles.Relationship between mass and energy: E = mc 24. Information for conversions between mass and energy:1 g = 6.02 x 10 23 amu 1 Joule (J) = c = 3.00 x 10 8 m/sEx. 1: Determine the energy in Joules of a mass defect of 1.00 amu.©RegliGower 37

Ex. 2: a. Determine the mass defect of Fluorine-19.(The atomic mass is 18.9984amu).b. Calculate the binding energy in J (use conversion factor between mass andenergy).c. Calculate the binding energy per nucleon.5. As nuclear binding energy per nucleon increases, the stability of the nucleus_____________.6. The element with the greatest binding energy per nucleon is ______________.F. Half Life: The average time it takes for ____________ of the unstable atoms in a sampleto decay.1. Exponential decrease of atoms _________________________________________.2. How much of a 500 g sample of Uranium-235 would be left after five half-lives?©RegliGower 38

3. A 16.00 mg sample of Radon-222 decays to 0.250 mg after 24 hours. Determine thehalf-life.4. The half-life of molybdenum-99 is 67 hours. How much of a 1.000 mg sample is leftafter 335 hours?G. Transmutation Reactions1. Bombardment: Bombard a target nucleus with other _________________ or nuclei.Example particles: _____,_____,_____,_____,_____2. Particle + nucleus → usually more than one _______________ formed.3. Examples:(a)(b)(c)*Chadwick – Discovery of the neutron*Rutherford (1919)-First artificial transmutation(d)4. Shorthand notation: (d = deuterium = ; T = tritium = )(a)(b)(c)(d)H. Fusion: Combining smaller nuclei to form _______________ nuclei.1. The combining of small nuclei to form larger nuclei increases the ____________ pernucleon. Therefore, stability _______________.(See Figure 2: binding energy)2. Nuclei smaller than _____________ will give off energy when they fuse together.3. Example: Sun reactions (made up mostly of H and He; T in interior ≈15 million °C)__________________4. When products are formed in a fusion reaction, mass is __________. This mass__________ turns into _________________.5. Fusion reactions take place at very high _________________.©RegliGower 39

6. They are often called _____________________.7. Reasons for high temperature:a. Positive nuclei ___________ one another.b. High _____ or speed needed to fuse the elements to overcome the repulsion.c. High _______ = High temperature (Millions of degrees Celsius)I. Fission: Large nucleus is ________________ into smaller nuclei + one ormore_______________.1. Example:2. Many fission products are formed. More than 30 different elements have been foundamong the products.3. Diagrams:4. The ______________ formed in the initial stages can induce ______________ inother .5. Nuclear fission takes place for elements with very large nuclei. Elements larger than_______can undergo fission. (See Figure 2).6. Practical applications of fission: naturally occurring __________________ andartificial _______________________.7. The mass of the starting material is __________ the sum of the masses of theproducts. The mass defect (missing mass) turns into _____________.8. Nuclear chain reaction: a _____________________ sequence of nuclear fissionreactions.Critical mass: the minimum mass of fissionable material needed for a________________ chain reaction.Subcritical mass: Less than the minimum mass. Too many neutrons willescape (will not be absorbed by the fissionable material), so _______ chain rxnwill occur.©RegliGower 40

9. Applications:a. Atomic Bomb: Self-sustaining nuclear chain reaction.i. Must contain __________ critical mass of U-235 (or Pu-239).ii. Critical mass must be kept in ____________ places before detonation!iii. Neutrons originate from a source at the center of the bomb_________the rxn.b. Nuclear Reactors: Controlled fission reaction. (8% of electricity in the U.S.).i. Advantage: Source of energy other than fossil fuels. No globalwarming gases.ii. Disadvantages: Produces highly radioactive products.Example: Strontium-90 (radioactive) = chemically like___________. (concentrates in the _____________).Iodine-131 = ______________ cancer.(within 10 mi. radius of nuclear power plants → I pills)J. Nuclear reactors: A device in which the controlled ____________ of a certain substanceis used to produce new substances and ___________________!!!1. _________________: Using the neutrons released during fission to cause othernuclei to undergo fission.2. Self-sustaining ___________ reaction keeps on going once it has begun.3. The neutrons released in a typical fission can be used to cause other nuclei tofission. If enough _________ or _________ is available, it is possible to create a__________________________. The minimum amount of fissionable material requiredfor a self-sustaining chain reaction is known as the _______________ ___________.The critical mass depends on a number of factors, among the factors:a. ________________: slows neutrons down which increases the # of favorablecollisions. The most common are _________________ (Chernobyl) and__________________ (D 2 O)(Used in Canada)(No need for enriched uranium).b. _________________ ie. 235 U & 239 Puc. _______________________: The % of fission material is increased (bydiffusion or centrifugation).d. In _____________ _____________ the chain reaction is controlled and theenergy is released gradually. In an “atomic bomb”, the chain reaction isuncontrolled and the energy release occurs in a few moments of time.4. __________ Power Planta. _______________(Cadmium/Boron) = Absorb neutrons (Stops rxn—Applyingthe brakes).b. Many use Uranium – 235 for fuel; fuel rods contain fissionable material.©RegliGower 41

c. Where does the electricity come from?__________occurs in the nuclear reactor core → __________ is produced→ heat causes liquid water to turn to ____________ → steam causes the_____________to turn → ________________ is generated → steam isturned back into liquid water in the _______________ (which uses coldwater from a river, lake, or ocean)→ warm water (non-radioactive) is sentback into the river, lake or ocean.Demo: Turn turbine—electricity is generated.5. ______________ Reactor: One in which fissionable material is produced at agreater rate than it is consumed.a. Converts ______________________ into ___________________b. Where does the ___________ come from?6. Problems of nuclear power.a. Thermal pollutions: Hot water exits into the ocean.b. Leakage of fission fragments: Highly radioactive.Examples: Chernobyl (Ukraine), Three Mile Island (USA)c. Spent fuel disposal: Yucca Mtn, NV.d. Only lasts for 30 years. Why?i. The build up of __________________.ii. The structural materials ________________.©RegliGower 42

K. Radioactive Dating:1. If the half-life of a radioactive isotope is known, an estimate of the _________ of anobject often times can be made. For example, the ratio of _____________ to carbon-12 in a living object is relatively constant. However, a living object _____________________ carbon-14 when it dies. Thus, knowing the half-life of carbon-14(_____________) and the object’s 14 C/ 12 C ratio, an estimate of the object’s age can bemade. Willard __________ won the 1960 Nobel Prize in chemistry for development in1947 of radiocarbon dating.2. Formation of Carbon-14: _______________________________________________3. Beta minus decay of Carbon-14: ________________________________________4. Useful for dating objects about ___________________ years old.5. How about older things? ___________________6. Geologists determine the amount of __________ remaining in a rock relative to theamount of daughter nuclei present to estimate the passage of time since the rocksolidified from molten material.©RegliGower 43

I. Background:A. Periodic TableUnit 3: Inorganic Nomenclature1. Column: __________________ or __________________. (Similar properties)2. Row: __________________.3. _____________: Left of staircase (Majority of the elements).4. ___________________: Right of staircase.• Exception: ______5. ___________________: Touching the staircase.B. Ions (Charged atoms)• Exception: ______ (metal).1. _____________: positively charged (lost e-).2. _____________: negatively charged (gained e-).C. Trends in the periodic table1. Using the planetary model – (simplified model of atom)2. Energy levels can contain a maximum of:1 st energy level: _____2 nd energy level: _____3 rd energy level: _____ (_____)3. _____________ are the keys to chemical bondsEx.Column IA (1) (_______________)Column VIIIA (18) (__________)H (____ e-)He (____ e-)Li (____ e-)Ne (____ e-)Na (____ e-)Ar (____ e-)Similarities:____________________________________©RegliGower 44

4. Atoms can gain or lose _____ to achieve a full outer shell (more stable).5. Atoms will do what is ______________ (least energy) i.e. Oxygen has 6 valence e-:easier to ___________ than to ____________.Group # of valence e- Gain or lose e- ChargeIAIIAIIIAIVAVAVIAVIIAVIIIA6. Label your periodic table!II. Binary Ionic CompoundsA. Background info1. Metal / ________________ (_______________ is always written first!)2. One element ______________ and the other _________________.3. _________________ of e-4. Charged ions ______________ one another (opposites attract).5. Compound is ______________.B. Examples of compound formation of binary ionic compounds:Ex. Sodium & chlorineEx. Calcium & bromineEx. Lithium & oxygenEx. Aluminum & sulfur©RegliGower 45

C. Shortcut to determining binary ionic compounds (Criss-Cross method):1. ________________ from charge becomes the subscript.2. All ionic compounds are ______________ (no + or -).3. Subscripts are written in the _______________ possible ratio.4. The number “1” is never written (It is implied).5. Examples:Lithium & oxygenAluminum & oxygenCalcium & oxygenMagnesium & nitrogenD. Nomenclature of binary ionic compounds (bi = 2).1. ____________ is named first (name of atom).2. ____________ is named second, ending changed to ___________.3. If the metal (cation) can have multiple charges, the charge is written as a romannumeral (IUPAC) or as the common name. (Fe, Cu, Co, Hg, Mn, Sn, Pb)4. Formula to name:a. Li 2 O _______________________________________b. Al 2 O 3 _______________________________________c. CaO _______________________________________d. Mg 3 N 2 _______________________________________e. Fe 2 O 3 _______________________________________f. SnO 2 _______________________________________g. CuCl _______________________________________h. MnN _______________________________________©RegliGower 46

