Welcome to the second term of Regents Chemistry at LaGuardia! If you have come this far, then you have survived the first term of Chemistry!
You'll find all of your class notes and homework assignments on this web page to assist you when you are studying for this class. Class notes are listed below. Homework assignments are listed in the Flashcards section.
If you need tutoring, check the Science Department bulletin boards for tutoring times.
Here's to another great semester and to all of you passing the Regents with flying colors!
Units Covered in Term 2:
1. Math of Chemistry
2. Stoichiometry
3. Solutions
4. Acids & Bases
5. Kinetics & Equilibrium
6. Redox & Electrochemistry
7. Organic
LaGuardia Chemistry Grading Policy
1. Exams 60 %
a. Midterm Counts as 2 Exams
b. Final Counts as 2 Exams
c. Ms. Pace's Policy: Lowest Exam Grade For Term Will Be Dropped
2. Homeworks 40 %
a. Late HW will have a penalty of 10 points for every day the assignment is not turned in.
Classroom Policy:
1. No Talking while the teacher is talking.
2. No Talking while other students are asking or answering questions.
3. Students should come to class on time. If a student is late they should quietly and discreetly sign the late book so they are not marked as cutting for the day.
4. No Graphing Calculators can be used on any exam. (Science Dept. Policy)
5. No Writing on Reference Tables used for exams.
6. Cheating will not be tolerated and disciplinary action, outlined in the LaGuardia High School Student Handbook, will be taken.
7. Handouts/Worksheets will be given out once. You are responsible for any additional copies if you lose your original.
8. Your name must be written on anything and everything handed in. If your name is not on your paper, (HW, Lab, Exam), then your paper will be discarded.
9. Tutoring is available. Please check the Science Department bulletin boards for tutoring days and times.
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apace2@schools.nyc.gov
Subject Line: Name, Chemistry Period ____
Body of Email: Name, Class Period, I agree with the Course Requirements as stated on our class website: www.schoolnotes.com/10023/mspace2.html.
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Week of 1/30/08:
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Lesson #1 Math of Chemistry January 31 - February 1, 2008
Aim: What is the Math of Chemistry?
I. Gram Formula Mass
a. Gram Atomic Mass - Mass of an atom in grams
i.e. Na = 22.98977 grams per mole = 23g/mol
b. Gram Formula Mass - Mass of a compound
i.e. sodium chloride - NaCl = 23g/mol + 35g/mol = 58g/mol
i.e. magnesium chloride - MgCl2 = 24g/mol + 2(35)g/mol = 94g/mol
i.e. magnesium nitrate - Mg(NO3)2 = 24g/mol + 2(14)g/mol + 6(16)g/mol = 148g/mol
II. Percent Composition - represents the composition as a percentage of each element in a compound compared with the total mass of the compound
Determine the %comp of each of the following compounds:
% Composition by Mass = Mass of Part/Mass of Whole X 100%
1. KMnO4
K = 25%
Mn = 35%
O = 40%
2. HCl
H = 2.8%
Cl = 97.2%
3. Mg(NO3)2
Mg = 16%
N = 19%
O = 65%
III. Mole Calculations
1. What is a mole?
a number, 6.02 x 10e23, 22.4L of a gas at STP, etc.
2. What is Avogadro's Number?
6.02 x 10e23
3. Why do we use moles in Chemistry?
to count particles, as a conversion factor between # particles and grams, etc.
4. If a sample contains 100g of magnesium chloride (MgCl2), then how many moles of magnesium chloride are in the sample?
# Moles = Given Mass (g)/ Gram Formula Mass (g/mol)
# Moles = 100g / 95 g/mol
# Moles = 1.05 mol
5. What is the mass of 2.3 moles of sodium sulfate (Na2SO4)?
2.3 moles = (x)/(142g/mol)
x = 326.6g
IV. Molar Volume - the volume occupied by 1 mole of any gas at STP = 22.4L
1. If a balloon contains 11.2L of He(g), then how many moles of He are in the balloon at STP?
If 22.4L = 1 mole of any gas, then 11.2L of He(g) = 0.5 moles
2. How many liters of Ne(g) are in 2.5 moles at STP?
2.5 moles x 22.4 L/mol = 56L
3. How many moles of Ar(g) are in 107.9 L at STP?
107.9L / 22.4 L/mol = 4.8 mol
4. How much volume does 32g of CO2(g) occupy at STP?
x = 32g / 44g/mol
x = 0.727mol
0.727mol x 22.4 L/mol = 16.28 L
V. Percent Error - indicates how closely the measurement agrees with an accepted value of the same quantity
% Error = ([Observed Value - Accepted Value]/Accepted Value)x 100%
(Note: "[ ]" represent absolute value bars. We are only interested in the magnitude of the percent error and not the algebraic sign.)
1. Observed Value - value from the experiment
2. Accepted Value - value from an accepted measurement of that quantity (from a book, scientist, teacher, etc.)
3. A student measures the density of an object to be 5.6 grams per liter. The accepted density of the object is 6.4 grams per liter. What is the percent error of this measurement?
([5.6g/L - 6.4g/L]/6.4g/L) x 100% = 12.5%
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Additional Math of Chemistry Practice Problems:
1. How much does 2 moles of lead(II)nitrate weigh?
Pb(NO3)2 = 331 g/mol
2 mol = X / 331 g/mol
X = 662 mol
2. How many moles are in 50g of ammonium thiocyanate?
NH4SCN = 76 g/mol
X = 50g / 76 g/mol
X = 0.65 mol
3. What is the percent composition by mass of copper in copper(II)chloride?
CuCl2 = 134 g/mol
% comp = 64 g/mol / 134 g/mol x 100%
= 0.48 x 100%
= 47.8%
4. Calculate the gram formula mass and the percent composition by mass of each element in the compound, I2.
254g/mol
I = 100%
5. Calculate the gram formula mass and the percent composition by mass of each element in the compound, calcium nitrite (Ca(NO2)2).
132g/mol
Ca = 30.3%
N = 21.2%
O = 48.5%
6. Calculate the gram formula mass and the percent composition by mass of each element in the compound, ammonium sulfate ((NH4)2SO4).
132g/mol
N = 21.2%
H = 6.1%
S = 24.2%
O = 48.5%
7. How many moles of nickel (III) carbonate will have a total mass of 165g?
8. What is the total mass of 0.75 mol of manganese (IV) oxide (MnO2)?
9. How many atoms are in 2.5 moles of Carbon?
2.5 mol x 6.02 x 10e23 = 15.05 x 10e23 or 1.505 x 10e24
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Converting with Moles Worksheet
1. In a chemical reaction 0.397 mole of ethyl chloride (C2H5Cl) is produced. What is the mass in grams of this amount of ethyl chloride?
2. A small bottle in the chemistry stockroom contains 43.35g of nickel (II) carbonate, NiCO3. How many moles of NiCO3 is this?
3. What is the equivalent in moles of 135L of ammonia (NH3) gas?
4. A nurse has been asked to get 0.0465 mole of quinine (C20H24N2O2). What mass of quinine should the nurse obtain?
5. During an electroplating process, 5.8625g of silver is deposited on a steel bar. How many moles of silver is this?
6. A helium-filled balloon has a total volume of 136,500 L at STP. How many moles of helium are in the balloon?
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Lesson #2 Empirical Formulas February 5-6, 2008
Aim: What is the difference between Molecular and Empirical Formulas?
1. Molecular Formula - The formula for a compound that exists as a molecule.
2. Empirical Formula – The formula that gives only the simplest whole-number ratio of the atoms that make up the compound.
i.e. Molecular/Chemical Formula = C6H6
Empirical Formula = CH
3) Empirical Formulas help us identify classes of compounds. For instance, all sugars like glucose (C6H12O6), have the same empirical formula CH2O, which is the reason why they are called carbohydrates.
4) How To Derive An Empirical Formula
i.e. A compound containing carbon, hydrogen, and oxygen is found to contain 9.1% hydrogen and 54.5% carbon. What is its empirical formula?
a.Since the percentage of an element is the number of grams per 100g of compound, we may assume a 100g sample of compound and convert the percents to grams directly.
b.Since the percent hydrogen and percent carbon do not add up to 100%, we conclude that the missing 36.4% is oxygen.
