Name:______
Date:______Per. #:______
AP Chemistry – Summer 2015
SUMMER ASSIGNMENT GET TEXTBOOK: Wed., June 3, 3:20 pm
Due: The first day of class, August 19.
I.Memorize all element symbols, ions (names, formulae, and charges),nomenclature of ionic and covalent compounds, solubility rules, and flame test colors on attached charts. There will be a test on these the second day of class. If you are unable to score at least an 80% on the test you may be recommended to drop the course. I strongly recommend you make flashcards and use them.
II.Read Chapters 1 & 2 in your textbook, take notes on the reading, and answer all the blue questions at the end. Check your answers in the appendix in the back of the book. I will expect you to understand all the material in Ch. 1 and 2 by the first day of school and some topics may appear on the test on the second day of school.
III. Use your textbook to complete the following problems. In this packet are explanations of chemical nomenclature and certain chemistry math topics as well as practice problems to supplement the work you are doing from the textbook.
IV.Join the class Facebook page. Search for “SLHS AP Chemistry: 2015-2016”. We use Facebook frequently for posting related videos and tutorials, problem sets, and answers to worksheets and study guides.
V.Mr. Rea’s email is and his phone number is 805-450-2021. Email or connecting through Facebook is preferred. Don’t call before August and call at a reasonable hour.
COURSE EXPECTATIONS:
AP chemistry is a fast-paced, challenging course for those students who can consistently work hard, pick up ideas quickly, and manage their time effectively. You will need to keep yourself on task, make certain you understand each day’s ideas before you leave the classroom, come to tutoring when you need help, and work on chemistry almost every day. Don’t tell yourself “I don’t have chemistry homework.” You may have a day where there is nothing due the next day, but you always have your book to read, you always have current and old chapters to review, you have online videos you can watch for enrichment, and you should have anAP prep book with which to practice.
Much like a college course, you are expected, to a large extent, to teach yourself the material. Class lectures provide enrichment, clarification, and a different perspective, but the best chemistry students are self-motivated and pursue knowledge and skill on their own without having to be told.
If you are willing to put in the time and effort every day, I will work with you to make certain you survive AP chemistry, maybe even get an A (it has been done!), maybe even get a 5! If you are smart, but lazy (you know who you are) your choices are to change your work habits or to ask your counselor to switch to CP chemistry or a different science class. Success in AP chemistry certainly requires intelligence, but perseverance and consistent effort are requirements as well.
Materials:
- a scientificcalculator (graphing recommended)
- chemistry binder with 5 tabbed dividers and a 100-sheet notebook or lined notebook paper
- A second notebook for lab data entry
- An AP Prep Book, think Princeton Review, Barron’s, etc. (Recommended but not required.)
Money:
All AP chemistry students are required to take the AP chemistry exam on Monday, May 2nd, 2016. It will cost a little over $90. (Warn your parents and save up now.)
I am assuming that you will be coming into this class in August with various scientific/mathematical skills. Some of these will be prior knowledge and some you will learn during the summer assignment. Review and prepare for the following:
Metric System
You know the metric system.
You know the meaning of the metric prefixes, kilo-, centi-, milli-, micro-, and nano-.
You know that there are other metric prefixes and can look them up if needed (mega, pico, etc.)
You can convert one measurement into another (e.g., 0.532 cg = ______mg).
You can convert squared or cubed units (e.g., knowing that 2.54 cm = 1 inch, 38.5 in2 = _____ cm2).
Dimensional Analysis & Showing Your Work
When you convert one unit to another, you can show your work using dimensional analysis or unit analysis.
You know that good examples of dimensional analysis are changing metric units, converting time units, or using density to convert mass to volume or volume to mass.
You know that you should always show enough work so that if your answer is incorrect, I can tell where you went wrong.
Scientific Notation
You can translate regular numbers into scientific notation and numbers written in scientific notation into normal notation.
You know the distinction between exponential notation and scientific notation.
Making Measurements
You can use a ruler or other measuring device to make a measurement to the correct number of significant figures, i.e. include all of the digits in the measurement that are a significant part of the measurement.
You can correctly assign a value when making a given measurement.
You always include a unit on a measurement or final answer.
You know the distinction between a measurement and a defined number (e.g., 12 things in a dozen, pi).
You can explain the difference between accuracy (how close a measurement is to a true or accepted value) and precision (how close a set of measurements are to each other).