5. Name to formula:a. Beryllimum fluoride _______________ _______________b. Potassium bromide _______________ _______________c. Tin (II) oxide _______________ _______________d. Cobaltic sulfide _______________ _______________e. Strontium iodide _______________ _______________6. Polyatomic Ion: A group of atoms with a _________ charge.Ex. (1) CN - =(2) NH + 4 =(3) OH - =Examples:a. Polyatomic ions will _______________ stay together as a group.b. If there is more than one polyatomic ion, it must be placed in_____________.Ions Formula Name©RegliGower 47

III. Helpful Hints to Memorize Oxyanions:A. In learning the formulas and charges of common oxyanions, start with the –ate form. Fromthere it follows that: hypo______ite = 2 less oxygens______ite = 1 less oxygen______ateper_____ate = 1 more oxygen**ALL forms have the SAME charge!**A Guide to Determine Whether the –ate Formula is –XO 3 or –XO 4A Guide to Determine What the Charge of the Oxy-Anion is:©RegliGower 48

B. Examples:Borate = _____________Nitrate = _____________Nitrite = _____________Carbonate = _____________Chlorate = _____________Perhlorate = _____________C. “Thio-“ = Sulfur replaces an oxygen.Ex. Sulfate = _____________Ex. Cyanate = _____________Thiosulfate = _____________Thiocyanate = _____________IV. Ternary Ionic Compounds: (compounds containing _____ or more elements).1. Name the _______________2. Find the appropriate name of the ____________________________________.3. Formula to name:a. Li 2 SO 4 ____________________________b. Fe(NO 3 ) 3 ____________________________c. CdC 2 O 4 ____________________________d. Cu 3 AsO 3 ____________________________e. Mn 2 SiO 4 ____________________________f. (NH 4 ) 2 SO 4 ____________________________4. Name to formula:a. Potassium thiocyanate _______________ _________________________b. Aluminum permanganate _______________ _________________________c. Plumbic acetate _______________ _________________________©RegliGower 49

d. Cobalt (III) oxalate _______________ _________________________e. Sodium hypochlorite _______________ _________________________V. Nomenclature of HydratesA. Hydrate: Ionic compound with ______________ molecules stuck in the ___________lattice. The water is included in the ____________ and formula.1. ZnSO 4 · 7 H 2 O _______________________________________2. CaCO 3 ·3 H 2 O _______________________________________3. Cu 2 C 2 O 4 · 2H 2 O _______________________________________4. Calcium chloride pentahydrate _______________ _________________________5. Cupric acetate monohydrate _______________ _________________________VI. Binary Molecular CompoundsA. Molecular (___________________) compounds1. Non-metal to _______________________.• ____________________ of staircase including hydrogen2. ___________________ of electrons.Ex3. Non-metals can often combine in several different ways.Ex.B. Nomenclature of binary molecular compounds:1. Greek prefixes are used:mono = hexa =di = hepta =tri = octa =tetra = nona =penta = deca =2. The prefix “___________” is omitted for the 1 st element.Ex. CO = __________________________3. For oxides the ending “_____ or _____” is omitted.a. N 2 O = ____________________________b. N 2 O 3 = ____________________________c. N 2 O 4 = ____________________________d. NO = ____________________________e. NO 2 = ____________________________f. NO 5 = ____________________________©RegliGower 50

IonicCompoundCovalent1. 1.2. 2.3. 3.Ex.VII. Nomenclature of Acids1. P 2 O 5 ______________________________2. NCl 3 ______________________________A. Acids: Compounds that contain _________________ as the positive ion (H + ).B. Exceptions: ___________ (water) & ___________ (hydrogen peroxide).C. Binary Acids: Acids that ______ ________ contain oxygen.1. Use prefix “___________”2. Add stem or full name of _______________.3. Add suffix “_______”.4. Add the word ________________.Ex.HBr = __________________________________HCl = __________________________________HCN = _________________________________D. Ternary Acids: Contain ____ or more elements __________________ oxygen.1. Acids formed with anions that contain ____________ become __________ acids.HNO 3 (NO 3-= ______________) __________________________HClO 4 (ClO 4 - = ______________) __________________________H 2 SO 4 (SO 4 2- = ______________) __________________________H 3 PO 4 (PO 4 3- = ______________) __________________________2. Acids formed with anions that contain _________ become ___________ acids.HNO 2 (NO 2-=______________)__________________________HClO 2 (ClO 2 - =______________) __________________________H 2 SO 3 (SO 32-=______________) __________________________3. Name to formula:a. cyanic acid _______________ _________________________b. dichromic acid _______________ _________________________c. hypochlorous acid _______________ _________________________d. hydrosulfuric acid _______________ _________________________©RegliGower 51

I. Atomic Weight (Mass)Unit 4: The MoleA. Atomic Mass Scale:1. Reference Unit: The _______________________ atom.a. Symbolb. 1 ________ atom = ______________ (exactly)c. amu = ______________________________ u (_______________________)d. 1 amu = __________ the mass of a _______________ atom.B. Mass of atomic particles in amu1. electron 9.11 x 10 -31 kg or ____________________ amu.2. proton = __________________ amu.3. neutron = _________________ amu.C. Atomic mass (Weight): An atom’s mass (_____________) in atomic mass units (amu).i.e. Carbon-12: 12.00000 amu 12.011 amu (Chart value)1. The atomic weight found on the periodic table is the ______________ of thenaturally occurring ___________ of an element. (Look on chart)2. Calculation of Atomic Mass (Weight) you must be given …a. The _________ of each ____________.b. Abundance (%) in _________ (varies slightly by location)3. Calculate the atomic mass of carbon.Isotopes % abundance Atomic massCarbon-12 98.890 12.00000 amuCarbon-13 1.1100 13.00335 amu4. The mass and % abundance are determined with an instrument called a__________ _____________. The abundance can _________ depending on wherethe sample is from. This is why the sig. figs. vary on the chart.D. Other calculations:1. Lithium has two isotopes. If lithium-6 has a mass of 6.015123and 7.42 % occurrence, what is the % abundance and mass of lithium -7?2. What is the % abundance of the two isotopes of chlorine if the atomic weight ofchlorine on the periodic table is 35.453? Chlorine-35 has a mass of 34.968853 amuand Chlorine-37 has a mass of 36.965903 amu.©RegliGower 52

E. Mass Spectrometry: Is the process involving the production of_________________________ from a sample, and their resulting ________________ (and__________________) according to their __________ to __________ ratio. (m/q)1. 3 distinct operations involved in mass spectrometry.a. __________________ of _______________ why? (________________)b. _________________ (by ____________ or _____________ fields)c. __________________ of + ions (photographically or electronically)2. Key Ideas: __________ + ions are deflected ___________. Results give the_________ and % ________________ of each isotope. (See diagrams & book formore info).II. Molar Mass & Avogadro’s NumberA. Definitions:1. Molar Mass (Gram atomic weight): the mass in _________ that is numerically equalto the mass in _____.a. mass of 1 carbon atom _______________b. mass of 1 mole of C _______________2. Mole = ___________________________ of anything.3. Avogadro’s number = _________________4. Molar Mass = The mass in grams of ____________ (6.02 x 10 23 atoms ormolecules).5. What is the relationship between atomic mass units and grams?B. Calculations based on the mole.1. What is the mass of a Titanium atom in amu? Grams?2. What is the mass of one mole of Titanium?©RegliGower 53

3. How many (a) moles and (b) atoms of sulfur are there in 255 g of Sulfur?4. How many (a) moles and (b) grams of Uranium are there in a billion U atoms?5. How many silver atoms are there in a solid silver ball that is 2.00 cm in diameter?(D Ag = 10.5 g/cm 3 )6. How many Cu atoms are there in a solid copper penny that has a mass of 3.06 g?7. What is the mass of one atom of silicon in (a) amu (b) grams (c) pounds?8. What is the mass of 1500 sulfur atoms in grams?9. A student measures a solid aluminum block: l = 4.84 cm, w = 3.68 cm, h = 2.14 cm.If the specific gravity of aluminum is 2.70, how many aluminum atoms does it contain?©RegliGower 54

10. Mercury has a specific gravity of 14.7. How many moles of mercury are in a bottlethat contains 3.55 pints of mercury?III. Molecular Mass & Percent CompositionA. Chemical Formulas: Express the composition of molecules (and ionic compounds) interms of the ___________ for the elements they contain.B. Molecular Formula: Show the ________ and _________ numbers of ________. Theyare the _________ formulas of molecules!C. Diatomic Molecules: Memorize the 7 elements that occur as diatomic molecule…D. Molecules containing more than two atoms are called _____________________________.Examples:E. Different forms of the same element are called _________________Examples:F. Empirical (simplest) Formula: Show the ________ and ________ of atoms.MFWater Glucose Hydrogen peroxide BenzeneEFG. Ionic compounds only have ______________ formulas, they exist as 3D lattice______________ but not as individual ______________. When _______________ areremoved or added to a neutral atom (or group of atoms) a charged particle is formed called an_______. A positive ion is called a __________ while a negative ion is called a ________.Substances made of cations and anions are called ___________ _______________. Ionsthat contain only one atom are ______________ while those containing two or more are called__________________.H. Structural Formula: Show the kind, number, and arrangement of atoms.waterethane©RegliGower 55