C = 54.5g H = 9.1g O = 36.4g
c.Calculate the number of moles of each element using the mole calculation equation.
C = 4.54 mol H = 9.1 mol O = 2.28 mol
The empirical formula is now: C4.54H9.1O2.28
d.Divide each result by the smallest number.
C 4.5/2.28 H 9.1/2/28 O 2.28/2.28
The empirical formula is now: C1.99H3.99O1
e.Round the numbers to the nearest whole number for the subscripts of the elements in the empirical formula of the compound.
The empirical formula is now: C2H4O
5) How To Derive A Molecular Formula From An Empirical Formula
i.e. A compound has the empirical formula CH2O, and its molar mass is determined in a separate experiment to be 180 grams per mole. What is the molecular formula of this compound?
a. Use the equation: Molecular Gram Formula Mass/Empirical Formula Gram Formula Mass = Empirical Formula Units
b. 180 g/mol / 30 g/mol = 6 empirical formula units
c. C 1x6 H 2x6 O 1x6 = C6H12O6 (glucose)
6) Practice Problems (Empirical Formulas)
a. A compound is analyzed and found to contain 36.70% potassium, 33.27% chlorine, and 30.03% oxygen. What is its empirical formula?
b. Determine the empirical formula of the compound that contains 17.15%C, 1.44%H, and 81.41%F.
c. The active ingredient in a photographic fixer solution contains sodium, sulfur, and oxygen. Analysis of a sample shows that the sample contains 0.979g Na, 1.365g S, and 1.021g O. What is the empirical formula of this compound?
7) Practice Problems (Molecular Formulas)
a. The empirical formula of a compound is NO2. Its molecular mass is 92 g/mol. What is its molecular formula?
b. The empirical formula of a compound is CH2. Its molecular mass is 70 g/mol. What is its molecular formula?
c. A compound is found to be 40.0%C, 6.7%H, and 53.3%O. Its molecular mass is 60.0 g/mol. What is its molecular formula?
d. A compound is found to be 64.9%C, 13.5%H, and 21.6%O. Its molecular mass is 74.0 g/mol. What is its molecular formula?
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Lesson #3 Equations February 7, 2008
Aim: What is a Chemical Equation?
I. Chemical Equation/ Reaction – One or more substances (the reactants) change into one or more new substances (the products)
a. Reactant --> Products [What does the arrow mean?]
b. Iron + Oxygen --> Iron (III) Oxide (Word Equation)
c. Fe + O2 --> Fe2O3 (Chemical Equation)
II. Symbols Used In Chemical Equations
a. + = used to separate two reactants or two products
b. --> = “yields”, separates reactants from products
c. <--> = used in place of --> for reversible reactions
1. ---> forward reaction
2. ---> reverse reaction
d. (s) = solid
e. (l) = liquid
f. (g) = gas
g. (aq) = aqueous solution, a substance that is dissolved in water
h. heat --> = indicates that heat is supplied to the reaction
i. catalyst --> = indicates that a catalyst was used in the reaction to change the reaction speed (rate)
III. Write the unbalanced chemical equation from the following word equations:
a. Solid Iron combines with Gaseous Chlorine to yield Solid Iron (II) Chloride
b. Liquid Hydrogen Peroxide decomposes to produce Liquid Water and Gaseous Oxygen
c. Solid Zinc combines with Aqueous Copper (II) Sulfate to produce Aqueous Zinc Sulfate and Solid Copper
d. A water solution of Barium Nitrate combines with a water solution of Sodium Sulfate to produce Solid Barium Sulfate and a water solution of Sodium Nitrate
e. Gaseous Hydrogen combines with Gaseous Oxygen to yield Liquid Water
f. Solid Potassium Bromide combines with Gaseous Fluorine to yield Solid Potassium Fluoride and Gaseous Bromine
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Lesson #4 Classifying Chemical Rxns February 8, 2008
Aim: How do I classify chemical reactions?
I. Direct Synthesis Reaction (Combination Reaction)–
2 or more reactants form one product
Fe (s) + Cl2(g) --> FeCl2 (s)
II. Decomposition Reaction – breakdown of a single reactant into two or more products
2H2O2 (l) --> 2H2O (l) + O2 (g)
hydrogen peroxide --> water and oxygen
III. Single-Replacement Reaction (Single-Displacement Reaction)– an uncombined element replaces another element that is part of a compound
Zn(s) + CuSO4(aq) --> ZnSO4(aq) + Cu(s)
IV. Double-Replacement Reaction (Double-Displacement Reaction)– two elements in 2 different compounds replace each other
Aim: How do I balance a Chemical Equation using smallest, whole-number coefficients?
I. The Law of Conservation of Matter = Matter can neither be created nor destroyed only changed from one form to another.
- Therefore when we balance a chemical equation we are making sure to follow this law. There must be the same number of atoms on the left side of the equation as there are on the right side.
(This is another type of reaction, called combustion (aka respiration), that we will study more during Term 2 when we cover Organic Chemistry.)
Left side of equation: 1 carbon atom, 4 hydrogen atoms, 2 oxygen atoms
Right side of equation: 1 carbon atom, 2 hydrogen atoms, 3 oxygen atoms
- We must use coefficients to balance equation. I.e. 5 NH3 means there are 5 molecules of ammonia or 5 MOLES of ammonia.
- Coefficients count the number of moles of an atom/molecule/compound that are present.
1. Balance equation for hydrogen atoms
CH4(g) + O2(g) -->CO2 (g) + 2H2O(l)
2. Balance equation for oxygen
CH4(g) + 2O2(g) --> CO2 (g) + 2H2O(l)
The equation is now balanced with smallest, whole-number coefficients.
What is the sum of the coefficients of the above equation? 6
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Lesson #5a Balancing Equations Practice
Aim: Balancing Chemical Reactions Practice Problems
I. Balance the following chemical reactions using smallest, whole-number coefficients.
II. Name the reaction type for the above reactions.
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Lesson 5b Stoichiometry Review
Aim: Stoichiometry Review for Exam #1.
1. Iron reacts with oxygen gas to produce iron(III)oxide. When this occurs slowly, it is known as rusting.
Fe + O2 --> Fe2O3
a) Given 25 moles of Fe, how many moles of O2 gas would be needed to completely react with this iron?
1. Balance the equation above using smallest, whole-number coefficients: 4Fe + 3O2 --> 2Fe2O3
2. Set up a relationship between Fe and O2 in this equation and Fe and O2 in the new equation in which there are 25 moles of Fe instead of 4 moles:
25 moles of Fe/4 moles of Fe = x moles of O2/3 moles of O2
3. x moles of O2 = 18.75 moles of O2
b) Given 25 moles of iron, determine the number of moles of Fe2O3 that would be produced.
1. Set up a relationship between Fe and Fe2O3 in this equation and Fe and Fe2O3 in the new equation in which there are 25 moles of Fe instead of 4 moles:
25 moles of Fe/4 moles of Fe = x moles of Fe2O3/2 moles of Fe2O3
2. x moles of Fe2O3 = 12.5 moles of Fe2O3
c) Determine the mass of Fe2O3 produced in part b.
# moles = given mass (g) / Gram Formula Mass (g/mol)
12.5 moles Fe2O3 = x / 160g/mol
x = 2000g Fe2O3 = 2kg Fe2O3 = 4.4 pounds of rust
d) Name the reaction type for this reaction.
Direct Synthesis, Synthesis, Combination
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Lesson #6 Mixtures February 13-14, 2008
Aim: What is the difference between a homogeneous solution and a heterogeneous mixture?
Motivation: Video - Solutions (Ionic & Molecular)
I. Mixtures = made up of 2 or more substances
1. Solute = substance that is dissolved or substance in smaller quantity
2. Solvent = substance doing the dissolving or substance in larger quantity
3. Examples:
Solute Solvent Examples
Solid Solid Metallic Alloy (Steel = Mn,V,W in Fe)
Solid Liquid Saline Solution (NaCl(aq))
Liquid Liquid Ethyl alcohol + water (or Oil in CCl4)
Gas Liquid CO2 + water (Soda = CO2(aq))
Gas Gas Air (N2 =solvent; O2,H2,etc. =solutes)
4. The addition of a nonvolatile solute to a solvent causes the boiling point of the solvent to increase and freezing point of the solvent to decrease.
a. Greater concentration of solute = greater effect on colligative properties of a solvent.