Significant Figures
You can determine the number of significant figures in a given measurement (i.e., You know whether a “0” in a measurement is significant or not.)
You can determine the precision in a calculation involving measurements when the measurements are written with the correct number of significant figures.
You can determine the precision in a calculation involving measurements when the measurements are written with notation.
AP Chemistry – Summer
Nomenclature Rules
I. Covalent Nomenclature
When naming covalent molecules, you must first look at the symbol and understand what all of the parts mean. We’ll use the example:
C2H4
The little numbers, called subscripts, tell you how many of each atom there are. Subscripts work on the element that comes before them. So in this example there are two (2) carbon atoms and four (4) hydrogen atoms in the molecule.
After you’ve got that straight, follow three easy steps to name the compound.
- Write the name of the more metallic element first.
- The least metallic elements are in the top right corner of the Periodic Table. The further down and to the left you go, the more metallic. If one element is a solid at room temperature and one is a gas, the solid is considered more metallic.
- Add PREFIXES to all elements.
- The one exception is that you never use mono- for the first element.
- Add –ide as a suffix to the last element only.
That’s it! If you can do that, you can name any covalent molecule. Let’s practice with C2H4.
- Carbon is more metallic, even though it is closer to the right, because it is solid.
At this point, the name of the molecule is carbon hydrogen.
- Check out the prefix table below:
1 = mono / 2 = di / 3 = tri / 4 = tetra / 5 = penta
6 = hexa / 7 = hepta / 8 = octa / 9 = nona / 10 = deca
Since we found that there are 2 carbons and 4 hydrogens in the molecule, the name at this point is dicarbontetrahydrogen.
- All wee need to do now is use –ide as a suffix at the end of the second element, hydrogen. The final name is dicarcon tetrahydride.
- Yay!
Here are a couple examples:
NO2 nitrogen dioxide PH3 phosphorus trihydride N2O4 dinitrogen tetroxide
Try these ones!
C3H8 / NH3 / H2O / N2O5II. Ionic Nomenclature
The rules to naming ionic compounds are similar to those for covalent molecules, but even easier because you DON’T USE PREFIXES. This is because ions can only form in specific ratios based on their charges, so the prefixes are considered redundant. For example:
When making a covalent molecule with nitrogen and oxygen, there are a few different ways that they can share electrons, so you can end up with N2O or NO2.
When making an ionic compound with calcium and fluorine, calcium always makes +2 cations and fluorine always make –1 anions. So the only possible ratio is CaF2.
Here are the rules:
- Name the metal first.
- Add –ide to the monoatomic nonmetal anion.
- If the anion is polyatomic (multiple atoms in a charged molecule, like NO3–), simply name it.
- NO PREFIXES
See, pretty easy. Check out these examples:
MgCl2 magnesium chloride Al2O3 aluminum oxide LiF lithium fluoride
Try these!
MgO / CaBr2 / Na2SO4 / K3NThe harder thing to do, since there are no prefixes, is to go backward and get the formula units of an ionic compound from the name. First: you need to be able to figure out the charge that each ion will be from the periodic table.
Group / 1A / 2A / 3A / 4A / 5A / 6A / 7A / 8ACharge / +1 / +2 / +3 / No ions / –3 / –2 / –1 / 0
For transition metals, you get a break because they always tell you the metal ion’s charge in the name. For example, in iron (III) oxide, the Roman numeral “III” tells you that the iron ion has a +3 charge.
Once you get the charges figured out, the second thing you need to do is switch the numbers, make them subscripts, and reduce to the simplest ratio. It’s easier than it sounds, look at the example below:
Aluminum oxide from the periodic table you get Al3+O2– switch numbers, make subscripts
Al2O3
Try these!
Calcium oxide / Iron (III) oxide / Magnesium fluoride / Lithium chromatePART 1: Nomenclature Practice
1) Write the names of the following covalent molecules. Remember to USEPREFIXES.
Symbol / a) H2S / b) CH4 / c) N2OName
Symbol / d) XeF6 / e) C4H8 / f) BrF
Name
Symbol / g) Cl2 / h) HF / i) P2I4
Name
2) Use the name of each covalent molecule to write the correct chemical formula.
