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Ms. Antonietta Pace
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Regents Chemistry 1 - Fall 2007
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FIORELLO H. LA GUARDIA HS
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Dear Students,
Welcome to the Wonderful World of Chemistry!
One of our goals this year should be to learn to appreciate chemistry and its affect on us in the "real world".
On this web page you'll find daily notes, homework assignments, and links to other valuable resources on the net. The list of homework assignments can be found in the Flashcards section.
Here's to a great school year!
Best,
Ms. Pace
P.S. My schedule for the fall term: TBA
My desk is located in Rm 523.
MAKE-UP LABS: TBA
Our Textbook is Prentice Hall Chemistry.
------------------------ATTENTION----------------------
Regents Chemistry 1 - Course Expectations
Units Covered in Term 1:
1. Classification of Matter
2. Energy
3. Atoms & Atomic Structure
4. Periodic Table & Trends
5. Bonding – Intramolecular and Intermolecular
6. Nuclear Reactions
7. Gas Laws
8. Kinetic Molecular Theory
LaGuardia Chemistry Grading Policy
1. Exams 50 %
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 15 %
a. Late HW will have a penalty of 10 points for every day they are not turned in.
3. Labs 20%
a. All labs must be done and handed in on time
b. If a lab is missed, students have one week to make up the lab
4. Class Participation 15%
a. Coming to class on time
b. Handing in work on or before the due date
c. Paying attention in class
d. Asking and/or answering teacher/student questions related to the course
e. Working problems out on the board
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.
*****************
To acknowledge that you have read the Course Requirements, please send an email to:
apace2@schools.nyc.gov
Subject Line: Name, Chemistry Period ____
Body of Email: Name, Class Period, I agree with the Course Requirements as stated on www.schoolnotes.com/10023/mspace.html.
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CLASSROOM NOTES
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Lesson #1 September 5-6, 2007
Aim: What is Chemistry?
I. Chemistry focuses on the structure and properties of matter and the energy that accompanies changes in matter.
II. Matter = anything that has mass and takes up space (volume)
A. Substances = Any form of matter
i. Elements = cannot be broken down into simpler substances by chemical means
1. atom = smallest particle of an element that can enter a chemical reaction
a. electrons = negatively charged particles that circle the atom's nucleus and weighs 1/1836 amu, less than 1 amu (atomic mass unit)
b. protons = postively charged particles that make up the atom's nucleus and weighs 1 amu
c. neutrons = neutrally charged particles that make up the atom's nucleus and weighs 1 amu
2. ion = charged atom of an element that has gained or lost electrons
a. Cation - postively charged atom that has lost electrons i.e. K+
b. Anion - Negatively charged atom that has gained electrons i.e. Cl-
c. Number of electrons is different from number of protons
d. Number of protons for an element does not change
3. isotope = radioactive atom of an element that has gained or lost neutrons i.e. C-12 vs. C-14
4. new element = if protons are gained or lost
ii. Compounds = consist of two or more elements bonded together and can be broken down into its elements by chemical means
i.e. water, H2O
B. Mixtures = 2 or more substances that are uniform & has a fixed composition
1. Homogeneous Mixture = aka Solutions - substances are uniformly dispersed throughout the mixture
Examples = sugar water, etc.
2.Heterogeneous Mixture = aka Suspensions - substances are not uniformly dispersed throughout the mixture
Examples = salad, ice cream, syrup, soda,whipped cream (not uniformly dispersed)
HW #1: Send Email Acknowledging Course Expectations, Due Friday, September 7th
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Lesson #2 Physical & Chemical Properties/Changes
September 7, 2007
Aim: What is the difference between physical and chemical properties/changes of matter?
I. Physical Properties
a. Length, Color, Temperature, Mass, Volume, Density, Malleability, Ductility, Conductivity, Melting & Boiling Points
b. Physical Changes – changes that do not affect the chemical character of a substance, only the physical appearance is changed and the same substance is still present
i. Cutting wood into smaller pieces
ii. Dissolving sugar in water
iii. Pouring liquid from one container to another
iv. Separating mixtures – any mixture can be separated by physical changes such as distillation and crystallization
v. PHASE CHANGES
II. Chemical Properties – when properties of matter depend upon the action of substances in the presence of other substances
a. Does Substance A burn?
b. Does Substance A help Substance B to burn?
c. Does Substance A react with Substance B?
d. Does Substance A react in water?
e. Does Substance A react with an acid or base?
f. Chemical Changes – produce new substances with new properties
(whenever a substance undergoes a change that produces one or more new substances with different properties from the original substance)
i. burning, digestion, fermentation, ionization, radioactive decay, decomposition, synthesis
III. Practice Problems –
Classify the following properties as chemical or physical.
a) Color -
b) Reactivity -
c) Flammability -
d) Odor -
e) Porosity -
f) Stability -
g) Ductility -
h) Solubility -
i) Expansion -
j) Melting Point -
k) Rusting -
l) Reacts with Air -
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Lesson #3 States of Matter September 10, 2007
Aim: What are the properties of the 3 states of matter?
1. solid = definite shape and definite volume
= consist of vibrating particles, with strong attractions to each other, arranged mostly in a rigid geometric pattern
2. liquid = indefinite shape and definite volume
= consist of particles that move in 3 dimensions, but still remain attracted to one another
3. gas = indefinite shape and indefinite volume
= consist of particles that move in 3 dimensions with little to no intermolecular attractions
These are studied in Chemistry because they exist at room temperature on this planet.
4. plasma = superheated gas, where atoms are stripped of electrons at temperatures over 5000 degrees Celsius
= most of the universe is made of plasma
= i.e. matter inside a neon tube
= nuclear fusion reactions only occur in plasmas
= YES! You can find this substance inside plasma television sets!
5. Bose-Einstein condensate = exist at very low temps near 0 Kelvin (Scientists have been able to reach temps of 200 billionths above 1K in the lab), atoms of an element collapse into each other to create a “superatom” that still retains the properties of the element. All the atoms are so close together that they end up occupying the same exact space.
= NO! This is Satyendra Nath Bose, NOT Amar G. Bose, founder of the sound system company Bose.
(To see a picture of a Bose-Einstein condensate visit:
http://math.nist.gov/mcsd/savg/papers/images/3D.00007.jpg)
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Lesson #4 Phase Changes September 12, 2007
Aim: What is a phase change?
I. Phase Change – when a physically distinct section of matter (that is set off from the surrounding matter by physical boundaries) changes into a different physically distinct section of matter.
- depends on the nature of the substance, the pressure, and the temp of the matter's environment
A. Gives the conditions of temp and pressure at which a substance exists as a solid, liquid, or gas
B. Melting = solid to liquid (heat is absorbed = endothermic)
C. Freezing = liquid to solid (heat is released = exothermic)
Melting and Freezing points will occur at the same temperature. It just depends if you are adding heat (Melting) or removing heat (Freezing).
D. Vaporization (Evaporation, Boiling) = liquid to a gas (heat is absorbed = endothermic)
E. Condensation = gas to a liquid (heat is released = exothermic)
Vaporization and Condensation points will occur at the same temperature. It just depends if you are adding heat (Vaporization) or removing heat (Condensation).