II. Types of Mixtures
A. Homogeneous – evenly distributed components in a mixture and particles do not settle out
i.e. solutions, saline solution (If you equally divide solution in Dixie cups to class, everyone will have same amount of salt to water ratio.)
B. Heterogeneous – unevenly distributed components in a mixture and particles will settle out
i.e. suspensions, salad (If divide onto plates, little Mary may not get a tomato and little Ezzy might not get an olive. Therefore, cannot divide equally.)
1. Colloidal Dispersions = heterogeneous mixture where particles are larger than those in solutions but smaller than those in suspension
i.e. Jello gelatin, smoke, aerosol sprays, whipped cream, butter, mayonnaise, milk, marshmallow, cheese
a. exhibit the "Tyndall Effect" = Light is scattered by particles in a substance forming a beam of light.
C. Solvation = dissolving / dissociation of solute in solvent
D. Hydration = solvation when solvent is water
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Lesson #7 Solubility February 15, 2008
Aim: What is Solubility?
I. Types of Solutions
1. Miscible = infinitely soluble
i.e. gases into gases
2. Unsaturated = solution has not reached its saturation point and can therefore hold more dissolved solute
3. Saturated = equilibrium between dissolved and undissolved solute
= the solvent has reached maximum holding capacity and any more solute added to the solvent will not dissolve
4. Super-saturated = dissolve more solute into solvent, passed the saturation point, when solvent is heated (for solid solutes), or when solvent is cooled (for gas solutes)
i.e. Put a Honey Bear in the fridge for a few days: Honey = eventually super-saturated honey solution will have the excess sugar crystallize at the bottom of the honey jar after a long period of time / to restore to store-bought condition = heat honey and then cool slowly
II. Solubility – the amount of solute added to a solvent that makes the solution saturated
A. Units = grams solute ÷ 100 grams of solvent
1. Solubility depends on temp and pressure of solute and solvent
Therefore: How Much solute dissolves, NOT How Fast solute dissolves
i.e. Stirring sugar in Iced Tea will make sugar dissolve faster, BUT will not change how much sugar Iced Tea can hold in solution.
i.e. Why does a soda become “flat” if opened and left unrefrigerated?
B/c High Temp and Low Pressure reduce solubility of CO2 which allows it to escape from liquid. (Solubility decreases.)
2. Solubility also depends on nature of Solute and Solvent
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Lesson #8 Reference Table G February 25, 2008
Aim: How do I use Reference Table G: Solubility Curves?
I. Solubility Curves – trace solubility of a substance with changing temp
1. Reference Table G - shows saturation points of specific solutes in 100g of water at certain temperatures
i.e. At 20°C, 38 grams of NaCl in 100g of water will form a saturated solution
Therefore: Solubility of NaCl @ 20°C = 38g NaCl per 100g water
At 20°C, 37 grams of NaCl in 100g of water will form an unsaturated solution
At 20°C, 39 grams of NaCl in 100g of water will form a supersaturated solution
At 20°C, 76 grams of NaCl in 200g of water will form a saturated solution
At 20°C, 19 grams of NaCl in 50g of water will form a saturated solution
2. Solids as solutes = generally : If Temp Increases, then Solubility of solid increases.
3. Gases as solutes = If Temp Increases, then Solubility of gas decreases. (And Gas escapes back into air.)
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Lesson #9 Reference Table F February 26, 2008
Aim: How do I use Reference Table F: Table of Solubilities in Water?
I. Reference Table F = Table of Solubilities in Water
This table is primarily used to determine the precipitate from a chemical reaction such as:
BaCl2 + Na2SO4 --> 2 NaCl + BaSO4 (precipitate)
1. Precipitate - substance that is insoluble in solution
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Lesson #10 Molarity February 29, 2008
Aim: How can I use Molarity to find the concentration of a solution?
I. Molarity
1. involves moles
2. shows the concentration of a solution
3. M = moles of solute/liters of solution (Table T)
4. Concentrated - strong, little solvent & lotsa solute, high molarity
5. Dilute - weak, watered down, lotsa solvent & little solute, low molarity
Example #1: Suppose we had 1.00 mole of sucrose (it's about 342.3 grams) and proceeded to mix it into some water. It would dissolve and make sugar water. We keep adding water, dissolving and stirring until all the solid was gone. We then made sure that when everything was well-mixed, there was exactly 1.00 liter of solution.
What would be the molarity of this solution?
Molarity = 1.00 mol sucrose/1.00 L soln
The answer is 1.00 mol/L. Or “One point oh oh molar”
Example #2 - Suppose you had 2.00 moles of solute dissolved into 1.00 L of solution. What's the molarity?
Molarity = 2.00 mol/1.00 L
The answer is 2.00 M.
Example #3 - What is the molarity when 0.75 mol is dissolved in 2.50 L of solution?
Molarity = 0.75 mol/2.50 L
The answer is 0.300 M.
Example #4 - Suppose you had 58.44 grams of NaCl and you dissolved it in exactly 2.00 L of solution. What would be the molarity of the solution?
Step One: convert grams to moles.
Step Two: divide moles by liters to get molarity.
Step Three: Dividing 58.44 grams by 58.44 grams/mol gives 1.00 mol.
Then, dividing 1.00 mol by 2.00 L gives 0.500 mol/L (or 0.500 M).
Do examples #5 and #6:
5) Calculate the molarity of 25.0 grams of KBr dissolved in 750.0 mL.
6) 80.0 grams of glucose (C6H12O6, mol. wt = 180. g/mol) is dissolved in enough water to make 1.00 L of solution. What is its molarity?
Practice Problems
1) Calcuate the molarity when 75.0 grams of MgCl2 is dissolved in 500.0 mL of solution.
2) 100.0 grams of sucrose (C12H22O11, mol. wt. = 342.3 g/mol) is dissolved in 1.50 L of solution. What is the molarity?
3) 49.8 grams of KI is dissolved in enough water to make 1.00 L of solution. What is the molarity?
HW #10: Molarity Worksheet #1-10, Due Wednesday, 3/21/07
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Lesson #11 Dilutions of Solutions March 3, 2008
Aim: How do I calculate molarity for a diluted solution?
1. 20mL of a salt solution has a concentration of 0.60M. If the solution is diluted to 40mL, then what is the new concentration?
M1V1 = M2V2
(0.60M)(20mL) = (x)(40mL)
x = 0.3M
2. 30mL of a 7M magnesium chloride solution is diluted. If the new solution’s concentration is 2M, then how much solvent was added?
3. A solution contains 5 moles of lithium nitrate in 200mL of solution. The new volume of the solution after the dilution is 300mL. What is the new concentration?
5moles/0.200L = 25M
(25M)(200mL) = (x)(300mL)
x = 16.7M
4. A solution contains 160g of potassium iodide in 200mL of solution. The new volume is 500mL after a dilution. What is the new concentration?
(160g)/(166g/mol) = 0.96 moles
0.96mol/0.200L = 4.8M
(4.8M)(200mL) = (x)(500mL)
x = 1.92M
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Lesson #12 Parts Per Million March 4, 2008
Aim: How do I calculate parts per million for a solution?
I. Parts Per Million (ppm) - a unit of measurement used to express the concentration of very dilute solutions.
1. "per cent" means out of a hundred,
"parts per million" (ppm) means out of a million
2. Reference Table T:
ppm = grams of solute/grams of solution x 1,000,000
(grams of solution = grams of solute + grams of solvent)
II. Practice Problems
1. If 0.025 grams of NaI are dissolved in 100 grams of water, then what is the concentration of the resulting solution, in parts per million?
ppm = 0.025g/100.025g x 1,000,000 = 250 ppm
2. What is the concentration of a solution in parts per million if 0.02 grams of LiCl are dissolved in 1000 grams of water?
ppm = 0.02g/1000.02g x 1,000,000 = 20 ppm
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Lesson #13 Acids & Bases Introduction March 5, 2008
Aim: What is the difference between acids and bases?