IV. Molecular Mass (weight)A. The molecular mass is the ________ of the atomic masses in (amu) of all the _________in a molecule. The molar mass is the same mass but in _________.1. What is the molecular mass of benzene, C 6 H 6 ? What is the molar mass?B. Substances that do not form molecules such as ionic solids do not have a molecularweight; they have formula weights.1. What is the formula mass of aluminum sulfite? The molar mass?2. Calculate the molecular mass of TNT, C 7 H 5 N 3 O 6 .3. What is the formula mass (amu) of Cr(OH) 2 • 18 H 2 O?C. Percent Composition: The percent by ________ of each ___________ in acompound.1. What is the % composition of Ba(CN) 2 ?2. What is the % composition of ___________________, Fe 2 (SO 4 ) 3 ?3. What % hydrogen is water?4. % H 2 O in a Hydrate: Ionic substances are usually formed in water solutions. Thewater molecules get trapped in the crystal structure and form ____________.Examples: Calcium chloride dihydrate Plumbic nitrate heptahydrate©RegliGower 56

5. Ba(CN) 2 • 2H 2 O is what % water by mass?V. Formula Determination:A. Review Problems:1. What is the % composition of Ca(OH) 2 ? FW = 74.102. Ca(OH) 2 • 2 H 2 O is what % water by mass?B. Determination of Empirical Formula:1. Laboratory analysis shows that a compound is 5.80 % H and94.2 % S by weight (mass). What is its E.F.?Step 1: Calculate the moles of each element present. (Assume 100 % = 100 g)Step 2: Determine the ratio between moles (smallest whole numbers) by dividingby the smallest value in step 1.2. Remember that these values are measurements and need not be exact. Theyshould be close. What would be the formulas for elements “X” and “Y”.X 1.5 Y 1.0 X 1.00 Y 1.03 X 1.00 Y 1.33 X 1.00 Y 1.253. (a) Laboratory analysis shows that a compound has the following composition:Zinc 52.0 %, Carbon 9.60 %, & Oxygen 38.4 %. Determine its empirical formula.(b) What is the empirical formula of a compound that is 52.9 % Al and 47.1 % oxygen?©RegliGower 57

C. Determination of the molecular formula:1. You must be given the ____________________ weight.2. What is the relationship between the empirical formula weight and the molecularformula weight?3. Molecular formula = Empirical formula x __________________4. Analysis shows that a compound is 30.4 % Nitrogen and 69.6 % Oxygen. If it has amolecular weight of 92 g, what is its Molecular formula?Step 1: Calculate the Empirical formula.Step 2: Calculate the Empirical formula weight.Step 3: Divide the Empirical formula weight into molecular weight (must begiven).Step 4: Determine the Molecular formula.5. Analysis of a compound shows it to be 38.71 % Carbon, 9.68 % Hydrogen, and51.61 % Oxygen. If further data reveals that it has a molecular weight of about 93 g,determine the Empirical and Molecular formulas of the compound?D. Determining the Formula of a Hydrate:1. Hydrate: Is an __________ compound that has _______ molecular trapped in itscrystal _________ structure. The formula of the hydrate can be determined by heatingup a known mass of the hydrated salt in a crucible and then measuring the mass of_____________ salt formed.©RegliGower 58

2. A barium iodide hydrate is heated in a crucible. What is the formula of thehydrate? (BaI 2 • x H 2 O)H 2 O°°°°°°10.407 g BaI 2 • x H 2 O (Hydrated) + Flame → “Anhydrous Salt” 9.520 BaI 2 + H 2 OStep 1: Find the grams of water lost.Step 2: Find the moles of water and anhydrous salt.Step 3: Find the mole ratio.3. In lab a student heats 5.20 grams of hydrated zinc sulfite in a 20.40 g empty crucible.After heating, the anhydrous ZnSO 3 compound and crucible have a combined mass of23.00 g. What is the formula of the hydrate?VI. MOLES II: Moles & FormulaA. General Information:1. 1 mole = _________________________ number of anything.Avogadro’s number = ________________________________.2. 1 mole of atoms = ____________________________ atoms.1 mole of molecules = ____________________________ molecules.©RegliGower 59

3. The atomic weight of oxygen is _________________ amu.a. One oxygen atom has a mass of ________________ amu.b. One mole of oxygen atoms (6.02 x 10 23 atoms) has a mass of_____________. This is called the __________ atomic weight.4. The molecular weight (mass) expressed in ________ represents the mass of a___________ molecule, while the molecular weight expressed in ____________represents the mass of __________ of molecules; this is called the gram molecularweight.5. Substance the do not form molecules, such as ionic solids, do not have amolecular weight. They have a formula weight that when given in grams is called thegram formula weight and represents the mass of ___________ or the compound.6. Road Map:B. Sample Problems1. 0.500 moles of Fe 2 (CO 3 ) 3 (________________________) contains how many molesof a. Fe b. C c. O2. Calculate the mass of one water molecule in (a) amu (b) grams3. (a) How much does 4.50 moles of water weigh in grams? (b) How many molecules isthis?4. 50.0 g of methanol, CH 3 OH, are (a) how many moles? (b) contain how manymolecules?5. How much does a mole of _____________ (C 6 H 12 O 6 ) weigh in kilograms?©RegliGower 60

6. How many water molecules are there in a pint of water?7. 50.0 g of ___________________ (Ca(OH) 2 ) contain how many (a) moles of oxygen?(b) oxygen atoms?8. What is the volume in mL of 5.55 x 10 24 water molecules?9. A spherical metal tank contains carbon monoxide gas under pressure. If the tankhas a diameter of 4.00 ft. and the density of the gas under pressure is 7.50 g/L, howmany molecules are there in the tank?VII. Balancing Chemical EquationsA. Chemical equations must be balanced due to the Law of Conservation of ________.(_______________). Reactants are written on the left of the arrow, ________________ arewritten to the right of the arrow.______________________→______________________B. Equations are balanced by changing ____________________, not by changing________________. For now all equations will be balanced using ____________ numbers.C. Certain elements will are found as __________________ molecules in nature. In achemical equation, they will be written in this form. The elements are easily remembered bylearning the name of the German guy:Mr. __________________________. The seven diatomic molecules are:_____, _____,_____, _____, _____, _____, and _____.©RegliGower 61

Example: Lighting a match:1. Observations:2. Word Equation:phosphorus + potassium chlorate → potassium chloride + diphosphorus pentoxide3. Skeleton Equation:P 4 + _________________→__________________+__________________4. Balanced Equation:D. Coefficients: Small, ________________ numbers placed before the chemical formula.They are multiplied by each atom in the compound.E. Law of Conservation of Mass: Matter cannot be _______________ nor ______________,so reactions must be ________________.F. Tips for balancing equations:1. Number of atoms of _______________ must equal number of atoms of __________.2. Coefficients are whole numbers written at the __________ of the substances.3. All atoms are ___________________ by the coefficients.4. Subscripts are ________________ changed.5. Keep polyatomic ions together as a ___________________ if unchanged fromreactants to products.6. Balance single elements __________.7. Use the even/odd rule.8. If an element is in ________________ compounds, balance that element last.G. Example:1. _____NaClO 3 → _____NaCl + _____O 22. _____Fe 3 O 4 + _____H 2 → _____Fe + _____H 2 OH. Sample Problems:1. Hydrogen + oxygen → water2. Zinc + hydrochloric acid → zinc chloride + hydrogen©RegliGower 62

3. Copper + silver nitrate → cupric nitrate + silver4. Ferric hydroxide → iron (III) oxide + water5. Ethane (C 2 H 6 ) + oxygen → carbon dioxide + water6. Calcium + water → calcium hydroxide + hydrogen7. Potassium + sulfuric acid → potassium sulfate + hydrogen8. Calcium nitrate + aluminum sulfite → calcium sulfite + aluminum nitrate9. Phosphoric acid is formed when crystalline diphosphorus pentoxide is dissolved inwater.10. When rust (ferric oxide) is dissolved in hydrochloric acid, it dissolves formingaqueous ferric chloride and water.11. Benzene, C 6 H 6 , is an organic solvent used to dissolve many organic compounds.Write an equation for the combustion of liquid benzene (carbon dioxide and water areformed).12. When a solution of aqueous plumbic nitrate and aqueous barium hydroxide aremixed, solid plumbic hydroxide and aqueous barium nitrate are formed.©RegliGower 63