F. Vapor Pressure (Table H=Vapor Pressure Graph)
a. Vapor Pressure is the measure of the force exerted by a gas above a liquid.
b. Increasing temperature of a liquid, increases the vapor pressure because it sends more particle from the liquid state to the vapor state by increasing average kinetic energy of the particles in the liquid.
c. When vapor pressure of a liquid equals the external pressure on the liquid, then the liquid has reached its boiling point.
d. This graph shows the boiling points of 4 liquids at different atmospheric pressures.
i. The lower the external pressure, the lower the boiling point temperature, the higher the vapor pressure and the particles need less kinetic energy to escape from the liquid.
ii. The higher the external pressure, the higher the boiling point temperature, the lower the vapor pressure and the particles need more kinetic energy to escape from the liquid.
i.e. #1) What is the vapor pressure of water at 105 degrees Celsius? a)100kPa b)101.3kPa c)110kPa d)120kPa
#2) As the pressure on a liquid is changed from 110.kPa to 120.kPa, the temperature at which the liquid will boil a)decreases b)increases c)remains the same d)depends on the liquid
#3) In a closed system at 40 degrees Celsius, a liquid has a vapor pressure of 50 kPa. The liquid's normal boiling point could be a)10 C b)30 C c)40 C d)60 C
#4) If the pressure on the surface of water in the liquid state is 30kPa, the water will boil at a)0 C b)30 C c)70 C d)100 C
#5) A sample of ethanoic acid at 100 degrees Celsius has a vapor pressure of a)53kPa b)100kPa c)101.3kPa d)125kPa
#6) Which substance has the greatest intermolecular forces of attraction between molecules?
a)propanone b)ethanol c)water d)ethanoic acid
#7) The vapor pressure of ethanol at its normal boiling point would be a)80kPa b)101.3kPa c)273kPa d)373kPa
#8) What is the normal boiling point of propanone? (approximately 55 degrees Celsius)
#9) Which sample has the highest vapor pressure?
a)ethanol@80C b)water@20C c)propanone@65C d)ethanoic acid@50C
G. Sublimation = solid to a gas (heat is absorbed = endothermic) i.e. Dry Ice, Air Freshener
H. Deposition = gas to a solid (heat is released = exothermic) i.e. Frost, contents of fire extinguisher when released
Sublimation and Deposition points will occur at the same temperature. It just depends if you are adding heat (Sublimation) or removing heat (Deposition).
II. Heating & Cooling Curves - http://www.wwnorton.com/chemistry/tutorials/ch11.htm (click on "view tutorial" for Section 11.6 Heating Curves)
A. Show a mathematical, visual picture of phase changes in a substance.
B. Latent Heat – The quantity of heat absorbed or released by a substance undergoing a change of state, such as ice changing to water or water to steam, at constant temperature and pressure. It is called latent because it is not associated with a change in temperature.
- On the Heating & Cooling Curves, Latent Heat appears as a straight line. This is because some particles have high energy and others have low energy during a phase change and when you take the average kinetic energy of the substance you get a single temperature. Hence, the straight line. Even though particles are moving at different speeds.
C. During a phase change, POTENTIAL ENERGY is increasing as heat energy is added to the system. However, KINETIC ENERGY is constant as temperature remains the same.
D. During "warming", KINETIC ENERGY is increasing as temperature is increasing. POTENTIAL ENERGY is constant.
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Non-Regents Supplemental Lesson: Phase Diagrams
Aim: What are Phase Diagrams?
I. Phase Diagrams - illustrate the relationships between pressure and temperature and how they relate to changes in states of matter (Phase Diagram - wine1.sb.fsu.edu/.../ Forces/Phase/Forces06.htm)
What do you expect to happen at the Triple Point?
1. Triple Point – where solid, liquid, gas exist in dynamic equilibrium for a substance at a specific temperature and pressure
a. the only set of conditions at which all 3 phases can exist in equilibrium with one another
2. critical point – only the gas phase can exist past this point, b/c kinetic energy of gases is extremely high and will overpower any pressure increase.
3. Normal Melting Point - mp temperature for a substance at 1 atm (101.3 kPa) of pressure or sea level
4. Normal Boiling Point - bp temperature for a substance at 1 atm (101.3 kPa) of pressure or sea level
5. STP = Standard Temperature and Pressure (Table A in Reference Tables)
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Lesson #5 Kinetic Molecular Theory September 17, 2007
Aim: What is Kinetic Molecular Theory (KMT)?
Do Now:
1. Which substance has a definite shape and a definite volume at STP?
a) Cl2(g) b) H2O(l) c) AlCl3(s) d) NaCl(aq)
2. Which substance is most likely to take the shape of and occupy the total volume of its container?
a) H2(g) b) CCl4(l) c) S(s) d) MgBr2(aq)
I. KMT - the understanding of the movement (kinetics) of particles (atoms/molecules/compounds).
- explains the effects of temperature and pressure on matter.
A. Heat added to a system
1. Temp increases b/c average kinetic energy increases
2. Particles move faster and spread out
3. More Collisions between particles and other particles and particles and walls of container.
4. The collisions result in a force being exerted which is called, "PRESSURE"
5. Pressure increases, which may in turn increase volume
6. Intermolecular forces become weaker as particles move farther from each other.
EXAMPLE: The Concord Jet would increase 8 inches in diameter during flight as the metal shell would expand due to the ultraviolet radiation at those high altitudes!
B. Heat removed from a system
1. Temp decreases
2. Particles move slower and condense
3. Less collisions
4. Pressure decreases
5. Volume decreases
6. Intermolecular forces become stronger and particles begin to "stick" together and change states of matter(phase changes i.e. liquid to solid)
EXAMPLE: The Concord Jet would decrease 8 inches in diameter during descent. Pilots who would leave flight manuals between their seats and the plane's walls would have a difficult time retrieving them after landing!
C. Entropy - disorder of a system.
1. The more disorder, the higher the entropy.
2. The more entropy, the further apart the particles are from one another.
3. Gases - high entropy
Solids - low entropy
II. Practice Problems
1. As a substance changes from a liquid to a gas, the average distance between molecules:
a) decreases b) increases c) remains the same
2. As a substance changes from a gas to a solid, the strength of the Intermolecular forces of the substance:
a) decreases b) increases c) remains the same
3. A camper inflates a mattress during a warm spring day. In the morning, the mattress appeared to have deflated. The camper found no leaks in the mattress. In terms of Kinetic Molecular Theory, describe what happened to the mattress.
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Lesson #6 Heat Equations September 19, 2007
Aim: What is Energy and How do I use the Heat Equations?
I. Energy – the ability or capacity to do work
A. Potential – stored energy
B. Kinetic – energy of motion
C. Specific Heat - the amount of heat it takes to raise the temp of 1.0g of a substance 1.0 degree C.
Specific Heat for liquid Water = 4.18 J/(g x degree C)
[Ref Table B]
II. Heat Equations in Ref Table T:
A. q = mC (Tf - Ti)
q= heat in joules
m= mass in grams
C= specific heat of substance in J/(g x degree C)
(Tf - Ti)= change in temperature in degrees Celcius (final temp - initial temp)
**Use this equation with problems containing two different temperatures.
B. q = mHf
a. Hf = Heat of Fusion (for H2O = 334 J/g) [Table B]
b. solids to liquids
c. liquids to solids
d. use for melting & freezing problems
**Use this equation with melting and freezing problems.
C. q = mHv
a. Hv = Heat of Vaporization (for H2O = 2260 J/g) [Table B]
b. liquids to gases
c. gases to liquids
d. use for vaporization & condensation problems
**Use this equation for evaporation/boiling and condensation problems.
D. Hv is always larger than Hf b/c more energy is needed to convert a liquid into a gas in order to free gas molecules from intermolecular forces of attraction between particles
III. Heat Equations Practice Problems:
1. A 3000 gram mass of water in a calorimeter has its temperature raised 5.0 degrees Celsius. How much heat energy was transferred to the water in joules?
2. How many joules of heat are absorbed when 50 grams of water at 100 degrees Celsius are completely vaporized?
3. If 20 joules of heat are added to 2.0 grams of water at 15 degrees Celsius, then what temperature did the water increase to?
OLD REGENTS QUESTIONS:
1. What is the total number of joules of heat energy absorbed by 15 grams of water when it is heated from 30 degrees Celsius to 40 degrees Celsius?
2. How many joules of heat are absorbed when 70.0 grams of water is completely vaporized at its boiling point?
3. The heat of fusion of a compound is 30 joules per gram. What is the total number of joules opf heat that must be absorbed by a 15.0 gram sample to change the compound from a solid to a liquid at its melting point?
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MORE Practice Problems
______________________
Name: _____________________ Date: ______ Period: ____
Chemistry 1 – Energy: Heat Equations
1. A 10g sample of H2O(s) melts. How much energy was required to completely melt the ice?
2. A 50g sample of H2O(s) melts. How much energy was required to completely melt the ice?
3. A 25g sample of H2O(l) evaporates. How much energy was required for the water to completely evaporate?
4. A 43g sample of H2O(l) evaporates. How much energy was required for the water to completely evaporate?
5. A 15g sample of water is at 90 degrees Celsius. If 9000 J of energy were used to heat the sample, then what is the final temperature of the water?