Name / a) carbon dioxide / b) diphosphorus pentoxide / c) pentacarbon decahydrideFormula
Name / d) phosphorus trihydride / e) dibromide / f) dinitrogen trioxide
Formula
Name / g) dicarbon hexahydride / h) silicon tetrahydride / i) carbon monoxide
Formula
3) Write the names of the following ionic formulae. For those containing transition metals, write two names, one with the Roman numeral denoting the charge, and one with the special Latin name (for example: “copper (II)…” and “cupric…”).
Symbol / a) Fe2O3 / b) PbCl4 / c) KBrName
Symbol / d) CaF2 / e) FeS / f) AgNO3
Name
Symbol / g) CuSO4 / h) Mn2S3 / i) Sr(OH)2
Name
4) Use the name of each ionic compound to write the correct chemical formula.
Name / a) magnesium bromide / b) aluminum nitride / c) iron (II) oxideFormula
Name / d) cuprous bicarbonate / e) mercuric chloride / f) sodium sulfate
Formula
Name / g) cobalt (II) nitrate / h) potassium permanganate / i) antimony (V) phosphate
Formula
More Chemical Formulae Practice: Complete the following table. (READ Section 2.8)
Helpful website:
Write formulae for the following: / Name each of the following:a. barium sulfate / a. CuSO4
b.ammonium chloride / b. PCl3
c. chlorine monoxide / c. Li3N
d. silicon tetrachloride / d. BaSO3
e. magnesium fluoride / e. N2F4
f. sodium oxide / f. KClO4
g. sodium peroxide / g. NaH
h. copper (I) oxide / h. (NH4)2Cr2O7
i. zinc sulfide / i. HNO2
j.potassium carbonate / j. Sr3P2
k. hydrobromic acid / k. Mg(OH)2
l. perchloric acid / l. Al2S3
m. lead (II) acetate / m. AgBr
n.sodium permanganate / n. P4O10
o. lithium oxalate / o. HC2H3O2
p. potassium cyanide / p. CaI2
q. iron (III) hydroxide / q. MnO2
r. silicon dioxide / r. Li2O
s. nitrogen trifluoride / s. FeI3
t.chromium (III) oxide / t. Cu3PO4
u. calcium chlorate / u. C2H6
v. sodium thiocyanate / v. NaCN
w. nitrous acid / w. HF
PART 2: Dimensional Analysis, Units and Prefixes, and Density (READ Sections 1.3, 1.6 and Appendix 6)
As AP chemistry students, you have two goals with problems. First, get the correct answer. Second, be able to show others WHY your answer is correct. Dimensional analysis meets both of these goals.
Dimensional analysis is always a given value and one or more conversion factors that allow you to determine the desiredunknown value.
Any mathematical fact can serve as a conversion factor.
1 hour = 60 minutes or
Ex.Convert 1.25 years to seconds:
1.Convert 2.83 days into seconds.
2.Convert 7.72 years into days.
3.Convert 0.0035 weeks into seconds.
4.Convert 180 days into minutes.
5.Convert your age into seconds
Density is often used as a conversion factor between the mass and volume of a sample. For example, the density of liquid mercury is 13.6 g/mL.
6.What is the volume of a 175-gram sample of mercury?
7.What is the mass of 1.00 gallon of mercury? [1 cup = 236.588 mL]
8.When I carry in a 5-gallon container of water from my car, I always wonder its weight. I looked it up on the web and found that 1 lb = 0.453542 kg and 1 qt = .946353 Liter. Calculate its weight in pounds.
Use Table 1.2 to find the definitions of unit prefixes to solve the following problems:
9. How many nanometers is 0.0357 m?
10. How many grams is 2489 micrograms?
11. How many centiliters is 98 L?
12. How many meters is 0.119 km?
13. How many nanograms is 6.22 10–4 g?
14. How many cL are in 1.01 102 kL?
15. How many μL are in 2.26 10–2 mL.
16. What is the density of a piece of wood with a mass of 3.50 kg and a volume of 5500 cm3?
17. If a rock has a density of 5.5 103 mg/mL, what is the density in g/L?
18. The highest temperature on the Earth’s surface ever recorded was 134°F, in Death Valley, CA. What is this temperature in °C and K? (See Appendix 6 in the back of the book for the conversion factors.)