VIII. Balancing Chemical Equations: Algebraic TechniqueA. Balance the following equation:Ca 3 (PO 4 ) 2 + H 2 SO 4 → CaSO 4 + H 3 PO 4STEP 1: Assign letter to unknown coefficients:a Ca 3 (PO 4 ) 2 + b H 2 SO 4 → c CaSO 4 + d H 3 PO 4STEP 2: Make a grid indicating the appearance of element or ion in each species of theequation. Use whole number and the coefficient to indicate the appearance.Ca 3a + 0 = c + 0PO 4 2a + 0 = 0 + dH 0 + 2b = 0 + 3dSO 4 0 + b = c + 0STEP 3: Reduce the equations:3a = c2a = d2b = 3db = cSTEP 4: Assume a = 1; solve for coefficients (**any variable can assumed to be 1)a = 1 b = 3 c = 3 d = 2STEP 5: Write equation with coefficients:B. Examples:Ca 3 (PO 4 ) 2 + 3 H 2 SO 4 → 3 CaSO 4 + 2 H 3 PO 41. ______Na 2 CO 3 + ______C + ______Sb 2 S 3 → ______Sb + ______Na 2 S + ______CO 22. ______C 6 H 6 + ______O 2 → ______CO 2 + ______H 2 O©RegliGower 64

3. ______HClO 4 + ______P 4 O 10 → ______H 3 PO 4 + ______Cl 2 O 74. ____MnO 4 + ____CaC 2 O 4 + ____H 2 SO 4 → ____CaSO 4 + ____Mn + ____CO 2 + ____H 2 O5. _____FeS 2 + ______O 2 → ______Fe 2 O 3 + ______SO 26. ______C 7 H 6 O 2 + ______O 2 → ______CO 2 + ______H 2 O©RegliGower 65

Unit 5: StoichiometryI. Stoichiometry = The ____________ relationships among reactants and products inchemical reactions.A. A chemical equation is a statement of __________________ fact. On the leftside of the reaction are the __________________, and on the right side the________________ of the reaction. Since no atoms are created nor destroyed in a__________________chemical reaction, the equation must be ________________.This means that the combined weight of the reactants is exactly ___________ tothe combined weight of the products. In terms of chemical laws, this is theLaw of _______________________ of matter (_____________________).B. Stoichiometric “road map” (Use the balanced chemical equation)Examples:(1) Calcium reacts with oxygen to form calcium oxide. If you are given 80.0 g of calcium, (a) how many grams of calciumoxide can be made? (b) How many moles of oxygen gas are required?©RegliGower 66

(2) Nitrogen and oxygen react to form dinitrogen pentoxide. If 112 g of nitrogen react with unlimited oxygen, (a) howmany moles of oxygen react and (b) how many grams of dinitrogen pentoxide are formed?(3) Aqueous sodium chloride and aqueous silver nitrate react forming solid silver chloride (a precipitate) and aqueoussodium nitrate. If one solution contains 100.0 g of silver nitrate and it reacts with unlimited sodium chloride, how manygrams of silver chloride can be made?(4) How many (a) grams and (b) molecules of carbon dioxide gas are produced when 100.0 g of methane gas (CH 4 ) areburned. Methane reacts with oxygen to form carbon dioxide and water. (c) If air is 22.0% oxygen by mass, what mass ofair is needed to supply enough oxygen?(5) Nitrogen reacts with hydrogen gas forming ammonia (nitrogen trihydride). How many (a) grams of hydrogen and (b)moles of nitrogen are required to form 1.00 pound of ammonia?(6) How many oxygen molecules are produced if 1.00 gallon of water is electrolyzed into hydrogen and oxygen gas?©RegliGower 67

Review Question:Nitrogen and hydrogen gas combine to form hydrazine (dinitrogen tetrahydride) which is used for rocket fuel. How manygrams of nitrogen gas are needed to form 155 g of hydrazine?II. Limiting ReagentA. Stoichiometric amounts: The proportions indicated in the ________________ rxn.B. Most reactions do not have stoichiometric amounts. Generally, one reactant will be__________________ before the other. The reactant that is depleted first is known asthe _________________________. The reactant that is left at the end of the reaction iscalled the __________________________.C. Analogy: How to make a cheese sandwich.2 slices of bread + 1 slice of cheese → 1 cheese sandwichIf you have 8 slices of bread and 6 slices of cheese, how many sandwiches canyou make? _________ (theoretical yield)What is the limiting reagent? _____________What is the excess reagent? _____________How much of the excess reagent is left at the end of the rxn? ___________D. Theoretical yield: The amount of product that forms if all of the __________________reagent has reacted. (This number is CALCULATED!)E. Actual yield: The amount of product that is actually made (Experimental).F. Percent yield: The comparison of the actual yield to the theoretical yield.G. Limiting reagent problems(1) How much hydrogen peroxide can be formed from 10.0 g of hydrogen and 125 g of oxygen? What is the limitingreagent?©RegliGower 68

(2) Determine the theoretical yield of water that can be formed from 10.0 g of hydrogen and 100.0 g of oxygen. What isthe limiting reagent?(3) How many grams of aluminum oxide can be formed from 100.0 g of aluminum and 100.0 g of oxygen gas?(4) A student does a lab in which she makes 2.17 g of plumbous nitrate. From her limiting reagent she calculated thatshe should have made 2.55 g. Determine: (a) theoretical yield, (b) actual yield, (c) % yield (d) absolute and relative error.(5) A sample of 100.0 g of hydrazine (N 2 H 4 ) burns in 280.0 g of oxygen gas and 97.2 g of water are formed. Calculatethe % yield of water. (Nitrogen dioxide is also produced.)©RegliGower 69

(6) In lab a student mixes a solution that contains 1.70 g silver nitrate and a solution that contains 0.832 g calciumchloride. Aqueous calcium nitrate and solid silver chloride are produced. After washing, filtering, and drying theprecipitate, he finds that he has obtained 1.25 g of silver chloride. What is the % yield?III. Reaction TypesSymbols: Solid (___) or (___); Liquid (____); Gas (____); Aqueous (dissolved in water) (____)A. Synthesis (__________________________)General form:1. element + element → compound (Use ions to form compound)a. sodium + chlorine →b. calcium + oxygen →c. lithium + sulfur →d. aluminum + oxygen →2. metal oxide + water → metal hydroxidea. Na 2 O(s) + H 2 O(l) →b. CaO(s) + H 2 O(l) →c. Al 2 O 3 (s) + H 2 O(l) →©RegliGower 70

B. DecompositionGeneral form:1. Compound → ________________________+_________________________a. FeCl 3 (s) →b. CuBr(s) →c. MgO(s) →2. Metal hydroxide → ___________________ + ___________________a. Be(OH) 2 (s) →b. Mn(OH) 2 (s) →c. CuOH(s) →3. Metal chlorate → ___________________ + ___________________a. KClO 3 (s) →b. Zn(ClO 3 ) 2 (s) →C. Double Displacement (___________________)General form:**Use solubility table for determining Ionic and Net Ionic equations.**Solubility Rules: (MEMORIZE!!)1. ALL alkali metal compounds are soluble.2. ALL ammonium compounds are soluble.3. ALL compounds containing nitrate are soluble.4. Most hydroxide compounds are INSOLUBLE. Alkali metal hydroxide are exceptions.5. Most chlorides, bromides, & iodides are soluble. The exceptions are:Ag + , Hg 2 2+ , Pb 2+©RegliGower 71

1. HCl(aq) + NaOH(aq) →Ionic:Net Ionic:________________________________________________________________________________2. AgNO 3 (aq) + BaCl 2 (aq) →Ionic:Net Ionic:________________________________________________________________________________3. FeCl 3 (aq) + KOH(aq) →Ionic:Net Ionic:________________________________________________________________________________4. Pb(ClO 3 ) 2 (aq) + NH 4 Cl(aq) →Ionic:Net Ionic:________________________________________________________________________________5. Ca(C 2 H 3 O 2 ) 2 (aq) + (NH 4 ) 2 CO 3 (aq) →Ionic:Net Ionic:________________________________________________________________________________©RegliGower 72

D. Single Displacement (____________________)(Fig. 4.15) Activity Series of MetalsLiKBaCaNa React with cold water to produce H 2MgAlZnCrFeCd React with steam to produce H 2CoNiSnPb React with acids to produce H 2HCuAgHgPtAuDo NOT react with water or acids1. Cation displacement - ****MUST use the Activity Series of Metals****General form:(Single element must be ____________ reactive than the element it replaces).a. Sn(s) + NaNO 3 (aq) →b. Zn(s) + H 2 SO 4 (aq) →c. K(s) + CaCl 2 (aq) →d. Na(s) + H 2 O(l) →e. Cu(s) + HCl(aq) →©RegliGower 73

2. Anion displacement - ****MUST use the periodic table****(Single element must be __________ reactive than the element it replaces.General form:Halogens (Column _______)Reactivity: _______>_______>_______>_______a. Cl 2 (g) + NaBr(aq) →b. I 2 (s) + KBr(aq) →c. F 2 (g) + CaCl 2 (aq) →E. Combustion (_________________) General form:1. C 2 H 6 (g) + O 2 (g) →2. C 3 H 8 O(l) + O 2 (g) →IV. Concentration of a solutionA. Concentration units1. Percent by mass = _____________________ x 1002. Percent by volume = _____________________ x 1003. Molarity = ____________________B. Example calculations(1) If 3.50 g of sugar is dissolved per 85.5 g of solution, what is the % sugar by mass?©RegliGower 74