6. A 15g sample of water is at 45 degrees Celsius. If 5000 J of energy were used to heat the sample, then what is the final temperature of the water?
7. If 45000 J of energy is required to completely melt ice, then how many grams of ice are present?
8. If 3800 J of energy is required to completely condense water, then how many grams of water are present?
9. How many grams of water require 23200 J of energy to increase its temperature from 25 degrees Celsius to 39 degrees Celsius?
10. How many grams of water require 15290 J of energy to increase its temperature from 47 degrees Celsius to 52 degrees Celsius?
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Lesson #7 Metric System September 20, 2007
Aim: How do I Convert Metric Units in Chemistry?
Motivation: The 1999 Mars Climate Orbiter (http://mars.jpl.nasa.gov/msp98/news/mco990930.html)
(http://mars.jpl.nasa.gov/msp98/news/mco991110.html)
I. Measurement and the Metric System
A. We use SI (Systeme International) units as an international standard of measuring in the sciences.
B. Seven Fundamental Quantities from which all measurement is based:
1. Length = Meter = 39 inches
We use various subdivisions of the meter: picometer, nanometer, millimeter, centimeter, decimeter
2. Mass = Kilogram = 2.2 pounds (on Earth), 1000 grams = measure the amount of matter an object contains (weight = force with which gravity attracts matter)
3. Temp. = Kelvin = measures the “hotness” of an object by averaging the average kinetic energy of the particles [273K = 0 degrees Celsius = 32 degrees Fahrenheit]
a) Conversion from Celsius to Fahrenheit:
°F = 9/5 °C + 32
b) Conversion of Celsius to Kelvin:
K = °C + 273
c) The Celsius Scale is based on water's freezing point & boiling point.
d) 0K is called, "Absolute Zero". This will theoretically occur when all kinetic energy turns to potential energy. The closest scientists have come to creating an environment at absolute zero has been 0.0000000005K!
4. Time = second (minute, hour, day, year)
5. Number of Particles = Mole (mol) = 6.02 x 10e23 = measures the number of particles (such as atoms or molecules) in a sample of matter (This measurement allows us to predict the behavior of matter.)
6. Electric Current = Ampere = measures the flow of electric charge
7. Luminous Intensity = Candela (cd) = measures the brightness of light
(Numbers 6 & 7 are commonly used in Physics.)
II. The Metric System - based on units of 10
One way to memorize metric units more easily is to use this pneumonic device:
King Henry Died By Drinking Chocolate Milk (where the "B" in "By" stands for Basic unit.)
K= Kilo
H= Hecto
D= Dekka
B= basic unit: meter / gram/ liter
D= Deci
C= Centi
M= Milli
or
Kids Hate Doing Math During Certain Months (where the 1st "M" in "Math" stands for meter.)
III. Metric Prefixes from Chemistry Reference Table C & p. 4 Barron’s Chemistry Review Book
Factor Prefix Symbol
10e-12 Pico p
10e-9 Nano n
10e-6 Micro ì
0.001 Milli m
0.01 Centi c
0.1 Deci d
10 Decka da
100 Hecto h
1000 Kilo k
10e6 Mega M
10e9 Giga G
(Letter "e", after the power of 10, stands for "exponent")
100mm = 1cm
100cm = 1m
1000mm = 1m
1000m = 1km
IV. Practice Problems
1. 50 mm = ________cm
2. 100 mm = ________cm
3. 500 mm = ________m
4. 50 cm = ________m
5. 3.0 cm = ________m
6. 4000 mL = ________L
7. 6.0 L = ________mL
8. 2.0 kg = ________mg
9. 2532 g = ________kg
10. 3.0 mL = ________L
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Lesson #8 Dimension Analysis September 21, 2007
Aim: What is Dimension Analysis?
Dimension Analysis - the conversion of values from one unit of measurement to another
You must use a ratio or conversion factor that is equal to the number 1. You can multiply anything by the number 1 without changing its value.
1. How many seconds are in one year? (Try doing this without referring to that song in the musical "RENT".)
60 sec/1 min x 60 min/ 1 hr x 24 hr/ 1 day x 365 d/1 yr = 31,536,000 s/yr
3. How many feet are in 15m?
15m/1 x 39in/1m x 1ft/12in = 48.75ft
3. If 1.41 US dollars = 1 Euro, then how many euros make up $10?
4. If 5 bracks = 2 snorks, and 7 snorks = 3 toggles, then how many toggles are in 1 brack?
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Lesson #9 Significant Figures September 24 & 25, 2007
Aim: What are Significant Figures?
I. Significant Figures - indicate the exactness of a measurement. They are the digits in the value of a measurement indicating the quantity to an accuracy justified by the measuring device and technique used to make the measurement.
II. Rules for a Number in a Value to Be Considered Significant
1. Digits other than zero are always significant
96 2 sig figs
61.4 3 sig figs
2. One or more final zeros used after the decimal point are significant
4.7200 5 sig figs
3. Zeros between two other significant digits are always significant
5.029 4 sig figs
4. Zeros used solely for spacing the decimal point are not significant. The zeros are place holders only.
7000 1 sig fig
0.00783 3 sig figs
5. How many significant figures are there in each of the following?
a. 903.2
b. 90.3
c. 900.0
d. 0.0090
e. 0.0900
f. 99
g. 0.0088
h. 0.049
i. 0.02
j. 40
k. 4.0
l. 40.
III. Rules for Adding and Subtracting Values
The answer many contain only as many decimal places as the least accurate value used in finding the answer.
i.e. 677 + 39.2 + 6.23 = 722.43 (Answer NOT in Sig Figs)
722 = Answer to the correct number of sig figs.
IV. Rules for Multiplication and Division of Values
The answer may contain only as many significant digits as the least accurate value used to arrive at the answer.
i.e. (1.1330) x (5.12600000) = 5.807758000000
1.1330 has 5 sig figs, 5.126000000 has 9 sig figs
The answer must include only 5 sig figs: 5.8078
V. Rounding Numbers
1. Look to the number after the last significant digit.
2. If it is 4 or lower, leave the value as is.
3. If it is 5 or higher, round the value up.
i.e. Round the atomic mass of Chlorine to the nearest whole number. 35.453 amu = 35 amu
VI. Practice Problems
1. If the mass of a substance is 343.2g, and its volume is 22.9 cm3, then what is the density of the substance expressed to the proper number of significant figures?
D=M/V = 342.2g/22.9cm3 = 14.98699g/cm3 = 15.0g/cm3
2. Express the answer to the following problem to the correct number of significant digits.
459.37 - 2.6 - 25.21 - 14.0 = 417.56 = 417.6
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Supplemental Lesson: Scientific Notation
Aim: How do I write in Scientific Notation?
I. Scientific Notation is an abbreviation for very BIG numbers or for very small numbers.
(Because Schoolnotes.com cannot display exponents, they will be written as: 10e-1, where the “e” denotes that a superscript should follow. I apologize for the confusion this may cause.)
I.e. 36000 = 3.6 x 10e4 = 3.6 x 10 x 10 x10 x 10
Numbers greater than 10 will have a positive exponent.
I.e. 0.0081 = 8.1 x 10e-3 = (8.1) ¸ (10 x 10 x 10) = 8.1 ¸ 10e3
Numbers less than One will have a negative exponent.
A. Multiplication =
(3.0 x 10e4) x (2.0 x 10e2) = (3.0 x 2.0) x 10e(4+2) = 6.0 x 10e6
B. Division =
3.0 x 10e4 ¸ 2.0 x 10e2 = (3.0 ¸ 2.0) x 10e(4-2) = 1.5 x 10e2
C. Adding =
5.40 x 10e3 + 0.60 x 10e3 = 6.0 x 10e3 [Make sure exponents are the same.]
D. Subtraction =
5.40 x 10e4 – 0.40 x 10e4 = 5.0 x 10e4 [Make sure exponents are the same.]