Motivation: Video - Acids, Bases, & Salts
I. Common Properties of Acids and Bases
1. Acids and Bases, when ionized, conduct electricity in water and are called “electrolytes”
2. dilute acids taste sour, dilute bases taste biter
3. aqueous solutions of bases feel slippery
4. acids + bases yield a neutral solution
HCl + NaOH --> NaCl + H2O
5. WATER is amphiprotic or amphoteric b/c it can behave as an acid in one reaction and a base in another reaction.
6. Acids, Bases, and Salts are all ionic compounds
II. The Arrhenius Acid Base Theory
1. Acid - any substance which delivers a hydrogen ion (H+), a hydronium ion (H3O+), a proton, to the solution.
Here is a generic acid dissociating, according to Arrhenius:
HA ---> H+ + A¯
HCl --> H+ + Cl-
a. see Reference Table K
b. oxides of nonmetals are considered acids
2. Base - any substance which delivers a hydroxide ion (OH¯) to the solution.
This would be a generic base:
XOH ---> X+ + OH¯
NaOH ----> Na+ + OH-
a. see Reference Table L
b. oxides of metals are considered bases i.e. CaO (lye)
*Remember: CH3OH is an alcohol, NOT a base! Any hydroxide group attached to a hydrocarbon will NOT produce a base.
3. When acids and bases react according to this theory, they neutralize each other through a double displacement reaction, forming water and a salt:
HA + XOH ---> H2O + XA
Keeping in mind that the acid, the base and the salt all ionize, we can write the following:
H+ + A¯ + X+ + OH¯ ---> H2O + X+ + A¯
Finally, we can drop all “spectator ions”, to get the following:
H+ + OH¯ ---> H2O
i.e. HCl + NaOH ---> H2O + NaCl
4. This covers hydrochloric acid, acetic acid, and so on and most of the bases like sodium hydroxide, potassium hydroxide, calcium hydroxide and so on. HOWEVER, Arrhenius’ theory did not explain why ammonia (NH3) was a base.
The Arrhenius theory of acids and bases will be fully supplanted by the theory proposed independently by Johannes Brønsted and Thomas Lowry in 1923.
III. Identify: Acid or Base?
1.HBr acid
2.NH4OH base
3.LiOH base
4.H2SO4 acid
5.Mg(OH)2 base
6.HI acid
7.HF acid
8.Ca(OH)2 base
9.KOH base
10.HNO3 acid
*11.NH3 BASE!!! (Please see Lesson #14 for explanation.)
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Lesson #14 Bronsted-Lowry Acid/Base Theory March 10, 2008
Aim: What is the Bronsted-Lowry Acid/Base Theory?
I. The Bronsted-Lowry Acid/Base Theory
1. An acid is a substance from which a proton can be removed.
a. An acid is a "proton donor."
2. A base is a substance that can remove a proton from an acid.
b. A base is a "proton acceptor."
3. proton - H+ ion, a hydrogen atom that loses its valence electron is just a proton
*Remember: proton, hydrogen ion, H+, H3O+ all mean the same thing
II. Examples:
1. In the reaction: NH3 + HBr --> NH4 + + Br- , identify the Bronsted-Lowry acid and the Bronsted-Lowry base.
HBr “donates” a proton to NH3, therefore:
HBr is the Bronsted-Lowry acid
NH3 is the Bronsted-Lowry base
2. In the reaction: H2O + NH2 -1 --> NH3 + OH -1, identify the Bronsted-Lowry acid and the Bronsted-Lowry base.
H2O is the Bronsted-Lowry acid
NH2- is the Bronsted-Lowry base
III. Identify: Acid or Base?
1.NH4 +1 acid b/c will be a proton donor
2.CO3 -2 base b/c will be a proton acceptor
3.NO2 -1 base b/c will be a proton acceptor
4.H30 +1 acid b/c will be a proton donor [aka Hydronium Ion]
*5.HS -1 acid & base (This substance is amphiprotic/amphoteric like water because it can both donate a proton and accept a proton, respectively.)
IV. Lewis Acids & Bases
1. Lewis Acid - electron pair acceptor
2. Lewis Base - electron pair donor (All Lewis Bases are exactly the same as Bronsted-Lowry Bases.)
Side Notes:
An acid-base reaction consists of the transfer of a proton from an acid to a base. KEEP THIS THOUGHT IN MIND!!
In an acid, the hydrogen ion is bonded to the rest of the molecule. It takes energy (sometimes a little, sometimes a lot) to break that bond. So the acid molecule does not "give" or "donate" the proton, it has it taken away. In the same sense, you do not donate your wallet to the pickpocket, you have it removed from you.
The base is a molecule with a built-in "drive" to collect protons. As soon as the base approaches the acid, it will (if it is strong enough) rip the proton off the acid molecule and add it to itself.
Now this is where all the fun stuff comes in that you get to learn. You see, some bases are stronger than others, meaning some have a large "desire" for protons, while other bases have a weaker drive. It's the same way with acids, some have very weak bonds and the proton is easy to pick off, while other acids have stronger bonds, making it harder to "get the proton."
**When an acid and a base come together they form a salt plus water.
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Lesson #15 pH & pOH March 13, 2008
Aim: What is pH (& pOH)?
I. pH = tells us the concentration of an acid or a base, based on the H+ ion concentration
i.e. If the [H+] = 0.0010 M = 1.0 x 10e-3, then the pH = 3.
1. pH was invented by a chemist in a soft drink bottling plant who was getting tired of writing concentrations on each bottle and started using an abbreviation:
p = for power of ten H = for H+ ion concentration
2. Find the pH for the following [H+]:
a. 0.1M pH 1
b. 0.000001M pH 6
c. 0.0001M pH 4
d. 0.00000000000001M pH 14
e. 1M pH 0
3. Find the [H+] for the following pH values:
a. pH 13 1.0 x 10e-13 M
b. pH 7 1.0 x 10e-7 M
c. pH 2 1.0 x 10e-2 M
d. pH -2 1.0 x 10e2 M = 100M
4. The degree of ionization determines the strength of acids and bases. Therefore, the more ions, the stronger the solution.
II. pH Scale
<--(0)---------------(7)-----------------(14)-->
acid neutral base
III. [Non-Regents] pOH = tells us the concentration of an acid or a base, based on the OH- ion concentration
i.e. If the [OH-] = 0.0010 M, then the pOH = 3.
IV. [Non-Regents] pOH Scale
<--(0)---------------(7)-----------------(14)-->
base neutral acid
1. Find the pOH for the following [OH-]:
a. 1 x 10e-2M pOH 2
b. 0.0000001M pOH 7
c. 1 x 10e-10M pOH 10
d. 0.01M pOH 2
V. [Non-Regents] pH + pOH = 14, [H+]x[OH-] = 1.0 x 10e-14
[Why does water have a neutral pH of 7?]
The pH of a solution plus the pOH of the same solution, when added together, will equal 14, ALWAYS!
i.e. #1) If the pH of a solution = 2, then what is the pOH of the same solution? pOH = 12 b/c 2 + X = 14
i.e. #2) If the [OH-] of a solution = 1.0 x 10e-4M, then what is the [H+] of the same solution? [H+] = 1.0 x 10e-10M b/c X x 1.0 x 10e-4M = 1.0 x 10e-14
i.e. #3) If the [H+] of a solution = 1.0 x 10e-11M, then what is the pOH?
The [OH-] = 1.0 x 10e-3, therefore the pOH = 3
VI.Practice Problems
1. If the [OH-] = 0.000010 M, then what is the pH of the solution?
[OH-] = 1 x 10e-5 M, therefore the pOH = 5
14 - 5 = pH
pH = 9
[H+]= 1 x10e-9 M, therefore [H+] = 0.0000000010 M
2. Find the pH, pOH, [H+], and [OH-] for the following:
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Lesson #16 Titrations March 14, 2008
Aim: How can I find the unknown molarity of an acid or a base?