(2) If there is 12.5 mL of alcohol dissolved in 128.4 mL of solution, what is the % alcohol by volume?(3) If 4.44 g of sodium chloride is dissolved in 122 mL of solution, what is the molarity?(4) Compare the ion concentrations of 1.20 M glucose, 1.20 M KCl, and 1.20 M Ba(NO 3 ) 2 .V. Preparation of a solution of known molarityA. Steps:1. Calculate the mass required.2. Weight this amount and put it into a __________________ flask.3. Add _______________ water and swirl to dissolve.4. After all the ________________ has dissolved, add water to the line.B. Example: Prepare 500 mL of 0.200 M potassium dichromateVII. Molarity calculationsUse dimensional analysis; start with the number other than the molarity (mass or vol.)(1) What is the molarity of a solution that contains 3.45 g of sodium hydroxide dissolved in 455 mL ofsolution?©RegliGower 75

(2) How many grams of ferric nitrate are there in 125 mL of 0.150 M solution?(3) What volume of 0.200 M silver nitrate solution is required to obtain 1.45 g of silver nitrate?(4) Explain how you would make 750 mL of 0.125 M barium chloride solution.(5) What is the concentration of each ION in 0.550 M solutions of:(a) potassium dichromate(b) ammonium phosphate(c) hydrochloric acid©RegliGower 76

VIII. Dilution of solutions:Concentrated solutions are often stored in chemical stockrooms and are called ________________solutions. Preparation of less concentrated solutions from a concentrated stock solution is done bythe process of __________________. This is accomplished by increasing the amount of______________(________________).Equation:Examples:(1) Explain how you would prepare 1.50 L of 0.0500 M potassium permanganate from a stock 1.50 Mpotassium permanganate solution.(2) Explain how you would prepare 250 mL of 0.250 M sulfuric acid solution from a concentrated 10.5M solution.(3) If 125 mL of distilled water is added to 115 mL of 0.100 M hydrochloric acid, what will be the finalconcentration of the solution?(4) What is the [Cl - ] if 125 mL of 0.100 M sodium chloride is mixed with 225 mL of calcium chloride.©RegliGower 77

IX. Quantitative Aspects of Reactions in Aqueous SolutionsA. Quantitative Analysis: the determination of the _________ or ___________________of a substance in a sample.B. Qualitative Analysis: The determination of the types of ___________ present in asolution.C. Gravimetric Analysis: An experimental procedure based on the measurement of___________.D. Precipitation Analysis: Finding the composition of a component of a compound bydepositing it from a solution as an _________________ compound called a __________.E. Quantitative Analysis Example: Determination of % by mass of Cl - in a sample.1. Weigh the sample; dissolve it in _____________ water. (Why distilled water?)2. Add enough silver nitrate of known concentration to ___________________ allof the Cl - as _____________.3. The silver chloride precipitate is separated from the mixture by _____________and then it is ______________ and _______________.4. From the amount of AgCl, determine the amount of _______ present.5. The % by mass of Cl - in the original sample is calculated by:F. How do you know which solution to use? Solubility rules!!!Solubility Rules:1. ALL alkali metal compounds are soluble.2. ALL NH 4 + compounds are soluble.3. ALL compounds containing NO 3 - , ClO 3 - , and ClO 4 - are souble.4. Most OH - compounds are INSOLUBLE. Alkali metal hydroxides and Ba(OH) 2 areexceptions.5. Most Cl - , Br - , and I - compounds are soluble. The exceptions are: Ag + , Hg 2 2+ , Pb 2+ .6. All CO 3 2- , PO 4 3- , and S 2- compounds are INSOLUBLE. Exceptions are alkali metaland ammonium compounds.7. Most SO 4 2- compounds are soluble. Exceptions are Ba 2+ , Hg 2+ , and Pb 2+ .1. Example: How could you precipitate PbI 2 ?©RegliGower 78

2. Example: How could you precipitate Mn(OH) 2 ?3. Example: How could you separate a mixture containing Cu 2+ , Ba 2+ , and Ag + ?Quant example problems:(1) In lab a student forms 5.34 g of silver chloride precipitate by reacting 1.775 g of an unknown chloride sample with anexcess of silver nitrate solution.(a) Write the net ionic equation.(b) Calculate the % by mass of Cl - .(c) If the formula is XCl 2 , what is element X’s molar mass and identity?(2) A student mixes 125 mL of 0.124 M aluminum nitrate and 175 mL of 0.104 M barium hydroxide.(a) Write the molecular, ionic, and net ionic equations. Identify the precipitate.(b) Determine the mass of precipitate formed.©RegliGower 79

(c) Calculate the concentration of excess cation (M).(3) A student wants to determine the molarity and % by mass of an NaCl solution. He slowly adds 0.100 M AgNO 3 untilhe observes that no more precipitate forms. He has added 12.45 mL of the silver nitrate to 100.0 mL of the NaCl solution.(sp. gr. = 1.00)(a) What is the molarity of the NaCl solution?(b) What is the % by mass of the NaCl solution(4) A 0.08750 g sample of an unknown compound containing phosphate ions is dissolved in water and treated withexcess calcium nitrate and 0.6764 g of precipitate forms.(a) What is the precipitate? Write the net ionic equation.(b) Calculate the % by mass of phosphate in the original sample.©RegliGower 80

(5) A student wishes to determine the amount of NaCl and KCl in a 1.557 g sample that is a mixture of both. Shedissolves the sample in distilled water and then reacts it with excess silver nitrate. He obtains 3.408 g of dry precipitate.Picture:(a) What is the precipitate? Write the net ionic equation.(b) What is the mass of chloride in the mixture?(c) What is the % of KCl and NaCl originally present in the sample?©RegliGower 81

Unit 6: Atomic Theory II /Periodic TrendsChapter 7: Atomic Theory III. History of the AtomA. ________________: Greek philosopher from 400 BC, “atomos”B. ________________: Atomic Theory, based on chemical behavior, around 1800C. ________________: Plum pudding model of the atom (First atomic model),after he discovered the _______ with the _______________________.D. ________________: Planetary model of the atom, after he discovered the ______with the _________________________ and ______ particles.Problems: The planetary model does not explain…1. why electrons do not ________ energy and _______ into the nucleus.2. the atomic line __________.3. ____________ trends of the elements.E. ________________: The Bohr Model of the atomF. ______________________________: Our current model of the atom.II. Radiant EnergyA. What is light? (1600s)1. Particle Theory: ___________________2. Wave Theory: _____________________B. What is light? (today)1. ________________________ dualityIII. Electromagnetic Radiation (Light behaves as a wave)A. Electromagnetic Theory1. James Clerk ________________ (~1870) proposed that accelerating_____________ produce a changing ____________ field, which produces achanging ________________ field, creating electromagnetic radiation(____________).2. Maxwell’s proposal was experimentally verified by _____________________using a ______________________.Diagram:3. Hertz created the first ________________ transmitter and receiver.©RegliGower 82

B. Electromagnetic Spectrum: The range of ________________ or ________________over which electromagnetic waves occur.BlueRedC. Characteristics of Electromagnetic Waves1. Speed: All EM waves travel at the speed of ___________ (in a vacuum)c = ____________________ m/s or ___________________ mi/sExample: How many kilometers is the sun from earth if it takes 8.1 minfor light to reach the earth from the sun?2. Frequency: The number of wave _____________ per second. The symbolscommonly used are ______ or ______. The units are _______ = _______ = ______.3. Wavelength: The length of one complete wave cycle, often measured fromcrest to crest, or trough to trough. The symbol used is ______. It is generallymeasured in __________.Diagram:4. The Wave Equation: Velocity = frequency x wavelength©RegliGower 83

D. Sample Calculations:1. If a red light has a wavelength of 725 nm, what is its frequency?2. What is the wavelength (in meters and Angstroms) of an FM radio station thathas a frequency of 101.5 MHz?IV. Quantum Theory (Light behaves as a particle)A. In 1900, ___________________ proposed that energy comes in ___________amounts called _______________.1. A quantum is a bundle, packet, or _________________ of energy.2. A quantum of light energy is called a ______________.B. The energy of a photon (“particle” of light):E =h =f =C. Graph of energy versus frequency:ED. Another unit of energy is the __________________(_____). The conversion factor is:fE. Using the wave equation (____________), the energy of a photon can be calculated interms of wavelength:©RegliGower 84

F. Electromagnetic Spectrum_________λ_________f_________E_________λ_________f_________E________________________________________________________________________________G. Sample Problems:1. Calculate the energy, speed, and frequency of monochromatic blue light thathas a wavelength of 490 nm.2. If a photon has an energy of 7.55 x 10 -19 J, what is its wavelength in nm?H. ____________________ further showed that light behaved as a ______________ withhis theory of the _______________________. He suggested that a beam of light is reallya stream of ___________________. He concluded that a photon of light contained aquantum of energy, E = hf. This work won him the Nobel Prize in Physics in 1905.V. Bohr Model of the Hydrogen AtomA. The ____________________ model (by __________________) did not explain: (1)the ________________ collapse problem, (2) periodic __________________, and (3)atomic _______ spectra.B. Physicist ______________________ applied the newly developed quantum idea (lighttravels as a packet of energy that we call a ______________) to the hydrogen atom.1. The electrons are found in _______________ levels outside thenucleus called _______________ or orbits.2. Electrons could only be found in these energy levels, and they could not loseenergy and collapse into the nucleus. Why? _______________________________.3. Energy levels were designated by the n, which referred to the______________quantum number.©RegliGower 85