Practice Problems =
1. 5.3 x 10e4 + 1.3 x 10e4 =
2. 7.2 x 10e-4 ¸ 1.8 x 10e3 =
3. 10e4 x 10e-3 x 10e6 =
4. 9.12 x 10e-1 – 4.7 x 10e—2 =
5. (5.4 x 10e4) x (3.5 x 10e9) =
6. (1.2 x 10e2) x (8.9 x 10e2) =
Answers =
1. 6.6 x 10e4
2. 4.0 x 10e-7
3. 10e7
4. 8.7 x 10e-1
5. 1.9 x 10e4
6. 1.1 x 10e5
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Lesson #10 Intro Atomic Theory October 2,3, 2007
Aim: What is the basic structure of an atom?
(Chapter 4 in Textbook.)
I. Models of the Atom
1. No one really knows what an atom looks like!
2. Therefore, we must believe in something we cannot see!
A. The Solid-Particle Model – (400 B.C.) by the Greek Philosophers
1. Democritus - Matter is discontinuous (atomic) and therefore can reach a point where it can’t be divided into smaller parts (This idea is part of the model of matter that we use today.)
2. Plato & Aristotle - Matter is continuous and can always be divided into smaller parts (Trumped Democritus' theory until Dalton.)
3. Atoms are solid particles
B. The Dalton Model – (1803) by John Dalton, English chemist & schoolteacher (Figure 4.2)
1. Matter is made of tiny, indivisible, unchangeable parts called atoms that cannot be destroyed or created
2. Each element is made of atoms that are identical to each other in all properties
3. Properties of one element are different from properties of another element
4. Atoms of different elements can physically mix together or chemically combine in simple whole-number ratios to form compounds
5. Chemical reactions occur when atoms separate, join, or rearrange. Atoms of one element cannot change into atoms atoms of another element as a result of a chemical reaction.
C. Protons - Eugene Goldstein (1886)
1. Did experiments with cathode-ray tubes and saw rays traveling in opposite direction from current that were positively charged
D. The Plum Pudding Model – (1897) by JJ Thompson (Figure 4.4)
1. Discovered the electron during experiments with cathode-ray tubes. Passed electrical current through gases at low pressures.
2. The electron was found to have a negative charge while floating in a “sea” of positive charge
E. The Nuclear Model – (1910) by Ernest Rutherford, former student of Thompson's (Figure 4.7)
1. The Gold Foil Experiment - bombarded gold foil sheets with nuclei of helium atoms and observed particles scattering
2. Discovery of the Atomic Nucleus
a. most of atom’s volume is empty space (particles passed through gold foil)
b. most of atom’s mass is in a very dense (particles ricochet off foil), positively charged (particles were deflected while passing through foil) nucleus
F. The Solar System Model – (1913) by Niels Bohr, Rutherford's Student (Figure 5.2)
1. Electrons move around the nucleus in circular orbits
2. Electrons exist in “allowed” orbits which is why they don’t lose energy and fall into the nucleus
3. Electrons’ orbits (energy level/orbital) have different amounts of energy depending how far away from the nucleus they are
4. The first ‘orbit’ has the least amount of energy
5. The farther away from the nucleus, the more energy the orbit has (excited state)
6. The outermost orbit (Valence Shell) contains the atom’s VALENCE ELECTRONS!!!
a. these valence electrons are important for bonding and determining chemical properties of an element
7. Atomic Emission Spectra/Spectral Lines/Bright Line Spectra
a. All atoms have a “fingerprint” of different wavelengths in the electromagnetic spectrum
i. an electron that is hit by a photon of light, absorbs the light's energy and then jumps to a higher energy level (Excited State = i.e. electron configuration of 2-7-2 for an atom of sodium)
ii. eventually the electron will fall back to its original lower energy level and emit the excess energy it absorbed from the light (Ground State = i.e. electron configuration of 2-8-1 for an atom of sodium)
iii. THIS EMITTED ENERGY corresponds to a specific wavelength in the electromagnetic spectrum
G. The Quantum-Mechanical Model (aka Wave-Mechanical Model) – (1927) by Erwin Schrodinger (Figure 5.5)
1. an electron behaves like both a particle and a wave
2. the electron’s location is described by the “probability” of where it will be located
3. Heisenberg’s Uncertainty Principle – The position and speed of a particle cannot both be known exactly at the same time. As one is known more precisely, the other becomes less certain. (Werner Heisenberg was Bohr's Student.)
H. James Chadwick (1932) - discovered the neutron
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Lesson #11 Intro to Element Symbols October 3, 2007
Aim: What do the Symbols on the Periodic Table Represent?
Motivation: Do you know the letters of the alphabet? It helps you to read better, right? Well, if you don’t know the symbols on the periodic table, then it will be difficult for you to understand reading chemistry.
I. The Language of Chemistry
A. elements - alphabet
B. compounds - words
C. formulas/equations - sentences
D. "reactants" - ingredients to a recipe
"-->" - cooking/baking
"products" - cake
Reactants --> Products
NaCl (s) + H2O (l) --> Na+ (aq) + Cl- (aq)
Solid sodium chloride mixed with liquid water will yield an aqueous solution of sodium ions and chloride ions.
II. Symbols for Equations
A. Substances:
(s) - solid
(l) - liquid
(g) - gas
B. Mixtures:
(aq) - aqueous (substance dissolved in water)
III. Periodic Table - list of all the elements, their weights, and the number of protons of each element in an organized manner.
A. The term “periodic” shows the cyclical patterns found in the elements.
B. Organized by Henry Mosley, b. 1887, so that it is organized by atomic number. He was killed in action in WWI in 1915.
C. shows trends, similarities, and differences of elements
1. has columns and rows.
2. The rows are called periods
3. The columns are called families or groups
4. Each group (column) of elements has similar chemical properties
5. Lists the most stable isotope of an element
IV. Elements - there are more than one hundred elements
A. first 83 have stable, non-radioactive isotopes (they also have radioactive isotopes, but most stable are listed on the table)
B. First 92 elements have been made by fusion nuclear reaction in stars
C. After Element 92 - all elements are "human-made" (read about physicist Glenn Seaborg)
D. Some of the symbols have one letter, some have two, but each element symbol has one and only one upper case letter in it.
1. So you write "Na", NOT "NA"!
2. i.e. Co is the element cobalt, CO is the compound carbon monoxide
3. i.e. Si is the element silicon, SI is the compound sulfur iodide
4. Chemtutor has a Quickquiz to help you learn the names and symbols of the elements = http://www.chemtutor.com/elem.htm
E. There are 2 elements on the Periodic Table that are not found on earth:
1. Pm = #61 Promethium = named after the Titan Prometheus who stole fire from the Gods and gave it to Humans and therefore he was eternally punished by being tied to a rock and having his guts ripped out by birds every day. (Since he was a Titan, he is immortal.) Element was identified by its spectral lines on the sun.
2. Tc = #43 Technetium = used to exist on Earth, but decomposed rapidly after earth’s formation.
Common Elements
Helium He Lithium Li Hydrogen H Sodium (Natrium) Na
Boron B Carbon C Silicon Si Calcium (Lime) Ca
Beryllium Be Fluorine F Neon Ne Sulfur (Brimstone) S
Phosphorus P Nitrogen N Aluminum Al Potassium (Kalium) K
Chlorine Cl Argon Ar Magnesium Mg Iron (Ferrum) Fe
Bromine Br Oxygen O Manganese Mn Copper (Cuprum) Cu
Cobalt Co Nickel Ni Chromium Cr Lead (Plumbum) Pb
Zinc Zn Krypton Kr Rubidium Rb Silver (Argentum) Ag
Iodine I Platinum Pt Cadmium Cd Tin (Stannum) Sn
Cesium Cs Barium Ba Francium Fr Antimony(Stibium) Sb
Bismuth Bi Arsenic As Strontium Sr Tungsten(Wolfram)W
Radon Rn Xenon Xe Polonium Po Gold (Aurum) Au
Radium Ra Uranium U Mercury (Hydrargyrum or Quicksilver) Hg
http://www.chemtutor.com/elem.htm
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Supplemental Lesson: Periodic Table
**(Some of the information here was not included in the class notes because it is not necessary when studying for the Regents.)**
Aim: What are the Basics of the Periodic Table of the Elements?