PART 3: Significant Figures and Uncertainty (READ Sections 1.4,1.5, and Appendix A1.5)
A. When takingmeasurements all certain digits plus the first uncertain number are significant.
Example: Your bathroom scale weighs in 10 Newton increments and when you step onto it, the pointer stops between 550 and 560. Your look at the scale and determine your weight to 557 N. You are certain of the first two places, 55, but not the last place 7. The last place is a guess and if it is your best guess it also is significant.
B. When given measurements in a problem, the numbers that are significant are the digits 1-9 and the 0 when it is not merely a placeholder.
1. When 0’s are between sig. fig., 0’s are always significant.
Example: 101 has 3 sig. fig. and 34055 has 5 sig. fig.
2. When the measurement is a whole number ending with 0’s, the 0’s are never significant.
Example: 210 has 2 sig. fig. and 71,000,000 also has 2 sig. fig.
3. When the measurement is less than a whole number, the 0’s between the decimal and other significant numbers are never significant (they are place holders).
Example: 0.0021 has 2 sig. fig. and 0.0000332 has 3 sig. fig.
4. When the measurement is less than a whole number and the 0’s fall after the other significant numbers, the 0’s are always significant.
Example: 0.310 has 3 sig. fig. and 0.3400 has 4 sig. fig.
5. When the measurement is less than a whole and there is a 0 to the left of the decimal, the 0 is not significant.
Example: 0.02 has only 1 sig. fig. and 0.110 has 3 sig. fig.
6. When the measurement is a whole number but ends with 0’s to the right of the decimal, the 0’s are significant.
Example: 20.0 has 3 sig. fig., 18876.000 has 8 sig. fig.
In case 4 and 6 the 0’s have no effect on the value (size) of the measurement. Therefore, these 0’s must have been included for another reason and that reason is to show precision of the measurement. Since these 0’s show precision they must therefore be significant.
In cases 2 and 3 removal of the 0’s DO change the value (size) of the measurement, the 0’s are placeholders and are thus not significant.
In case 5 the 0 is completely unnecessary, it is neither a placeholder nor adds to the accuracy of the measurement.
UNCERTAINTIES IN CALCULATIONS
1. When adding or subtracting numbers written with the notation, always add the uncertainties and then round off the value to the largest significant digit. Round off the answer to match.
Example: (22.4 .5) + (14.76 .25) = 37.16.75 = 37.2 .8
The uncertainty begins in the tenths place… it is the last significant digit.
2. When adding or subtracting numbers written in significant figures, show the uncertainty by rounding the answer to match the largest place with uncertainty.
Example: 267 + 11.8 = 278.8 = 279
The least accurate original measurement is only accurate to the ones place.
3. When multiplying or dividing measurements written in significant figures, show the uncertainty of your calculations by rounding off your answer to match the same number of significant figures as your least precise measurement (the measurement with the least number of significant figures).
Example: 477.85 32.6 = 14.657975 = 14.7
32.6 is the least accurate measurement with only 3 significant figures.
NOTE: There are two types of precision: “absolute precision” and “relative precision.”
Example: 322.45 x 12.75 x 3.92 = 16116.051 = 16100
All the measurements are accurate to the hundredth place (absolute precision) but the answer is rounded to 3 significant figures because 3.92 has only 3 significant figures (relative precision).
In Summary:
Adding and Subtracting / Multiplying and dividing#’s with notation / Rule 1 / Don’t Do This Case
#’s with significant figures / Rule 2 / Rule 3
A. Indicate the number of significant figures then round each to the number of significant figures indicated.
For example:
1.234has ______4___ significant figures and, rounded to2 significant figures, is ___1.2____
1.0.6034has ______significant figures and, rounded to2significant figures, is ______
2.12,700has ______significant figures and, rounded to2 significant figures, is ______
3.12,700.00has ______significant figures and, rounded to1significant figures, is ______
4.0.000983has ______significant figures and, rounded to2 significant figures, is ______
5.123342.9has ______significant figures and, rounded to5 significant figures, is ______
6.6.023 x 1023has ______significant figures and, rounded to2 significant figures, is ______
7..005600has ______significant figures and, rounded to1 significant figures, is ______
8.10000.5006has ______significant figures and, rounded to5 significant figures, is ______
9.2.0 x 10-3has ______significant figures and, rounded to1 significant figures, is ______
10.3.456110has ______significant figures and, rounded to3 significant figures, is ______
B. Given calculations with the calculator answer, write the answers with the appropriate number of significant figures.