C. Diagram:D. Atomic Line Spectrum: The Bohr Model was able to explain the _________________line spectrum.1. When ______________ is added to an electron, it will jump to _______________energy levels, known as the ____________________ energy state.2. When the excited electron falls down to a ____________ energy level,it releases ______________ in the form of a ______________ of light. Thefrequency or wavelength of that photon could be determined by the equations:3. The larger the gap in energy levels, the larger the _____________ of thephoton, therefore the _______________ the wavelength, and the _______________the frequency.4. The spectral lines are due to light emitted by the falling electron. Since thespacing between energy levels is constant, the light emitted for a particularelectron jump will always be the same _________________ and _______________.©RegliGower 86

Spectral lines: Lyman Balmer Paschen BrackettSeries Series Series SeriesLight: ______ ______ ______ ______Jumps from excitedstates to: ______ ______ ______ ______**Note: n = 1 is known as the ___________________ state. All others are __________________states.Line spectrum (_______________________)E. Bohr Calculations1. Each spectral line corresponds to difference in ________________(_____)between two permissible ___________________.a. Calculation of energy, f, wavelength of photon emitted (or absorbed):©RegliGower 87

. Calculation of the radius of orbits (R n ):c. Calculation of the energy of each energy level (shell), E n :F. Sample Problems:(1) Calculate the radius and energy of the 3 rd shell of the hydrogen atom.(2) What are the frequency and wavelength of the photon emitted when an excitedelectron jumps from the 3 rd shell to the ground state in a hydrogen atom?(3) (a) What is the frequency and wavelength of the photon absorbed when anelectron jumps from the 2 nd to the 6 th shell? Given: E 6 = -0.378 eV, E 2 = -3.40 eV.©RegliGower 88

(b) How would the frequency and wavelength of the emitted photon compare tothe absorbed photon if the electron jumped back from the 6 th to the 2 nd ?(4) What frequency photon is needed to cause an electron in the hydrogen atom tojump from the ground state to the 6 th shell?(5) An excited hydrogen atom emits a photon of 95.0 nm when an electron jumps from ahigher energy level to the ground state. From what shell did the excited electron jump?VI. Quantum Mechanical Model of the AtomA. Problems with the Bohr Model1. It could only explain the atomic line spectrum for _______________________elements:2. There was no proof of _________________ orbits.3. It did not explain the electron _______________ problem of circular orbits.4. Could not explain chemical __________________.B. Changes from the Bohr Model1. Treats the electron as a ______________ and as a ________________.2. Develops a model for the _____________________ of finding an electron ina space outside the nucleus, called an ___________________ or_________________________.3. We now think of an electron as: ____________, _____________, ____________©RegliGower 89

C. Scientists who contributed to the Quantum Mechanical Model1. Louis __________________(1929)a. Wave-particle duality of _________________.b. If a wave can behave like a particle, then a ________________ canbehave like a wave.c. Electrons show diffraction patterns very similar to _____________,which verifies the wave properties of matter. See Figure 7.14 (p. 277).d. Derivation of de Broglie’s wave equation:e. The wave equation could be applied to all systems, but is only_____________________ for very small objects. For large objects,like a baseball, the wavelength would be so small, that it would beundetectable.f. Example: Calculate the wavelength of a 2.00 kg ball moving 5.00 m/s.g. Example: Calculate the de Broglie wavelength of an electron moving10.0% the speed of light. (Write the answer in meters and Angstroms).h. de Broglie’s idea explains why electrons can only be present in specific____________________ (e.g. electrons can only be on the rungs of theladder, but never between), and why electrons in an energy level do notlose ________________.©RegliGower 90

i. de Broglie thought that electrons behaved like _________________waves, and that the wave must fit the _____________________ of theorbit exactly. Otherwise, the wave would _______________ itself out.Standing waves(a) circumference = integral # of λ(b) circumference ≠ integral # of λ, so thewave would cancel itself out.j. Each orbit corresponds to a _______________ wavelength, thereforeelectrons can only have particular amounts of energy. Since a waveis continuous, electrons do not ____________ energy.2. Werner Heisenberg: The Heisenberg __________________ _________________a. It is not possible to know the exact ___________________ and___________________ of an electron at the same time.b. As the uncertainty of the position decreases, the uncertaintyof the momentum (mass x velocity) __________________.c. This mathematical equation shows that an electron cannotmove in a well-defined ___________ as Bohr thought. If an electronmoved in a circular orbit, we could know the position and momentumat the same time.d. We do not know the exact path of an electron, but we can call theregion of space where the electron has a high _________________of being found the ____________________ or electron cloud.Diagram:©RegliGower 91

3. Erwin Schrödinger: The Schrödinger ______________ Equationa. The equation (which involves advanced calculus) describes thebehavior and _______________ of electrons.b. It incorporates both the __________________ and __________________nature of matter.c. The wave equation can only be solved for _______ electron systems. Allothers are only approximated.d. The wave function = _______ (psi).e. ψ 2 = _________________ of finding an electron in a particle region ofspace.f. The solutions to the wave equation yields the ______________numbers.D. Quantum Numbers1. The ________ quantum numbers are used to describe the most _____________location of the electron in the atom.2. Analogy: Where is the most probable place to find you at 3 a.m.?______________ _______ ___________ ________Principal Energy Level Sublevel Orbital Spin(n) → (l) → (m) → (s)3. The four quantum numbers:a. Principal Quantum Number (n)i. Gives the principal ________________ of the electronii. Possible values: n = 1, 2, 3, …_____iii. Represents the average ______________ from the nucleus. Thelarger the number the _______________ the electron is from thenucleus.b. Angular Momentum Quantum Number (l)i. Determines the ______________ of the electron.ii. Gives the “shape” of the ___________iii. The value of l depends on the value of n: l = 0, 1, 2, …______©RegliGower 92

iv. Examples of shapes:l = 0 (s orbital)l = 1 (p orbital)l = 2 (d orbital)l = 3 (f orbital)8 – lobed orbital of complex shapes (not shown in our textbook)l = 4 (g orbital), l = 5 (h orbital), l = 6 (i orbital)**** There are currently no atoms large enough for electrons to be presentin any of these sublevels.v. Chart of sublevel types for each energy level©RegliGower 93

Value of n Possible Values of l Types of Orbitals12345c. Magnetic Quantum number (m l )i. Gives the _____________________ of the angular momentum.ii. It describes the magnetic ____________ generated by anelectron as it moves around the nucleus.iii. The values of m depend on l: __________ to __________.iv. It gives the ___________________ of the orbitals in space.v. Chart of Orbital Orientations for Each SublevelValue of l Type of Sublevel Possible Values of m l # Orbitals0123d. Spin Quantum Number (s)i. Every electron behaves in certain respects as though it is aspinning charged sphere. There are _______ possible directions ofspin.ii. Possible values of s: ________ and ________.iii. Electrons within an orbital must have _________________ spins.iv. Diagram showing electron spin©RegliGower 94

****Electrons are represented with arrows to indicate spin direction.v. Stern and Gerlach experimentally verified spin with the followingexperiment:vi. A beam of hydrogen atoms is passed through a ______________field. Half of the atoms are deflected ____________, the other halfare deflected downward. This means that on average, half of theatoms have electrons spinning in one direction, the other half arespinning in the opposite direction.4. Pauli Exclusion Principle: No two electrons can exist in exactly the___________ quantum state. Every electron in atom will have a _______________set of quantum numbers.E. Sample Questions:1. What are the four subshells? Which quantum number represents subshell?2. How many (a) orbitals, and (b) electrons maximum are there in each subshell?(i) s (ii) p (iii) d (iv) f3. Complete the table:©RegliGower 95

Value of n Values of l Subshells # Orbitals Max. # electrons1 0 1s 1 22344. What is the maximum number of electrons that could occupy:(a) a 1s orbital?(b) a 2p subshell?(c) the 3 rd energy level?(d) a 2p orbital?5. How many orbitals are there in:(a) a 2p subshell(b) a 3f subshell?(c) the 4 th energy level?(d) the 1 st energy level?6. Give the four quantum numbers for all of the electrons in a full:(a) 6s orbital(b) 3p subshell(c) 4f orbital7. In which subshell is each of the following electrons found?(a) n=5, l=1, m=0, s=+ ½ (b) n=4, l=3, m=2, s=- ½ (c) n=3, l=4, m=5VII. Electron Configuration: Shows how the electrons are arranged in an atom.©RegliGower 96

A. The electron configuration is written in the following format:C (______electrons): 1s 2 2s 2 2p 2number of electrons in that sublevelsublevel (l)Energy level (n)(The sum of all of the superscripts should equal the __________ # of electrons.)B. The orbital diagram uses arrows to designate the electrons:CC. How to write an electron configuration:1. Start filling electrons in the subshells of _______________ energy first, thenbuild up (aufbau principle).2. Only ________ electrons per orbital.3. Subshell # orbitalsspdf4. Electrons within an orbital much have ______________ spins to followthe ___________ Exclusion Principle.5. Hund’s Rule: If there are multiple orbitals of the same energy, each orbitalis first filled with only ______ electron, before they begin to _________ up.D. Auf Bau Diagramexample: (4 electrons) ___ ___ ___ ___ ___ ___E. Examples:©RegliGower 97