I. Representative Elements
A. Properties of Metals
1. Metallic bonds result from the easy flow of electrons between the closely spaced energy levels of metal atoms.
2. each metal cation is bonded to a sea of valence electrons.
3. the flow of electrons around the metal atoms gives the properties of malleability (the ability of a metal to be shaped without breaking) and ductility (the ability of metal to be drawn into a wire.)
4. can withstand the stress of being moved against each other since they have electrons in common.
B. Alkali Metals i.e. Lithium and sodium
1. give up one electron in a chemical compound resulting in their high reactivity
2. they are not found in nature, and difficult to obtain in their pure form because they react with air, water, and other substances.
3. soft (can be cut with a knife), have many metallic properties, soluble in water, can conduct heat and electricity, make colorless solutions
4. can form hydroxides with water.
5. Compounds containing hydroxide are alkali or are basic.
C. Alkali Earth Metals i.e. Calcium and barium
1. Alkali Earth Metals, aka alkaline earth metals, give up two electrons in a chemical compound.
2. too reactive to be found in their pure state in nature.
3. denser and not as soft as alkali metals.
4. Magnesium = can be used as a firework because when exposed to a flame, it produces a bright white light.
D. Aluminum
1. third most abundant element in the Earth's crust and usually is found combined with oxygen.
2. was once more expensive that gold or silver b/c difficult to extract in its pure form.
3. good electrical conductor, highly reactive, can form lightweight alloys, and is malleable.
II. Transition Metals
A. Properties of Transition Metals
1. The transition metals, aka transition elements, occupy the space between Group 2 and Group 13
2. common characteristics = color, formation of complex ions, and multiple oxidation states (the number of electrons they give off varies).
B. Some Transition Metals
1. Chromium (Cr)
a. known to prevent corrosion and for the red color in rubies
b. In the 1996 Olympics opening ceremonies, chromium was applied to trucks used in one of the many acts performed to reflect light to illuminate what was happening on the field.
c. can be applied to steel objects to improve their appearance. It can be found in chromite, FeCr2O4.
2. Iron (Fe)
a. fourth most abundant element in the earth's crust.
b. can be found in meteorites.
III. Metalloids
A. Properties of Metalloids
1. called semiconductors and can therefore be used in micro-circuitry, electronics, and computers due to this property.
2. boundary = metalloids: boron, silicon, germanium, arsenic, antimony, tellurium, and polonium. Aluminum and astatine are not metalloids.
C. The Halogens
1. reactive nonmetals in Group 17
D. Noble Gases
1. generally inert gases.
2. An inert material will not react with another material.
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Lesson #12 Avogadro's Number October 4, 2007
Aim: How do I Calculate the Number of Atoms of an Element From Moles?
1 mole = 6.02 x 10e23
= a number NOT a measurement
1 mole of an inert gas will occupy 22.4 L of space at STP (Standard Temperature & Pressure - See Reference Table A).
1) How many atoms of carbon are in 1.5 moles?
1.5 x 6.02x10e23 = 9.03 x 10e23
2) How much does 1.5 moles of carbon weigh?
1 atom of Carbon weighs 12 amu.
1 mole of Carbon weighs 12 grams.
1 mole = 6.02 x 10e23 = 12g of carbon
0.5 moles = 3.01 x 10e23 = 6g of carbon
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1.5 moles = 9.03 x 10e23 = 18g of carbon
or 1.5 moles = x / (12g/mol)
x = 18g
I. Gram Formula Mass = Atomic Mass
- the mass of a mole of an atom in grams
- To obtain the gram formula mass of a compound, you must add up the gram atomic masses of all the individual atoms in the compound.
1) NaCl
Na = 23 g/mol, Cl = 35 g/mol
23 gmol + 35 g/mol = 58 g/mol
2) NaNO3
Na = 23 g/mol, N = 14 g/mol, O = 16 g/mol
23 g/mol + 14 g/mol + 3(16 g/mol) = 85 g/mol
3) Ca(OH)2
Ca = 40 g/mol, O = 16 g/mol, H = 1 g/mol
40 g/mol + 2(16 g/mol) + 2(1 g/mol) = 74 g/mol
II. Converting Grams to Moles
1. Mole Calculations Equation on Table T:
2. # moles = given mass (g) / gram formula mass (g/mol)
i.e. How many moles of NaCl are in 116g of NaCl?
# moles = (116 g)/(58 g/mol) = 2 moles
III. Converting Moles to Number of Particles
i.e. How many NaCl particles are in 2 moles of NaCl?
1 mole = 6.02 x 10e23 atoms
1 mole of NaCl = 6.02 x 10e23 NaCl molecules
2 moles of NaCl = 2(6.02 x 10e23) = 12.04 x 10e23 NaCl molecules
1) How many atoms are there in 1 mole of C6H12O6 molecules?
a) 24 (6.02 x 10e23)
b) 24
c) 12 (6.03 x 10e23)
d) 12
2) How many Helium atoms are in 59 grams of Helium?
3) How many mole of Copper are in 126 grams?
4) A student masses 493 grams of Gold. How many atoms of gold has the student massed?
5) A jeweler has 22 moles of Silver. How many kilograms of silver does she have.
6) How many grams of Selenium are in 5.5 moles?
7) How many moles of Radon are in 239 grams?
8) How many atoms of Cobalt are in 126 grams?
9) How much do 3.73 x 10e24 atoms of Potassium weigh?
10) How much does 1.9 moles of copper(I)sulfate weigh?
copper(I)sulfate = Cu2SO4
2Cu = 2(64)=128g/mol
S = 32g/mol
4O = 4(16)=64g/mol
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Cu2SO4 = 224g/mol
1.9 mol = (x)/(224g/mol)
x=425.6g
11) How many moles are in 176 grams of tin(IV)nitrate?
tin(IV)nitrate = Sn(NO3)4
Sn = 119g/mol
4N = 4(14)=56g/mol
12O = 12(16)=192g/mol
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Sn(NO3)4 = 367g/mol
x = (176g)/(367g/mol)
x = 0.48 mol
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Lesson #13 Lewis Dot Structures October 9, 2007
Aim: How do I draw Lewis Dot Diagrams?
I. Lewis Dot Diagrams –
a. Show how valence electrons are involved in bonding
1. Valence Electrons – electrons in outermost orbit (or energy level) that are involved in bonding
2. Valence Electrons are found in s and p sublevel orbitals.
b. Follow the OCTET RULE = most atoms want 8 e- in their valence shell.
i.e.
Na Cl
2-8-1 2-8-7
Na+ (Na loses 1 electron) Cl- (Cl gains 1 electron)
2-8 2-8-8
Na transfers one electron to Cl to give Na+ and Cl-.
Na+ + Cl- -- > NaCl
Na loses its one valence electron to Cl thereby allowing its third energy level to be vacant and its new valence shell (2nd energy level) to have 8 e-. Cl therefore gains one electron and adds it to its seven valence e- to complete its Octet.
II. Writing Lewis Structures
a. Put element symbol in a box (for practice only)
b. Only 2 valence e- can fit on any one side of the box
c. 2 e- on one side = paired electrons
d. 1 e- on one side = unpaired electron
e. Start filling up orbitals, or sides to the box, with one electron on each side first
f. Then pair up the lonely e- on each side with the remaining e-.