I. Titration = the process used to find the unknonwn molarity of an acid or a base by adding measured volumes of:
1. an acid of known molarity to a base of unknown molarity,
or
2. a base of known molarity to an acid of unknown molarity,
until the endpoint (equivalence point) or neutralization (pH 7) is reached. An indicator will tell us when this happens.
Moles of acid = Moles of base
Moles of H+ = Moles of OH-
II. MaVa = MbVb
[M (acid) x (# of H+)] x mL (acid) = [M (base) x (# of OH-)] x mL (base)
III. Examples
1. How many mL of 4.00M NaOH are required to exactly neutralize 50.0 mL of a 2.00M solution of HCl?
a. 25.0 mL b. 50.0 mL c. 100.0 mL d. 200.0 mL
HCl NaOH
mL acid x M acid = mL base x M base
50 x 2 = ? x 4
100 = 4X
X = 25 mL (the mL of base needed to neutralize the acid)
(If the acid was H2SO4, then you would have to multiply the Molarity of the acid by 2 b/c of the 2 H+ ions in H2SO4.)
2. In an acid-base titration, standard 0.10 M KOH is used to determine the concentration of an unknown HNO3 solution.
a. Write an equation for the neutralization reaction.
HNO3 + KOH --> H2O + KNO3
b. Name the salt produced.
KNO3 = potassium nitrate
c. If 50 mL of KOH is needed to neutralize 200 mL of HNO3,then what is the concentration of the acid?
mL acid x M acid = mL base x M base
(200 mL) x M acid = (0.10 M) (50.0 mL)
M acid = 0.025 M
IV. Indicators - substances that indicate the degree of acidity or basicity of a solution through characteristic color changes
a. see Reference Table M
b. Litmus is a dye made from lichen plants and is an excellent indicator used for identifying both acids & bases.
i.e. If bromthymol blue is used as an indicator in a titration, where a solution of known molarity of a base is added to a solution of unknown molarity of an acid, then what will the color change be?
From yellow to blue.
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Lesson #17 Salts March 18, 2008
Aim: What is a Salt?
I. Salt
1. ionic compound
(metal cation + nonmetal anion)
i.e. NaCl (Remember, NaCl is NOT the ONLY salt in the world!)
2. not an acid or a base
3. may be formed as a product by the neutralization reaction of an Arrhenius acid and an Arrhenius base
4. form electrolytes in solution
II. Is It A Salt? Yes or No
1. CH4 - No 6. MgI2 - Yes 11. Ba(NO3)2 - Yes
2. CaCO3 - Yes 7. NH3 - No 12. Cu3(PO4)2 - Yes
3. LiOH - No 8. H2SO4 - No 13. FeBr2 - Yes
4. KBr - Yes 9. RbF - Yes 14. Cu(OH)2 - No
5. HCl - No 10. H2O - No 15. WF6 - Yes
III. Practice Problems
1. Write the equation for the neutralization of chlorous acid by potassium hydroxide. Name the salt produced.
HClO2 +KOH --> KClO2 + H2O salt = potassium chlorite
2. An acid-base neutralization produces calcium nitrate as the salt. What are the names of the acid and base?
Ca(OH)2 + 2HNO3 --> Ca(NO3)2 + 2H2O
The acid is nitric acid. The base is calcium hydroxide.
3. What are the name and the formula of the salt produced by the reaction of magnesium hydroxide and chloric acid?
Magnesium chlorate; Mg(ClO3)2
4. Write the products of the following neutralization reactions, balance the equations, and name the salts produced.
a. H2SO4 + KOH -->
b. H2SO4 + Ba(OH)2 -->
c. H3PO4 + Ca(OH)2 -->
d. H3PO4 + NaOH -->
e. HNO3 + Mg(OH)2 -->
f. HNO3 + Al(OH)3 -->
5. Name the following salts.
a. NaHCO3 d. Na2SO4
b. Na2CO3 e. KCl
c. NH4NO3 f. NH4Br
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Lesson #18 Potential Energy Diagrams March 24-25, 2008
Aim: How can one illustrate the process of a chemical reaction?
I. Potential Energy Diagrams = relationship between activation energy and heat of reaction for a chemical reaction
1. Activated Complex = temporary composite of atoms or molecules that will eventually form the products of the rxn.
= highest point on diagram is potential energy of this composite
= energy needed to form this composite = activation energy
2. Activation Energy - minimum amount of energy needed to form the Activated Complex and start the chemical reaction
3. Catalyst = atom or molecule that speeds up a chemical reaction when added to the rxn.
= lowers peak of Activated Complex b/c takes less energy to make
= Heat of reaction and potential energy of products don’t change.
4. Heat of Reaction = The heat released or absorbed in a chemical reaction. It is the difference between the Potential Energy of the Products and the Potential Energy of the Reactants
Heat of Reaction = PE Products - PE Reactants
A. Exothermic Rxn = Reactants --> Products + Heat Energy
1. self-sustaining i.e. burning match
2. spontaneous
3. Heat of Reaction is negative
4. Usually a reaction in which bonds are formed.
5. Feel warm to the touch as heat is released into the system.
B. Endothermic Rxn = Heat Energy + Reactants --> Products
1. non-spontaneous
2. Heat of Reaction is positive
3. Usually a reaction in which bonds are broken.
4. Feel cold to the touch as heat is removed from the system and absorbed into compounds in order to break bonds
II. Factors that Affect Reaction Rates
1. # of collisions = The more collisions b/t reacting species – atoms, molecules, ions, or other particles, the faster the reaction will occur.
- energy = reactants must have enough energy to make activated complex
- Therefore, more effective collisions, faster rate.
- No collisions, no reaction.
2. Nature of Reactants = kinds of bonds that need to be broken and/or made
- I.e. Ionic bonds = react faster in water and at room temp. b/c least amt. of bond rearrangement.
- I.e. Breaking bonds and making new ones (i.e. oxygen + hydrogen = water) react at slow rate at room temp.
3. Concentration - High conc. = rate of rxn. increases b/c possibility for collisions inc.
4. Temp inc. = rate of rxn. increases b/c particles move faster
- # collisions per unit time increases
- higher percentage of effective collisions
5. Surface Area - more surface area, faster rxn
6. Catalyst - atom, ion, compound that speeds up a chemical rxn without affecting the products
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Lesson #19 Le Chatelier’s Principle March 26-27, 2008
Aim: What is Le Chatelier’s Principle?
Video: Chemical Equilibrium - Reaction Tendencies
I. Le Chatelier’s Principle = When a system at equilibrium is disturbed by a stress, the position of the equilibrium will shift to relieve that stress. (i.e. Thermostat)
Dynamic Equilibrium for a Chemical Reaction:
The rate at which products are made is the same as the rate at which reactants are made.
Note: For LeChatelier's Principle to work, the chemical system MUST be at equilibrium first.
A. Changing Concentration – if rxn is at equilibrium and you decrease the concentration of a substance, the reaction will shift towards the removed substance.
- if rxn is at equilibrium and you increase the concentration of a substance by adding some of that substance, then the reaction will shift away from the increased substance.
B. Adding An Inert Gas - add an inert gas to increase the total pressure of the system, nothing will change. The equilibrium pressure will NOT change.
Inert = Inactive = Stable
C. Change In Temperature - if a reaction at equilibrium has a temperature change, due to an increase or decrease in the temperature of the system, then the equilibrium position will change.
- if the temperature increases then the reaction will proceed toward the side that consumes energy.
For example, in the equation:
N2 + 3 H2 <=====> 2 NH3 + 92 kJ
the reaction will proceed to the left (towards the side consuming energy.)
[The Reaction Will Move in an Endothermic Direction.]
- Likewise, if the temperature is decreased, then the reaction will proceed toward the side that does not consume energy (Exactly opposite that of the above example.) (Note: Please note that the equilibium position will change but the equilibrium constant will NOT change with a change in temperature.)
D. Changing The Volume Of A Container - if a reaction at equilibrium has a pressure change (due to an increase or decrease in container size), then the equilibrium position will change.
- if the container size increases then the reaction will proceed towards the side with the most number of gaseous moles.
- if the container size decreases then the reaction will proceed towards the side with the least number of gaseous moles.