1. Boron2. Phosphorus3. Potassium4. IronF. Using the Periodic Table to determine electron configurations:***Label your periodic table with the energy subshells***G. Kernal Configurations©RegliGower 98

1. Use the preceding _______________ as the kernel and place it in brackets.2. Examples:NickelSiliconTungstenH. Irregularities with the Transition Metals: Half-filled and filled ______ subshells aremore stable than other configurations, so some atoms have irregular electronstructures.MEMORIZE THESE EXCEPTIONS!!1. Examples:ChromiumCopperSilverGoldI. Valence Electrons: The electrons in the outermost _____ and _____ subshells.1. The value of _____ must be the same for the s and p.2. The number is never greater than _______ (___________)3. Elements can be represented with the dot formula:Li Be B C N O F Ne©RegliGower 99

VIII. Review of Quantum Mechanical ModelA. Example: NeB. n 2 = # __________ in an energy level.2n 2 = # __________ in an energy level.C. Electron configurations of transition metal cations (lose e- from VALENCE first!)Fe Fe 2+ Fe 3+Sn Sn 2+ Sn 4+I. History of the Periodic TableChapter 8: Periodic TrendsA. In 1869, ___________________________ organized the elements by increasing______________________. The organization showed similar ________________ forelements in the columns. He left gaps where some elements had not yet beendiscovered and predicted their __________________ based on the properties ofother elements in the column. For example, there was a gap below Aluminum. Hecalled the undiscovered element “___________________” (first element belowaluminum). He predicted its properties. When the element was discovered (andcalled __________________), its true properties were compared to the predictedproperties of eka-aluminum:Eka-AluminumGalliumAtomic mass 68 amu _______Melting Point Low _______Density 5.9 g/cm 3 _______Formula of oxide Ea 2 O 3 _______©RegliGower 100

II. Periodic Classification of the ElementsIII. Electron Configuration of Cations and AnionsA. Examples: Na Na +Al Al 3+F F -O O 2-B. Isoelectronic: Atoms and ions that have the _______ number of electrons and the________ ground-state electron configuration.IV. Atomic RadiusA. Atoms do not have __________________ boundaries, so it is difficult to determineatomic size.B. Atomic radius is defined as _______________ the distance between two nucleiin adjacent atoms.©RegliGower 101

C. Periodic trend in atomic radius:1. Within a group: Atoms increase in size going ___________ the columnbecause elements in each successive period have an extra energy level.Examples: NaK2. Within a period: Atoms increase in size going from _________ to _________in a row!!a. WHY? As the number of ____________ in the nucleus increaseswithin a period, the electrons are pulled inward with a greater __________.b. As Z eff _______, atomic radius _______.c. Z eff = Effective nuclear charge = Z – inner electronsd. For neutral atoms: Z eff = # of ____________________.e. Example: Calculate the Z eff of Li and O. Which atom would be larger?3. General trend of atomic radius:_______________V. Ionic RadiusA. Cations (___________): Are _______________ than the corresponding ____________atom.1. Example: Calculate Z eff for Na and Na+2. As Z eff ______, ionic radius _______.B. Anions (___________): Are _______________ than the corresponding neutral atom.©RegliGower 102

1. Example: Calculate Z eff for F and F -2. Z eff remains _____________________ as electrons are added to an atom.3. Why do atoms get larger when they gain electrons? More e- ______________.C. Example: Circle the larger atom.VI. Chemical Reactivity(1) Br, I (2) O, O 2- (3) Ca 2+ , Ca (4) V, Fe (5) C, SiA. Nonmetal reactivity: _______________ Most reactive nonmetal:_______________Same trend for Nonmetallic character.B. Metal reactivity: _______________ Most reactive metal:________________Same trend for Metallic character.WHY is cesium more reactive than lithium?©RegliGower 103

VII. Ionization Energy (IE)A. The minimum energy needed to remove an _____________ from a gaseous stateatom, generally measured in kJ/mol.X (g) + ionization energy _______ (g) + _______ (g)B. Gaseous atoms are used because they are uninfluenced by their neighbors.C. In a multi-electron atom, more than one electron can be removed from the atom.When the second electron is removed, we call it the ________________ IE, etc.X + (g) + IE 2 _______ (g) + _______ (g)Second IEX 2+ (g) + IE 3 _______ (g) + _______ (g) Third IED. It takes ___________ energy to remove the second electron than the first, becausethe positively-charged ion has a larger attractive force towards the remaining electrons.E. Ionization Energy DataF. Ionization Energy Periodic Trend1. Within a group: IE _____________ going down a column. Why? Asan atom gets larger, the electron is ______________ to remove because thereis less pull to the nucleus (more e- shielding and larger distance from nucleus).2. Within a period: IE generally increases going from __________ to ________on the periodic table.Which group has the lowest IE?_____________Why?Which group has the highest IE?_____________Why?©RegliGower 104

3. Within a period there are _____ exceptions to the general trend.a. Group IIIA has a LOWER ionization energy than Group ______.b. Group VIA has a LOWER ionization energy than Group ______.Examples:Be vs. BN vs. O*** Half-filled and Full subshells are relatively stable ***G. By studying the IE of an element, you can determine the number of ___________ e-of an atom. Outer (valence) electrons take significantly ________ energy to removethan an inner electron.Example: BeVIII. Electron Affinity (EA)A. The negative of the energy change that occurs when an electron is ___________ byan atom in the gaseous state to form an ____________.X (g) + e - X - (g)F (g) + e - F - (g) Energy = -328 kJ/mol; EA = +328 kJ/molB. The more POSITIVE the EA, the more easily an atom accepts an electron. Negativeelectron affinities are for atoms that do not want to accept an electron.C. Which group has the highest EA? ______________ Why?D. Why do the Alkali Metals have a relatively high EA?©RegliGower 105

Unit 7: Chemical BondingChapter 9I. IntroductionA. What is a chemical bond? A _________ that holds atoms together.B. Why do atoms bond? To become more ___________. Maximum ____________occurs when electrons are the same as a _____________.C. Forces of _________________> Forces of __________________D. Types of bonds:1. Ionic: _____________ of electrons.2. Covalent: ________________ of electrons.3. Metallic: _______________ electrons.II. Valence electronsA. The ________________ electrons (usually the ____ and ____ of the highest shell).B. These are the electrons involved in _____________.C. Transition elements sometimes use ____ subshell electrons for bonding.D. Octet rule: Atoms tend to form bonds until the yare surrounded by _____ valence e-.Exceptions: He Li + Be 2+III. Ionic BondingA. The _____________________ of electrons from a ____________ to a ____________.Cations and __________ are formed which _______________ one another with an________________ force.B. Lewis dot formula for ions:SodiumFluorineMagnesiumOxygenSodium ionFluorideMagnesium ionOxideB. Examples:Sodium chlorideMagnesium chloride©RegliGower 106

C. Ionic Structure1. Crystal ___________: 3-D arrangement of ____ and ____ ions in an ionic solid.2. Lattice energy: The energy _____________ when 1 mol of an ionic solid is________________ into its ions.a. Always ____________________ (energy absorbed).b. Lattice energies are generally large values, which shows that:i. ionic bonds are _____________.ii. ionic solids have relatively _______ melting points.iii. ionic compounds are generally ____________ at room temperature.c. Amount of energy depends on:i. the ___________ between ions.ii. the ___________ of the ions. **(This is has a larger effect)d. Coulomb’s Law: + r -e. Which compound has the highest lattice energy?(i) CaO or CaS? (ii) LiH or CaH 2 ?f. Why is the order of lattice energies: LiF > LiCl > LiBr > LiI?g. Which has a higher lattice energy: MgCl 2 or NaCl? Why?©RegliGower 107

D. Ionic compounds are generally __________________ in water, and the solution ___________electricity because the compounds are __________________.E. Molten (liquid) ionic compounds are also ___________ conductors of electricity, becausein the liquid state, charged particles (________) can flow.IV. Covalent BondingA. The ____________ of electrons between __________________.B. Structure:1. individual _______________ (e.g. H 2 O)2. 3-D lattices in network covalent bonds (e.g. _________________)C. Covalent compounds have attractive forces between atoms (______________) andattractive forces between molecules (________________________).D. Intermolecular forces between molecules are generally ___________, so covalentcompounds are generally in the _____________ or __________ phases, or possibly a lowmelting__________.E. Covalent compounds are generally ___________ conductors of electricity because noions are present.F. Types of covalent bonds:(shared pairs of electrons are shown as _____ or _____.)1. Single bond: _____ pair of shared electrons.Example: H 22. Double bond: _____ pairs of shared electrons.Example: O 23. Triple bond: ______ pairs of shared electrons.Example: N 2G. Bond length and energy:Bond Type N-N N=N N≡NBond Length (Å) 1.47 1.24 1.10Bond Energy (kJ/mol) 163 418 9411. As the number of bonds _____, the bond length ______ (______________ relationship)2. As the number of bonds _____, the bond energy ______ (_____________ relationship)3. Why can’t N 2 be directly used by plants? What forms of nitrogen can be usedas a plant nutrient?©RegliGower 108