III. The Ground State
a. If an element is said to be "in the ground state", then all of its electrons are found in the lowest energy levels possible.
b. Therefore, when drawing Lewis Structures of Atoms in the Ground State:
1. Pair off the first 2 e-
2. Then put one electron on each of the three sides until you reach the correct number of valence e- for that atom.
IV. The Excited State - when electrons are hit by a photon of light they become "excited" and jump to a higher energy level, and the electron is said to be in the "excited state"
V. Practice
Draw Lewis Structures for the following elements:
1. P
2. Na
3. Al (in the ground state)
4. Mg (in the excited state)
5. Xe
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Lesson #14 Writing Chemical Compounds October 10, 2007
Aim: How do I Use the Elements on the Periodic Table to Write Chemical Formulas? (Part I: Naming Binary Compounds)
I. Intro
A. First letter always capital, second is never, if waiting to be named they have three letters derived from their atomic numbers
B. Most atoms found alone in nature, others with a partner(s):
1. Diatomic Elements: HOFBrINCl
a. These elements are never found alone & in their electrically neutral states
2. Polyatomic: ozone O3, P4, S8, buckyball C60
II. Atoms are electrical in nature and are assigned oxidation numbers (or oxidation states)
i. Ox states can be positive, negative, zero
ii. Ox states show what the electrical charge of the atom will be if it is combined with another atom
iii. Atoms that remain alone (free elements) have an ox # = zero
1. HOFBrINCl
2. noble gases (Group 18 elements)
3. solid metals
4. single elements that are electrically neutral
III. Binary Compounds (rules)
i. Elements with positive ox numbers combine with elements with negative ox numbers (Cation + Anion)
ii. the algebraic sum of the oxidation numbers = 0, this makes the compound electrically neutral
iii. Element with positive ox number (Cation) is usually always written first
iv. i.e. H20 , hydrogen assigned ox # = +1 & oxygen assigned ox # =-2
1. these are the charges of H and O in water
2. you can find these charges on the periodic table in the upper right hand corner of the element box under: Oxidation States
3. H+, O2-
v. Practice:
1. i.e. Na+ + F- yields NaF
2. i.e. Ca2+ + O2- yields CaO
3. i.e. Ca2+ + 2 F- yields CaF2
IV. More Practice Problems
1. Write the following binary compounds:
a. Li+ + O2-
b. Al3+ + Cl-
c. Fe3+ + S2-
d. H+ + N3-
2. Answers
a. Li2O
b. AlCl3
c. Fe2S3
d. NH3 - ammonia
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More Practice
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Write the Compounds and Draw the Lewis Structures for the Following:
1. Calcium + Bromine
2. Sodium + Iodine
3. Cesium + Oxygen
4. Magnesium + Fluorine
5. Potassium + Hydrogen
(Note: Hydrogen will use its negative oxidation state when bonding with any Group 1 metal)
6. Potassium + Sulfur
7. Francium + Bromine
8. Lithium + Chlorine
9. Beryllium + Sulfur
10. Strontium + Oxygen
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Lesson #15 Naming Binary Compounds October 11, 2007
Aim: How do I name Binary Compounds?
I. Binary Compounds use the names of two elements that make up a compound.
1. Element with positive oxidation number is written first (CATION)
2. Element with negative oxidation number is written second (ANION)
3. Suffix –ide added to end of name
II. Examples
1. NaCl is named sodium chloride
2. CaO = calcium oxide
3. AlF3 = aluminum fluoride
III. Popular Anions (Elements with Negative Oxidation Numbers)
1. H- hydride
2. F- fluoride
3. Br- bromide
4. C4- carbide
5. P 3- phoshide
6. I- iodide
7. N3- nitride
8. S2- sulfide
9. O2- oxide
10. Cl- chloride
IV. Elements with more than one positive oxidation number
These elements can form more than one binary compound with the same element.
i.e. Fe 2+ and Fe 3+ can form two oxides: FeO and Fe2O3, both are called iron oxide. – Can tell the difference b/t by naming one Iron (II) Oxide [FeO] and Iron (III) Oxide [Fe2O3].
V. Roman Numerals for Naming Compounds
1. I
2. II
3. III
4. IV
5. V
6. VI
7. VII
VI. Name the following:
1. copper (II) sulfide = CuS
2. iron (III) bromide = FeBr3
3. tin (IV) oxide = SnO2
4. aluminum nitride = AlN
5. MgF2 = magnesium fluoride
6. KH = potassium hydride
7. NiI2 = nickel (II) iodide
8. N2O = nitrogen (I) oxide i.e. nitrous oxide = laughing gas
9. CO = carbon monoxide
10. CO2 = carbon dioxide
11. SO3 = sulfur trioxide
12. CCl4 = carbon tetrachloride
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Lesson #16 Polyatomic Ions October 15, 2007
Aim: How do I name compounds containing Polyatomic Ions?
I. Polyatomic Ion = group of at least 2 atoms with an overall electrical charge
= behaves as a single unit with a fixed oxidation number
II. Some Common Polyatomic Ions
Ref Table E
S2- Sulf-ide
SO3 2- Sulf-ite (electric charge is distributed throughout all atoms in the ion)
SO4 2- Sulf-ate (electric charge is distributed throughout all atoms in the ion)
When you see an electrically neutral compound with 3 or more capital letters, you know that there is a polyatomic ion present in the compound.
III. Name each of the following:
a. CuSO4 = copper (II) sulfate
b. AlPO4 = aluminum phosphate
c. NaHCO3 = sodium hydrogen carbonate
IV. Write the formula for each of the following:
a. potassium nitrate = KNO3
b. lead (II) carbonate = PbCO3
c. ammonium acetate = NH4C2H3O2
d. magnesium chlorite = Mg(ClO2)2
V. More Practice Problems: Name each of the following =
a. NaI sodium iodide
b. FeCl3 iron (III) chloride
c. P4O10 phosphorus (V) oxide
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Lesson #17 Basics of Bonding October 22, 2007
Aim: What are the different types of bonds between atoms?
(Intramolecular forces - forces within a molecule, forces between single atoms)
I. Chemical bonding involves the transfer or sharing of valence electrons between two or more atoms.
A. Ionic Bonds – one or more valence electrons are transferred from one neutral atom or group of atoms (polyatomic ions) to another, resulting in a cation (positively charged particle as a result of the loss of electrons) and an anion (negatively charged particle as a result from the gain of electrons).
- Opposite charges attract, cation and anion bind together through an electrostatic attraction.
- Ionic bonds are strong, poor conductors as solids, good conductors as liquids, soluble in polar solvents, i.e. NaCl.
- Metal + Non-metal
B. Covalent Bonds (Molecular Bonds)– sharing of valence electrons.
- When the electron pair is evenly distributed between bonding atoms resulting in a non-polar covalent bond.
- When the electrons are held more closely by one atom over the other, then there is a polarity that results and the bond is a polar covalent bond.
- Covalent bonds are very strong, non-polar covalent substances are insoluble in polar solvents (i.e. oil (non-polar) and water (polar) don't mix) and are unable to conduct electricity.
- Non-metal + Non-metal
C. Metallic Bonds - Metal + Metal
– Valence electrons are not associated with any specific atom or pair of atoms, but wander freely among adjacent metal atoms. They are positive ions in a sea of valence electrons. This free-electron character (mobile electron) of a metal explains why metals are excellent conductors of heat and electricity.
- Transition Metals (Groups 3-12) form compounds that make colored solutions i.e. paint
- Weak bonds
D. Coordinate Covalent Bonds - When both bonding electrons being shared between two atoms come from only one of those atoms.
i.e. NH3 + H+ --> NH4+ Nitrogen donates its electron pair to a Hydrogen ion.
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Lesson #18 Electronegativity October 24, 2007
Periodic Trends Lesson 1 of 3
Aim: What is Electronegativity?
I. Electronegativity – describes the ability of an atom, in a compound, to attract electrons (Reference Table S)
a) measures how hungry an atom is for electrons, the higher the electronegativity - the hungrier the atom
b) Increases as you move down a group
c) Increases as you move from left to right across a period
d) Increases as you move from the lower left corner of the Period Table to the upper right corner of the Periodic Table
II. Bonds & Electronegativity =
(This technique works best with binary compounds. Otherwise you have to examine the compound's geometric shape and symmetry.)
1. Ionic - if the difference in electronegativities between the bonded atoms is 100%
- if difference between electronegativities is 1.7 or greater, then the bond is ionic
2. Covalent - if both atoms have identical electronegativities
- if the difference is between 0.3 and 1.7, then it is polar covalent
- if the difference is less than 0.3, then it is non-polar covalent.
III. Examples
1. Cl2 3.2(Cl) - 3.2(Cl) = 0 non-polar covalent
2. H2O 3.5(0) - 2.1(H) = 1.4 polar covalent
3. NaCl 3.2(Cl) - 0.9(Na) = 2.3 ionic
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Lesson #19 Atomic Radius October 29, 2007
Periodic Trends Lesson 2 of 3
Aim: What is an Atomic Radius?