E. Pressure Changes and their Effect on Equilibrium -
A pressure change potentially affecting the position of equilibrium can be accomplished two different ways.
1. The first way is to make the pressure in the reaction vessel higher by introducing an inert gas like argon. The pressure does go up, but this has no effect on the position of the equilibrium. Keep that in mind. NO EFFECT. End of discussion.
2. The second way to change pressure is to change the volume of the container up or down. Changing the volume up will reduce the pressure and reducing the volume will send the pressure up. The position of the equilibrium will be changed, except in one particular circumstance, demonstrated in example #6.
F. ENTROPY - The DISORDER of a system.
1. Gases have high entropy because the arrangement of particles is random.
2. Solids have low entropy because the arrangement of particles is ordered.
II. EXAMPLES
Example #1 - which way will the equilibrium shift if more H2 is added to this reaction at equilibrium:
N2 + 3 H2 <==> 2 NH3
Answer - the H2 amount goes up (by adding it), therefore according to LeChatelier's Principle, the reaction will try and use up the added H2. It does so by shifting the position of equilibrium to the right. This makes more NH3 by using up N2 and H2
Example #2 - using the same reaction, which way will the equilibrium shift if some NH3 is removed from the reaction when it is at equilibrium.
Answer - according to LeChatelier's Principle, the chemical system will attempt to replace the lost NH3. The stress was to remove NH3, so the opposite is to replace it. The equilibrium position will shift to the right in order to replace some of the lost NH3.
Example #3 - which way will the equilibrium shift if the system temperature goes up (heat is added):
2 SO2 + O2 <==> 2 SO3 + heat
Answer - even though heat is not a thing, for the purposes of LeChatelier's Principle, you can treat it as if it has physical existence. Since heat is added, the reaction will shift to try and use up some of the added heat. In order to do this, the reaction must shift to the left.
Example #4 - using the same reaction, which way will the equilibrium shift if heat is removed (that is, the temperature goes down).
Answer - the reaction will attempt to do the opposite of what the stress was. Since the stress was to remove heat, the reaction will shift to the right to generate more heat (replacing only a part of what was lost).
Example #5 - the container holding the following reaction (already at equilibrium) has its volume suddenly reduced by half. Which way will the equilibrium shift to compensate?
PCl3 + Cl2 <===> PCl5
Answer - since the volume went down, this means the pressure went up. The reaction will try to lessen the pressure by shifting to the side with the lesser number of molecules. This means a shift to the right because for every PCl5 molecule made, two molecules are used up. The lesser the number of total molecules in the container, the lesser the pressure.
Example #6 - the container holding the following reaction (already at equilibrium) has its volume suddenly increased. Which way will the equilibrium shift to compensate?
H2 + Cl2 <===> 2 HCl
Answer - neither side is favored over the other since both sides have the same number of total molecules (two). No matter which way the reaction shift, the total number of molecules would remain unchanged. In cases like this, where there is an equal number of molecules on each side, the equilibrium would remain unchanged by the change in pressure (in either direction).
Example #7 - the system below is already at equilibrium when a catalyst is added to the system. What happens to the position of the equilibrium? Does it shift right, left, or no change?
PCl3 + Cl2 <===> PCl5
Answer - there will be no change in the equilibrium. BOTH (with emphasis on both) the forward and the reverse reactions are speeded up. A catalyst just gets you to equilibrium faster, it doesn't affect the final position of equilibrium like changing the concentration would.
What is the effect of each of the following changes?
a) decrease in temp
b) increase in pressure
c) increase in [NO]
d) decrease in [H2O]
Answers: a) shift left
b) shift right
c) shift right
d) shift left
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Lesson #20 Oxidation Numbers March 31, 2008
Aim: How do I find oxidation numbers on the Periodic Table?
I. Oxidation Numbers = Oxidation States
1. show the electrical charge of an atom if it has lost or gained electrons
2. Cation - positively charged ion whose atom has lost electrons, lost negative charge
3. Anion - negatively charged ion whose atom has gained electrons, gained negative charge
4. Free elements - atoms whose oxidation number is zero
a. HOFBrINCl
b. Noble Gases
c. solid metals, etc.
d. any uncombined element
5. Find the oxidation numbers for each element in the following compounds:
a. NaClO4
b. KClO
c. PbSO4
d. Cu2CrO4
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Lesson #21 Redox April 2, 2008
Aim: What are Redox Reactions?
I. Redox - contraction for "REDuction" and "OXidation"
A. Occur everyday in:
1. bleaching
2. photography
3. rusting
4. batteries
B. Reduction - when an ion, atom, or group of atoms gains electrons, thereby reducing the charge and oxidation number
(Reduction Half Rxn) i.e. Cu 2+ + 2 e- --> Cu 0
C. Oxidation - when an ion, atom, or group of atoms loses electrons, thereby increasing the charge and oxidation number
(Oxidation Half Rxn) i.e. Cu 0 --> Cu 2+ + 2 e-
II. "LEO" the lion says, "GER"
L = lost
E = electrons
O = oxidation
G = gained
E = electrons
R = reduction
or you can use another pneumonic device to help you remember the difference between reduction and oxidation:
O = oxidation
I = is
L = lost
R = reduction
I = is
G = gained
or make up your own!
III. Redox Equations
1. MnO2 + 4HCl --> MnCl2 + Cl2 + 2H2O
Reduced element = Mn from +4 charge to +2 charge
Oxidized element = Cl from -1 charge to 0 charge
2. 16HCl + 2KMnO4 --> 8H2O + 2KCl + 2MnCl2 + 5Cl2
3. Mg(NO3)2 + Ni --> Mg + Ni(NO3)2
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Lesson #22 Electrochemistry April 3, 2008
Aim: What is a Voltaic Cell (aka Galvanic Cell)?
I. Voltaic/Galvanic Cell - uses a spontaneous redox rxn to produce an electric charge by the transfer of electrons.
i.e. a battery
1. Spontaneous Redox Rxn - favors products in reaction, releases heat, Delta H = -, Exothermic Reaction
2. contains one oxidation half reaction cell (anode) connected to one reduction half reaction (cathode) by a salt bridge (a tube containing a strong electrolyte such as potassium sulfate).
3. may also be connected by a porous barrier
4. Salt Bridge/Porous Barrier permit the migration of ions between solutions
5. FAT CAT - electrons flow From Anode To CAThode
6. Red Cat = reduction occurs at the cathode
7. Table J is used to identify the anode and cathode of an electrochemical cell. The element highest on the chart donates electrons (anode) to the element lowest on the chart (cathode).
8. An Ox follows the FAT CAT to the Red Cat
II. Picture of a Voltaic Cell: http://en.wikipedia.org/wiki/Galvanic_cell
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Lesson #23 Voltaic vs. Electrolytic Cells April 4, 2008
Aim: What is an Electrolytic Cell?
I. Electrolytic Cell - nonspontaneous redox reaction which needs electricity in order to produce a chemical reaction
A. Endothermic
B. Does NOT favor products
II. Uses for Electrolytic Cells
A. Purification of Metals - obtaining group 1 & 2 metals by reducing their fused salts
B. Decomposition of Compounds
i.e. 2H2O(l) + electricity --> 2H2(g) + O2(g)
C. Electroplating - the cathode is always the object that gets electroplated by the anode metal
III. What is the difference between a Voltaic Cell and an Electrolytic Cell?
A. Voltaic Cell - aka Electrochemical Cell, rxn is spontaneous, chemical rxn generates electricity, Exothermic, Cathode = +, Anode = -
B. Electrolytic Cell - rxn is non-spontaneous, electricity generates chemical rxn, Endothermic, Cathode = -, Anode = +
C. Picture of an Electrolytic Cell http://www.usoe.k12.ut.us/curr/science/core/assess/chem/cop.jpeg
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Lesson #24 Intro to Organic Chemistry April 28, 2008
Aim: How does Organic Chemistry Differ from Inorganic Chemistry?