V. ElectronegativityA. The ability of an atom to ______________ a shared pair of electrons toward itself.B. Elements with high _______ (accept electrons easily) and high _______ (do not loseelectrons easily) also have _________ electronegativity.C. The general trend of electronegativity is the same as IE: _______________D. Electronegativity is a relative concept and can only be measured relative to otherelements. Therefore it has no ___________, and the highest value was set at _______.E. Chart of electronegativities:F. Element with highest electronegativity: _______ Lowest electronegativity: _______VI. Bond PolarityA. Non-Polar Covalent (_____________ Covalent) Bond1. Electrons are shared ________________.2. The bond has 100% covalent character and 0% ___________ character.3. The electronegativity difference between atoms in the bond is __________.4. Example: H 2B. Polar Covalent Bond1. Electrons are shared _________________.2. The electronegativity difference between atoms in the bond is < _______.3. Example: HCl©RegliGower 109

C. Ionic Bond1. Electrons are transferred from one atom to the other.2. The electronegativity difference between atoms in the bond is ≥ _______.3. Example: LiFD. Examples: What type of bonds are each of the following:(a) KF (b) HBr (c) Br 2 (d) H 2 S (e) CsClVII. Ionic or Covalent CharacterA. Only bonds between identical elements are purely covalent. All others have some % ioniccharacter and some % _______________ character.B. Example from period 2 of the types of bonding between atoms:VIII. Oxidation Numbers (See page 129 in text):The charge an atom would have if electrons were completely transferred to the __________electronegative atom in the bond. The “relative” charge of an atom.A. Rules to assign oxidation numbers:1. In free elements, each atom has an oxidation state of _______.(i.e. H 2 , Br 2 , Li, etc.)2. For monatomic ions, the oxidation state is equal to the __________ on the ion.(i.e. Li+ has an oxidation state of +1)3. The oxidation state of oxygen in most compounds is ________, but in peroxidesits oxidation state is _________.4. The oxidation state of hydrogen when it is bonded to a non-metal is ______. Whenit is bonded to a metal its oxidation state is ______.(i.e. HCl = +1; LiH = -1)©RegliGower 110

5. Fluorine has an oxidation state of ______ in ALL its compounds. Other halogenscan have ____________ oxidation states when bonded to oxygen.6. In a neutral compound, the sum of the oxidation states is ______. In a polyatomicion, the sum of the oxidation states is the ___________ of the ion.B. Examples: Determine the oxidation states of each element in the following:LiH BeH 2 H 2 O HF OF 2 H 2 O 2CaCl 2 Cl 2 O 5 Cl 2 O 7 ClO 2 PbO 2 SiO 2NO 2 N 2 O NO 3-NO 2-MnO 4-C 2 O 42-IX. Periodic Table Relationships:Column## ValenceLewisdot formulaMost likely #of bonds# lone pairsIA 1IIA 2IIIA 3IVA 4VA 5VIA 6VIIA 7VIIIA 8X. Lewis Structures of Covalent Compounds: A representation of the covalent bonding in a molecule.A. Covalent bonds are shown as __________. Example: H 2B. Lone pairs of electrons are shown as _________. Example: O 2C. ONLY ____________ electrons are shown.©RegliGower 111

EXAMPLES:D. General steps for drawing Lewis structures:1. Sum the valence electrons in the compound.2. Add _____ for each negative charge. Subtract 1 for each _____ charge.3. Generally place the element that makes the _________________ number of bonds inthe center.4. Draw ____________ bonds to the other atoms off of the central atom.5. Place electrons around the ___________ atoms until an __________ is reached.6. If you run out of electrons, start forming _____________ or ___________ bonds.7. If you have EXTRA electrons after all have octets, place them on the __________ atom.8. In the end, all atoms should have an octet that need an octet (____ is an exception),and the total number of electrons should be placed on the molecule.1. CH 4 2. O 3 3. CO 2 4. CO 32-PRACTICE!!!1. HOCl 2. NF 3 3. TeCl 24. ClO 3-5. C 2 H 2 6. XeF 27. N 2 O 4 8. HOCN 9. PO 33-10. HNO 2 11. CH 2 O 12. CH 3 COOH©RegliGower 112

XI. Formal ChargeA. (Determined for each atom)1. Formal charge = ___________________________________.2. Use: To decide between _____________ Lewis structures. (Smaller formal charge=____________)3. For neutral compounds,____________________________________________________.4. For cations,_________________________________.5. For anions, _________________________________.6. Examples:CO 2Formal ChargeStructure 1:OCOStructure 2:Net Charge =OCOAmmonium ionNet Charge =Nall HNet Charge =Cyanate ionWhich structure is moreplausible? Why?XII. Resonance FormsA. Two or more equally ____________ structures of the same compound.B. Example: Ozone O 3C. Neither form alone accurately describes the _______________ of ozone.D. The bond lengths between each oxygen atom are the _________, which means that neitherbond is a ___________ bond or a ____________ bond, but a hybrid of the two.©RegliGower 113

D. Example: benzene (C 6 H 6 )XIII. Exceptions to the Octet RuleA. Odd number of ____________________1. Example: Nitrogen monoxide _________ Valence electrons:____________2. Free radical = ________________________! (SMOG)B. Incomplete octet (_________on central atom) : Elements with 3 or less valence electrons.1. BF 32. BeCl 23. HgCl 2C. Expanded octet (_________ on central atom): Occurs when (1)_________central atom and (2)surrounding atoms are very ___________________ (ex. ____, ____, ____)1. P2. SXIV. Coordinate Covalent Bonds: A covalent bond in which ____ atom provides both of the _________electrons.A. Ammonia (________) Ammonium ion (________)Formal charge of N?B. Water (_________) Hydronium ion (_________)Formal charge of O?©RegliGower 114

Chapter 10I. Molecular GeometryA. Ideal Geometries: The ____________ atom has NO lone pairs of electrons.TypeGeometryExampleCompoundExampleStructureBondAngleAX 2AX 3AX 4AX 5AX 6II. VSEPR Theory: Valence Shell Electron ___________ _________________ TheoryA. Double or ___________ bonds can be treated like single bonds.B. If a molecule has _________________ structures, VSEPR can be applied to any of them.C. Lone pairs of electrons repel atoms ____________ than bonding pairs. Therefore,lone pairs take up more __________ than atoms.TypeGeometryExampleCompoundExampleStructureBondAngleAX 3 EAX 2 E 2AXE 3©RegliGower 115

MEMORIZE THE SHAPES FOR THE TEST!III. Polar Molecules: Molecules with a ________________ moment (dipole arrows don’t cancel out!)A. Determination of molecular polarity1. Draw Lewis structure.2. Determine ______________ of molecule.3. Determine polarity of ____________.4. Draw dipole arrows.5. Determine net dipole (_________________________________________________).©RegliGower 116

B. Examples:1. CO 2 Type = _____________ Shape = _____________________2. H 2 O Type = _____________ Shape = _____________________3. BF 3 Type = _____________ Shape = _____________________4. CH 2 F 2 Type = _____________ Shape = _____________________5. XeF 2 Type=______________ Shape = ____________________6. NH 3 Type = _____________ Shape = _____________________©RegliGower 117

IV. Hybridization of Atomic Orbitals:The _________________ of atomic orbitals in an atom (generally the_______________ atom) togenerate a set of _______________ orbitals.A. Hybrid orbitals: a combination of _______ or more atomic orbitals.Example: Carbon generally forms ____ bonds, but only has ____ lone electrons. Why??Electron configuration of carbon:Atomic orbitals of valence electrons:Hybridization of carbon:When carbon has four single bonds, it is _________ hybridized.B. Procedure for Hybridizing Atomic Orbitals:1. Hybridization of orbitals only applies to _______________.2. Hybridization is the mixing of at least _______ or more nonequivalent atomic orbitals.Hybrid orbitals have different _________ than the atomic orbitals:Atomic OrbitalsHybrid Orbitals:s p sp sp3. The number of hybrid orbitals is equal to the number of pure _________ orbitals.4. Hybridization requires an input of _____________. However, this energy is more thanrecovered during __________ formation (an ________________ process).5. Covalent bonds are formed by an __________________ of _____________ orbitals withatomic orbitals:C. Example hybridizations:Hybrid# SingleBondsExample Atomic Orbitals Hybrid Orbitals Shapespsp 2sp 3sp 3 dsp 3 d 2©RegliGower 118

D. Example: Predict the hybridization of the central atom of each of the following:1. HgCl 2 2. AlI 3 3. PF 3E. Hybridiztion of s, p, and d orbitals: For elements in the _____________ period and beyond,the s, p, and d orbitals all contribute to hybridization. Expanded ____________ can beexplained by the hybridization of the d orbitals. Because there is no ________ level,elements of the ____________ period cannot form expanded octets.1. Example: Phosphorus in PBr 52. Example: Sulfur in SF 6V. Diagram of Hybrid Orbitals©RegliGower 119

VI. Hybridization in Molecules Containing Double and Triple Bonds:A. Types of covalent bonds:1. Sigma bond (_______): The bonds formed when the ______________ orbitals overlapwith the atomic orbitals.2. Pi bond (_______): The bonds formed when ______ orbitals overlap above and belowthe plane.Example: C 2 H 4B. Types of Hybridization for Carbon:Bonding Hybrid Atomic Orbitals Hybrid OrbitalsSigma/PibondsC. Example: Determine the hybridization of each carbon and the shape of each molecule:(1) CH 2 O, (2) C 2 H 2.©RegliGower 120