- Tells us how big or small an atom is (Reference Table S)
I. Atoms have roughly a spherical shape
1. spheres have radii
2. chemists take half of the internuclear distance, (from center of one atom to center of another) b/c they cannot measure radius of a single atom
II. Radius increases as you move down a group
1. as period number increases, then more electrons fill more principle occupied energy levels
III. Radius decreases as you move from left to right across a period
1. as group number increases, the number of protons increases and so does the positive charge of the nucleus
2. number of electrons also increases, but since the electrons are still in the same energy level, and the number of principle occupied energy levels stays the same, then the electrons can be pulled closer toward the nucleus, therefore decreasing the radius of the atom.
IV. Ionic Radius
1. Cations = If an atom loses an electron, its radius will decrease in size
2. Anions = If an atom gains an electron, its radius will increase in size
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Lesson #20 Ionization Energy October 30, 2007
Periodic Trends Lesson 3 of 3
Aim: What are Ionization Energies?
I. First Ionization Energy = measures the ease with which the atom loses its first electron to become an ion.
1. There is an inverse relationship between size of the atom (atomic radius) and the 1st Ionization Energy.
2. The larger the atom, the more weakly held the outer e- are, the more easily the outer e- can be removed from the valence shell
3. As you go down a group, ionization energy decreases
4. As you go across a period from left to right, ionization energy increases
5. Increase Ionization Energy = difficult to remove an electron (these atoms mostly form anions).
6. Decrease Ionization Energy = easy to remove an electron (these atoms mostly form cations).
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Lesson #21 Intermolecular Forces November 2, 2007
Aim: What are Intermolecular Forces?
I. Intermolecular Forces - forces of attraction between molecules/compounds and whose strength determines the state/phase of matter a substance will exist as, such as a solid, liquid, or gas
1. weaker than the covalent bonds, ionic bonds, and metallic bonds.
2. act between the positive region of one polar molecule and the negative region of another.
3. the crystals that form as a result of these bonds, generally do not conduct electricity.
4. all substances would exist in gaseous form if it were not for IM forces.
A. Ion-Dipole Interactions - Ions are attracted to a polar molecule i.e. Na+, Cl- in H2O
B. Dipole-Dipole Interactions - Polar molecules are attracted each other (i.e. H2O (l) and H2O (s))
C. Weak Van der Waals Forces (aka "London Dispersion Forces")
1. weakest of all IM forces.
2. They consist of instantaneous dipole moments that charge a molecule for just an instant.
3. All substances have Van der Waals forces.
4. Non-polar compounds can exist as a liquid (i.e. hexane (C6H14)) or a solid because of these forces.
D. Hydrogen Bonding - dipole-dipole interactions between compounds containing hydrogen
– bond between ions and/or polyatomic ions that involves bonding with hydrogen. i.e. Water
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Lesson #22 Average Isotopic Mass November 19, 2007
Aim: What is Percent Composition By Mass, Average Isotopic Mass, and Percent Error?
Motivation: Intro to Nuclear Chemistry - Alexander Litvinenko & Polonium
I. Percent Composition By Mass
1. Percent Composition equation on Ref Table
2. % composition by mass = mass of part/mass of whole x 100%
i.e. What is the percent composition by mass of sodium in sodium chloride?
% composition by mass = (23 g/mol) / (58 g/mol) X 100% = 39.66%
Therefore, 39.66% of the mass of NaCl belongs to just Na.
II. Average Isotopic Mass = is the average of the atomic masses of all the element's isotopes, as found on Earth, that is weighted by the isotope's percent abundance in nature.
i.e. #1
Isotope Atomic Mass (amu) % Natural Abundance
Neon-20 19.99 90.9
Neon-21 20.99 0.3
Neon-22 21.99 8.8
19.99(0.909) + 20.99(.003) + 21.99(0.088) = 20.17 amu = This value is closestto the mass of the more abundant isotope.
i.e. #2
Isotope Atomic Mass (amu) % Natural Abundance
Magnesium-24 23.99 79.3
Magnesium-25 24.99 10.1
Magnesium-26 25.98 10.6
i.e. #3
Isotope Atomic Mass (amu) % Natural Abundance
X-10 10.0 10.0
X-11 11.0 20.0
X-12 12.0 70.0
i.e. #4
What is the average isotopic mass of element X if its masses by abundance are as follows:
20% X-20
20% X-21
10% X-22
50% X-23
Therefore:
0.20 x 20 g/mol = 4.0 g/mol
0.20 x 21 g/mol = 4.2 g/mol
0.10 x 22 g/mol = 2.2 g/mol
0.50 x 23 g/mol = 11.5 g/mol
4.0 g/mol + 4.2 g/mol + 2.2 g/mol + 11.5 g/mol = 21.9 g/mol
III. Percent Error
1. Percent Error equation on Ref Table
2. % error = (measured value - accepted value)/accepted value x 100%
i.e. If a student masses a beaker and finds it weighs 100g, and the teacher says its true mass is 105g, then what is the student's % error?
% error = (100g - 105g)/105g x 100% = 4.7% error
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Lesson #23 Nuclear Chemistry I November 27, 2007
Aim: What is Nuclear Chemistry and Radioactive Decay?
**You will be using Table N & Table O for this unit.
I. Nuclear Chemistry - involves changes in the atomic nucleus
A. Natural Radioactivity
1. Elements #1-92 were all created in stars
2. Elements 93+ were all created in a laboratory (aka Transuranium Elements)
3. Elements #1-83 have at least one stable, non-radioactive isotopes. They also have unstable, radioactive isotopes, but only the most stable isotope is shown on the Periodic Table.
i.e. carbon-14 vs. stable carbon-12
i.e. hydrogen-3 vs. stable hydrogen-1, etc.
4. Elements #84+ are all radioactive and have no stable, nonradioactive isotopes. However, the most stable radioactive isotope of the element is shown on the Periodic Table.
5. When an element is radioactive, it will undergo a spontaneous disintegration of its atomic nucleus
6. Transmutation = the process in which an element changes into another element
II. Radioactive Decay = radioactive nuclei break down naturally in a series of steps to “correct” the ratio of neutrons to protons to make the atom stable
1. Parent & Daughter Nuclei
a. The Radioactive Element going through decay is the PARENT NUCLEUS.
b. The Element the PARENT transmutates into is called the DAUGHTER NUCLEUS.
2. When Mass of Radioisotope vs. Time is graphed, an exponential decay curve is revealed
3. There is only one paricle on the reactants side of a natural transmuation equation
A. Alpha Decay - Isotope of Element loses subatomic particles equivalent to a Helium nucleus or alpha particle. Decrease atomic weight by 4 amu and atomic number by 2 amu.
B. Beta Decay (B-) - Isotope of Element loses an electron. Atomic Weight stays the same, increase atomic number by 1 amu.
1. This happens because a neutron breaks down into a proton and an electron.
2. An electron = A Beta Particle
C. Positron Decay (B+) - Isotope of Element loses a positron. Atomic Weight stays the same, decrease atomic number by 1 amu.
1. Positron = anti-matter of an electron = positively-charged electron
D. Gamma Decay
Thorium-230 --> Radium-226 + Helium-4 + Gamma Particle
1. high-energy photon is emitted
2. gamma particle has no mass & no charge
3. high penetrating power and very dangerous
a. can be stopped by several meters of concrete or several centimeters of lead
Alpha & Beta - Decay:
http://dbhs.wvusd.k12.ca.us/webdocs/Radioactivity/Writing-Alpha-Beta.html
Beta + Decay:
http://dbhs.wvusd.k12.ca.us/webdocs/Radioactivity/Writing-Positron-EC.html
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Lesson #24 Nuclear Chemistry II November 30, 2007
Aim: What is a half-life period?