I. Organic Chemistry is: the study of carbon and its compounds
1. Carbon has 4 valence electrons and so has 4 bonding electrons and so can have a total of 4 covalent bonds
2. Carbon can have single, double, or triple bonds
A. Bonding
1. Inorganic = Ionic and Polar Covalent
2. Organic = Covalent with little to no polarity
B. Solubility
1. Inorganic = soluble in polar solvents, like water
2. Organic = soluble in non-polar solvents, like oil
C. Electrolyte Behavior
1. Inorganic = can conduct a current in solution due to ionization
2. Organic = do not conduct a current in solution (except Organic Acids)
D. Melting and Boiling Points (related to Intermolecular Forces)
1. Inorganic = melt and boil at high temps
2. Organic = melt and boil at lower temps
E. Rate of Reaction (related to what occurs in the Activated Complex)
1. Inorganic = react rapidly due to simpler and more “flexible” bonds (more bond rearrangement)
2. Organic = do not react rapidly due to strong covalent bonds with little to no "flexibility" (bond rearrangement)
II. Alkanes – hydrocarbons that contain only single bonds between carbon atoms [Tables P & Q]
1. Formula: CnH2n+2 , where n = # carbon atoms
2. All Alkanes are Saturated Hyrdrocarbons because they contain the maximum number of hydrogen atoms possible to saturate the compound.
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Lesson #25 Isomers and IUPAC April 29, 2008
Aim: What are Isomers?
I. Isomers are compounds with the same molecular formula, but different structural formulas
1. butane vs. isobutane (C4H10)
2. pentane has 3 isomers
n-pentane (normal pentane) isopentane neopentane
3. hexane has 5 “constitutional isomers”, which are 2 or more compounds with different structures, but the same molecular formula
4. heptane has 9 isomers
5. decane has 75 isomers
II. IUPAC Naming System
1. Determine the “parent structure” by finding the longest unbroken chain of carbons
2. Modify name by including functional groups attached to parent structure
3. If no groups attached, it is “normal”, label “n-” in front of name
4. Number the parent chain so that the functional groups, double/triple bonds, are closest to Carbon #1.
5. Order the functional groups alphabetically on the parent chain
III. Hydrocarbon Functional Groups
Methyl = CH3-
Ethyl = CH3CH2- or C2H5-
Propyl = CH3CH2CH2- or C3H7-
IV. Naming Multiple Functional Groups
2 = di
3 = tri
4 = tert
5 = pent
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Lesson #26 Alkene Series May 7, 2008
Aim: What are Alkenes?
I. Alkenes – hydrocarbons that contain at least one carbon-carbon double bond
1. Formula: CnH2n
2. Alkenes are considered unsaturated hyrdrocarbon compounds because they do not contain the maximum number of hydrogen atoms possible to saturate the compound.
i.e. Propene C3H6
i.e. C4H8 1-butene and 2-butene
II. Stereoisomerism – isomers whose atoms have different arrangements in space
i.e. 2-butene has 2 stereoisomers
cis-2-butene trans-2-butene
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Lesson #27 Alkynes May 7, 2008
Aim: What are Alkynes?
I. Alkynes – hydrocarbons that contain at least one carbon-carbon triple bond
1. Formula – CnH2n-2
2. Alkynes are considered unsaturated hyrdrocarbon compounds because they do not contain the maximum number of hydrogen atoms possible to saturate the compound.
i.e. H-C _= C-H ethyne or acetylene
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Lesson #28 Cyclo-Hydrocarbons & Benzene Rings May 8, 2008
Aim: What are Cyclo-Hydrocarbons, Benzene Rings, and Halocarbons?
I. Cyclo-Hydrocarbons - hydrocarbons arranged in a ring instead of in a chain
II. Benzene Series – aka aromatic hydrocarbons has the formula CnH2n-6
1. benzene C6H6
2. toluene (methylbenzene) C7H8 – methyl group replaces a hydrogen
III. Halocarbons - hydrocarbons that have at least one halogen attached to the parent structure (see Table R)
IV. Nomenclature
1. If the ring has more carbons than the carbon chains branching off of it, then it is the parent structure and the chains are the functional groups.
2. If the carbon chain has more carbons than the ring, then it is the parent structure and the ring is the functional group.
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Lesson #29 Alcohols May 9, 2008
Aim: What are Alcohols?
I. Alcohols – hydrocarbons that contain at least one hydroxyl functional group OH-
A. Classified by
1. # hydroxyl groups
2. way hydroxyl groups are attached to parent chain
3. ends in suffix “-ol”
B. Alcohol Series
1. methanol – solvent, aka methyl alcohol
2. ethanol – antiseptic & alcoholic beverages, aka ethyl alcohol
3. Condensed Structural Formulas
- methanol CH3OH - ethanol CH3CH2OH
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Lesson #30 Aldehydes, Ketones, Ethers May 9, 2008
Aim: What is the Carbonyl Group?
I. Carbonyl Group – carbon atom double bonded to an oxygen atom
A. Aldehydes – carbon of the carbonyl group is bonded to at least one hydrogen
- always found on the first carbon of the parent chain
-suffix “-al”
B. Ketones – carbonyl with no hydrogen atoms
-suffix “-one”
ketone propanone (aka acetone-solvent in nail polish remover)
C. Ethers – R1-O-R2 , can be made by dehydrating primary alcohols
2CH3CH2OH --> C2H5-O-C2H5 + H2O
Diethyl ether – once a popular surgical anesthetic
- the smaller chain comes first in the name, followed by the longer chain and then the suffix, "ether"
i.e. CH3CH2-O-CH3 methyl ethyl ether
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Lesson #31 Carboxylic Acids and Amines May 13, 2008
Aim: What are Carboxylic Acids and Amines?
I. Carboxyl Group : Organic Acids
1. carboxyl is a contraction of carbonyl and hydroxyl
2. abbreviated as COOH
3. acidic properties
4. suffix “-oic acid”
5. methanoic acid (aka formic acid = ant venom) HCOOH
6. ethanoic acid (aka acetic acid = 5-6% solution is vinegar) CH3COOH
II. Organic Amines
1. Amino Group -NH2 compound
2. amine = hydrocarbon with one or more hydrogen atoms replaced by an amino group
3. aminomethane or methanamine
4. amino acids = organic acid with one or more amino groups
a. 2-aminoethanoic acid = glycine
b. 2-aminopropanoic acid = alanine
- carbon is “chiral” b/c bonded to 4 differ groups
5. Amides
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Lesson #32 Substitution & Addition Rxns May 14, 2008
Aim: What are substitution and addition rxns?
Do Now: Name and Draw the following organic compounds
1. CH3CH2CH2CHNH2CH2COOH
2. CH3CH2CH2CHNH2CH3
3. 2,2-dichloro-3-methylheptanoic acid
I. Organic rxns are slower than inorganic rxns
1. non-polar covalent bonding in organic compounds allows for atoms to be held more tightly due to the even sharing of electrons (strong bonds, weak intermolecular forces)
2. polar/ionic bonding in inorganic compounds allows for atoms to be held more loosely due to the uneven sharing of electrons (weak bonds, strong intermolecular forces)
II. Substitution
1. one type of atom is replaced by another
i.e. halogens can substitute the hydrogens of alkanes
2. alkanes only
2. after addition triple bond becomes a double bond
3. faster rxn than substitution rxns
4. characteristic of unsaturated compounds
5. if add hydrogen to unsaturated compound = hydrogenation, process used to solidify liquid vegetable oils
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Lesson #33 Fermentation May 15, 2008
Aim: What is fermentaton?
I. Fermentation = biological process in which oxidation occurs in the absence of oxygen
II. Yeast enzymes are living organisms that act as catalysts in fermentation
III. Fermented glucose --> ethanol and carbon dioxide
yeast enzymes
zymase
C6H12O6 --------------> 2 C2H5OH + 2 CO2
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Lesson #34 Oxidation, Polymers, & Polymerization May 15, 2008
Aim: What are Oxidation and Polymerization Reactions?
I. Oxidation (aka Combustion/Respiration)= saturated hydrocarbons react w/ O2 at high temps to produce
CO2 & H2O
1. CH4 + 2O2 --> CO2 + 2H2O
2. makes alkanes suitable as fuels
3. in limited supply of O2, carbon monoxide or carbon soot are produced
a. 2CH4 + 3O2 --> 2CO + 4H2O
b. CH4 + O2 --> C + 2H2O
II. Polymerization
1. Polymer – large molecule of many repeating units called monomers
2. proteins (amino acids), cellulose/starch (glucose), polyethylene (ethene), nylon, polyester
3. Condensation Polymerization – monomers joined by dehydration rxn
a. monomer must have at least 2 functional groups
b. silicones, nylons, polyesters
c. HOCH2CH2OH + HOCH2CH2OH --> HOCH3CH2-O-CH2CH3OH + H2O
4. Addition Polymerization – monomers joined by addition rxn
a. double and triple bonds produce single and double bonds
b. “~” indicates repeating monomer units
c. polyethylene
d. n(CH2=CH2) --> (~CH2CH2~)n
alkene --> alkane
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Lesson #35 Esterification & Saponification Reactions May 16, 2008
Aim: What are Esterification and Saponification Reactions?
I. Esterification - the formation of an ester by combining an alcohol with an organic acid through a dehydration process
RCOOR' = general formula for an ester, the carbon is double bonded to one oxygen and single bonded to the other oxygen
The "ic" in ethanoic acid turns into "ate" in the ester. The "ate" part of the name is the chain the includes the carbonyl carbon. The "yl" part of the name belongs to the chain that is attached to the oxygen with a single bond.
Sodium stearate is the soap. It has a non-polar, hydrophobic end ("water-fearing"), which is attracted to organic substances (stains, oil, etc.) and a polar, hydrophilic end ("water-loving") that is attracted to water and washed down the drain. This dual property is what makes up detergents.
Aim: What is the Kinetic Molecular Theory of Liquids?
Do Now: Describe what happens at the particle level when a liquid condenses.
I. Liquids
1. occurs when temp of gases decreases and pressure of gas increases (this squishes the particles more closely together)
2. Particles DO NOT exhibit ideal behavior
a. have measurable volume
b. exert forces on each other
3. takes shape of its container
II. Factors Affecting Changes of State of Liquids
1. The nature of the liquid (i.e. ethanol evaporates faster than water)
2. Liquid temp. (hot water evaporates faster than cold water)
3. The surface area of liquid (1.0L water in a puddle evaporates faster than 1.0L water in a bottle)
III. Vapor Pressure – pressure evaporated liquid exerts on the surface of the liquid (Table H)
1. depends on nature of liquid and its temp
2. rises slowly at low temps
3. when container is open to atmosphere, vapor pressure increases until equals pressure of atmosphere above liquid = liquid BOILS (boiling point)
IV. Practice with Table H
1. At what temperature will propanone boil if the pressure is 150 kPa?
2. Which liquid will boil first at 1 atm according to Table H?
3. What will the pressure be if ethanol boils at 55 degrees Celsius?
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Lesson #__ Colligative Properties
Aim: What is the effect of adding a solute to a solvent?
I. Colligative Properties = factors that depend on the number of solute particles present and the nature of the solvent. (These properties do not depend on the identity or type of solute used.)
1. Vapor Pressure Lowering
2. Boiling Point Elevation
3. Freezing Point Depression
A. Vapor Pressure Lowering
1. vapor pressure = pressure exerted by a vapor (gas) that is in dynamic equilibrium (moving balance) with its liquid in a closed system.
OR It is the pressure produced when vapor particles above a liquid, in a sealed container, collide with the container walls and a dynamic equilibrium exists b/t the vapor and the liquid.
2. Solution with nonvolatile (not easily vaporized) solute has lower vapor pressure than pure solvent
i.e. NaCl in water. Na+ and Cl- ions make it more difficult for water molecules to escape into their vapor form, therefore decreasing vapor pressure
B. Boiling Point Elevation
1. boiling point = temp at which vapor pressure of liquid phase equals the atmospheric pressure
2. B/c adding solute to solvent lowers vapor pressure, more kinetic energy must be added to get solvent molecules into gas phase
3. B/c more kinetic energy must be added, the solvent will boil at a higher temperature
C. Freezing Point Depression
1. freezing point = when particles of a solid take on an orderly pattern
2. Adding solute to solvent will lower freezing point temperature b/c solute gets in-between the particles and hinders the formation of this “orderly pattern”
3. Therefore, more kinetic energy must be removed from the solution for it to solidify.
4. i.e. When add salt to icy sidewalks to make ice melt and keep water on sidewalks from freezing quickly again (by lowering the freezing point temp.)
D. The more solute added to solution, the more exaggerated the colligative properties will be.
E. Electrolytes = solutes that dissociate (dissolve) into ions and conduct electricity in a solution
1. Ionic solutes are electrolytes that can dissolve in polar solvents
2. Polar covalent solutes are non-electrolytes that can dissolve in polar solvents
3. Non-polar covalent solutes are non-electrolytes that can dissolve in non-polar solvents
4. "Like dissolves Like"
II. Examples of Colligative Properties (Old Regents Curriculum)
1. Water boils at 100°C and freezes at 0°C. If you add 1 mole of solute (i.e. glucose) to 1 kg of water (solvent), the boiling point raises to 100.52°C
and the freezing point decreases to –1.86°C. (Note that 1 mole solute (molecular) ÷ 1kg solvent = 1 Molal solution.)
A. Primary, Secondary, Tertiary Alcohols
1. If the alcohol group is attached to the first carbon, then it is a primary alcohol.
2. If the alcohol group is attached to a middle carbon with 2 "R" groups, then it is a secondary alcohol.
3. If the alcohol group is attached to a middle carbon with 3 "R" groups, then it is a tertiary alcohol.
B. Diols & Triols
1. Diols - contain 2 hydroxyl groups on the parent structure
i.e. 1,2 - ethanediol
2. Triols - contain 3 hydroxyl groups on the parent structure
This is a single displacement reaction. Iron (II) will displace Lead (II) because they are both cations and cannot bond with each other.
Fe 2+ + Pb(NO3)2 ---> Pb 2+ + Fe(NO3)2
IV. Isotopes and Average Atomic Mass
Determine the average atomic mass of the following mixtures of isotopes:
1. 80% I-127, 17% I-126, 3% I-128
2. 50% Au-197, 50% Au-198
3. 15% Fe-55, 85% Fe-56
4. 99% H-1, 0.8% H-2, 0.2% H-3
5. 95% N-14, 3% N-15, 2% N-16
6. 98% C-12, 2% C-14
V. Mixed Mole Problems
1. How many grams are there in 1.5 x 10e25 molecules of CO2?
2. What volume would the CO2 in Problem #1 occupy at STP?
3. A sample of NH3 gas occupies 75.0 liters at STP. How many molecules is this?
4. What is the mass of the sample of NH3 in Problem #3?
5. How many atoms are there in 1.3 x 10e22 molecules of NO2?
6. A 5.0g sample of O2 is in a container at STP. What volume is the container?
7. How many molecules of O2 are in the container in Problem #6? How many atoms of oxygen is this?
_____________________________________________________________
Ms. Pace’s Chemistry Extra Credit Assignment
- To Replace Second Lowest Exam Grade. (Lowest Exam Grade gets dropped automatically.)
- Due ________________.
For a grade of 95%:
1. Build a molecular model of an over-the-counter or prescription drug. (Do not use any type of food product!)
2. Write a 1 page paper describing the chemical mechanism of the drug (how it works). Must also include:
a. Chemical formula
b. 1-2 symptoms it alleviates
c. 1-2 risks of taking it
d. area of body it directly affects
3. Must be a work of fiction, story, allegory, song, play, etc.
4. Must be typed, 12 point font (Times New Roman or Arial), 1.5 spacing, 1 inch margins
For a grade of 100%:
Must include #1-4 above AND must be presented/performed in front of the class as a story, song, skit, interpretive narrated dance, etc.
* If you want to work with a partner(s), then for every additional person, there will be an additional page written to the story.
Chemistry Regents Prep Old Regents Exams(interactive), Additional Chemistry Resource Sites...
Links For All Science Subjects This site was created by the elegant Dr. Judy Brickell. Though she is no longer with us, her legacy lives on through the painstaking hours she spent organizing this site to help all students in all areas of the sciences.
Electrolytic Cell A diagram of an electrolytic cell used for electroplating.
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