I. Half- Life – time required for a substance to decay to one-half of its initial value
i.e.#1= If there is a 100g sample of a radioactive element and the half-life of this element is 5 days, then after 5 days half of the sample (50g) will decay and the other half (50g) will still be radioactive element. In another 5 days there will be half of this amount or 25g of original sample left.
i.e.#2= 1.00g of C-14 --->5730 years---> 0.500g of C-14 --->5730 years---> 0.250g C-14 --->5730 years---> 0.125g C-14
It takes 3 half-lives, 17,190 years, for 1.00 gram of C-14 to decay to 0.125 grams of C-14
A. There are two equations on Table T of the Reference Tables which will help you answer questions on Half-Life.
1. Fraction Remaining = (1/2)t/T
2. Number of Half-Life Periods = t/T
t = total time elapsed
T = half-life
B. Practice Problems
1. What fraction of a sample of Cr-51 will remain after 168 days?
# half-lives = 168 days/28 days per half-life = 6 half-lives
fraction remaining = (1/2)6 = 1/64th of original sample
2. If a sample of Cr-51 has an original mass of 52.0g, then what mass will remain after 168 days?
mass remaining = original mass x fraction remaining
= 52.0g x 1/64 = 0.813g
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Aim: Nuclear Chemistry Practice Problems December 3, 2007
1) Polonium-210 has a half-life of 138.39 days.
Suppose 1 gram of Po-210 was used to poison Alexander Litvenenko. How many days will have passed if 1/64th of the original sample still remains?
2) What is the mass of the remaining Po-210 in Question #1?
3) Po-210 is considered an "alpha emitter". What is the daughter nucleus for the decay reaction of Po-210?
4) If 3.0g of Sr-90 in a rock sample remained in 1999, approximately how many grams of Sr-90 were present in the original rock sample in 1943?
5) What is the mass of K-42 remaining in a 16g sample of K-42 after 37.2h?
6) What is the total mass of Rn-222 remaining in an original 160mg sample of Rn-222 after 19.1 days?
7) In 6.2h, a 100g sample of Ag-112 decays to 25.0g. What is the half-life of Ag-112?
8) A sample of I-131 decays to 1.0g in 40 days. What was the mass of the original sample?
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Lesson #25 Fission & Fusion December 4, 2007
Aim: What are Fission and Fusion reactions?
I. Fission - a form of artificial transmutation that involves the splitting of a heavy nucleus to produce lighter nuclei
A. When a heavy element captures a neutron, it gets split into smaller elements and more free neutrons are released with energy.
1 neutron + U-235 --> Ba-142 + Kr-91 + 3 neutrons + energy (4.9 x 10e9 kcal/mole)
1. U-235 & Po-239 are the only fissionable isotopes
2. 1 kg of U-235 will release energy equivalent to 20,000 tons of dynamite
3. Matter is converted to energy
B. Chain Reaction - a nuclear reaction (a controlled fission reaction) in which one step supplies energy or reactants for the next step
i.e. 3 neutrons produced from U-235 fission reaction can then split 3 more U-235 nuclei that will release 9 more neutrons & energy, etc.
C. Critical Mass - minimum quantity of fissionable material needed for a chain reaction to occur
1. Detonation of an Atomic b0mb - uncontrolled fission reaction
II. Fusion - involves combining of light nuclei to produce a heavy, more stable nucleus at high temperatures and pressures
A. Occurs naturally in the sun
B. When matter is converted into energy
C. When atoms, such as hydrogen, combine to form larger atoms.
Hydrogen-2 + Hydrogen-2 --> Helium-4
Hydrogen-1 + Hydrogen-1 --> Hydrogen-2 (deuterium)
Hydrogen-1 + Hydrogen-2 --> Helium-3
III. Some Uses for Radioisotopes
A. Carbon-14 = used in the study of organic materials and mechanisms
B. Iodine-131 = used in the diagnosis of thyroid disorders
C. Technetium-99 = used in the diagnosis of brain tumors
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Lesson #26 Real Gases Vs. Ideal Gases December 7, 2007
Aim: What is the difference between Ideal and Real Gases?
*Even though ideal gases do not exist, the concept of an Ideal Gas is important because it sets a standard for a gas that follows the gas laws at ALL conditions of pressure and temperature. It is also used to explain the behavior of a gas sample.
I. Ideal Gases
1. particles have mass but negligible volume
2. particles move randomly and in straight, parallel lines (therefore never collide with each other)
3. particles have no intermolecular forces of attraction between them
4. particles undergo perfect elastic collisions in which energy is not lost
5. The absolute Kelvin temperature is directly proportional to the average kinetic energy of the particles
II. Real Gases
1. volume occupied by gas is equal to the volume of its container
2. particles are easily compressed because of the space between them
3. particles move randomly (straight lines, perpendicular lines, wavy lines, etc.)
4. particles have weak intermolecular forces of attraction between them (Van der Waals/London Dispersion Forces)
5. collisions can be elastic or non-elastic
6. particles behave like an Ideal Gas ONLY AT high temperatures and low pressures
7. Noble Gases behave the most like Ideal Gases
III. Pressure = force/area
1. Pressure is the pushing force of particles on each other and on container walls during collisions(elastic or non-elastic).
2. 1 atm = 760 torr = 760 mm Hg = 101.325 kPa (pascal is the metric measurement) = pressure at sea level = standard pressure
3. Evangelista Toricelli = said that air exerts a pressure
a. created 1st barometer – long tube, closed at one end, and filled with mercury and then inverted into a pool of mercury. (Always stayed in tube at 760 mm = atmospheric pressure [760 torr])
4. The number of particles, volume, and temperature all affect the pressure of a gas.
IV. Practice Problems
1. What real gas comes closest to having characteristics of ideal behavior?
2. Why do real gases deviate from ideal behavior?
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Lesson #27 Boyle's Law December 10, 2007
Aim: What is Boyle's Law?
I. Boyle’s Law (Robert Boyle) – inverse relationship between pressure of gas and its volume at constant temperature
a. PV = k , P = pressure, V = volume, k = constant at specific temp for a certain sample of air
b. Since using same gas = P1V1= P2V2
For example: P1 & V1 are the pressure and volume of a gas before it expands
P2 & V2 are the pressure and volume after expanding
I.e. If P1 = 10 atm, V1 = 3L, V2= 7L, and you need to find the new pressure, P2
P2 = 4.3 atm
I.e. As the volume of a gas at STP doubles, what will be its new pressure?
Due to the inverse relationship between volume and pressure, as Volume doubles, Pressure halves.
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Lesson #28 Charles’ Law & Guy Lussac's Law December 7, 2007
Aim: What is Charles’ Law and what is Guy-Lussac's Law?
I. Jacques Charles – volume of a gas increases linearly with its temperature, at constant pressure
V = bT
V = volume, b= a proportionality constant, T= temp in Kelvin
- Temp is measured in Kelvin => K= 0C +273
ALWAYS convert temperature to Kelvin when working on problems involving gases.
(One reason for this rule is because if one temperature is 0 degrees Celsius, then you will have a zero in a denominator for one of your fractions and end up with an answer that is undefined. Kelvin temperatures are always positive values above zero.)
II. Practice problems:
1. A balloon is inflated to 5L at 20 degrees C. What will be the new volume of the balloon if it is put in a freezer at 0 degrees C?
V1/T1 = V2/T2
(5L)/(293K) = (V2)/(273K)
2. If you took a balloon outside that was originally at 20 C at 2 L in volume, and it heated up to 29 C, then what would its new volume be?
Examples:
The equation V = bT can be rearranged to V/T = b. Now, you know that since you are dealing with the same gas at constant pressure (always assume constant pressure, unless if the problem tells you otherwise), the constant, b, will remain the same. So:
V1 / T1 = b = V2 / T2
or V1 / T1 = V2 / T2.
V1 = 2 L, T1 = 20 C, T2 = 29 C, and you must solve for V2.
WAIT!!! You have to convert the temperatures to Kelvin first!! So:
T1 = 20 C = 20 C + 273 = 293 K
and T2 = 29 C = 29 C + 273 = 302 K
Now, you can plug in the numbers and solve for V2. (2 L / 293 K) = (V2 / 302 K), and V2 = 2.06 L
3. What volume will a 300mL gas sample at STP occupy when the temperature is doubled?
III. Guy-Lussac = direct relationship between pressure and temperature, at constant volume
a. P1/T1 = P2/T2
b. Explains why your car tires may blow if pressure builds too much | |
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