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Ch. 5 Notes---Scientific Measurement
Qualitative vs. Quantitative Qualitative measurements give results in a descriptive nonnumeric form. (The result of a measurement is an _____________ describing the object.) *Examples: ___________, ___________, long, __________... Quantitative measurements give results in numeric form. (The results of a measurement contain a _____________.) *Examples: 4’6”, __________, 22 meters, __________... Accuracy vs. Precision Accuracy is how close a ___________ measurement is to the ________ __________ of whatever is being measured. Precision is how close ___________ measurements are to _________ ___________. adjective short heavy cold number 600 lbs. 5 ºC single true value several each other
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Practice Problem: Describe the shots for the targets.
Bad Accuracy & Bad Precision Good Accuracy & Bad Precision Bad Accuracy & Good Precision Good Accuracy & Good Precision
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Significant Figures Significant figures are used to determine the ______________ of a measurement. (It is a way of indicating how __________ a measurement is.) *Example: A scale may read a person’s weight as 135 lbs. Another scale may read the person’s weight as lbs. The ___________ scale is more precise. It also has ______ significant figures in the measurement. Whenever you are measuring a value, (such as the length of an object with a ruler), it must be recorded with the correct number of sig. figs. Record ______ the numbers of the measurement known for sure. Record one last digit for the measurement that is estimated. (This means that you will be ________________________________ __________ of the device and taking a __________ at what the next number is.) precision precise second more ALL reading in between the marks guess
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Significant Figures Practice Problems: What is the length recorded to the correct number of significant figures? length = ________cm 11.65 (cm) length = ________cm 58
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For Example Lets say you are finding the average mass of beans. You would count how many beans you had and then find the mass of the beans. 26 beans have a mass of grams. 44.56 grams ÷26 = grams So then what should your written answer be? How many decimal points did you have in your measurement? Rounded answer = 2 1.71 grams
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The SI System (The Metric System)
Here is a list of common units of measure used in science: Standard Metric Unit Quantity Measured kilogram, (gram) ______________ meter ______________ cubic meter, (liter) ______________ seconds ______________ Kelvin, (˚Celsius) _____________ The following are common approximations used to convert from our English system of units to the metric system: 1 m ≈ _________ kg ≈ _______ L ≈ 1.06 quarts 1.609 km ≈ 1 mile gram ≈ ______________________ 1mL ≈ _____________ volume 1mm ≈ thickness of a _______ mass length volume time temperature 1 yard 2.2 lbs. mass of a small paper clip sugar cube’s dime
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The SI System (The Metric System)
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kilo- hecto- deka- deci- centi- milli-
Metric Conversions The metric system prefixes are based on factors of _______. Here is a list of the common prefixes used in chemistry: kilo- hecto- deka deci- centi- milli- The box in the middle represents the standard unit of measure such as grams, liters, or meters. Moving from one prefix to another involves a factor of 10. *Example: 1000 millimeters = 100 ____ = 10 _____ = 1 _____ The prefixes are abbreviated as follows: k h da g, L, m d c m *Examples of measurements: 5 km 2 dL 27 dag 3 m 45 mm mass cm dm m grams Liters meters
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380 km = ______________m 1.45 mm = _________m
Metric Conversions To convert from one prefix to another, simply count how many places you move on the scale above, and that is the same # of places the decimal point will move in the same direction. Practice Problems: 380 km = ______________m mm = _________m 461 mL = ____________dL cg = ____________ dag 0.26 g =_____________ mg ,000 m = _______km Other Metric Equivalents 1 mL = 1 cm L = 1 dm3 For water only: 1 L = 1 dm3 = 1 kg of water or mL = 1 cm3 = 1 g of water (1) How many liters of water are there in 300 cm3 ? ___________ (2) How many kg of water are there in 500 dL? _____________ kilo- hecto- deka deci- centi- milli- 380,000 4.61 0.0004 260 230 0.3 L 50 kg
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Metric Volume: Cubic Meter (m3)
10 cm x 10 cm x 10 cm = Liter
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Ch. 4 Problem Solving in Chemistry
Dimensional Analysis Used in _______________ problems. *Example: How many seconds are there in 3 weeks? A method of keeping track of the_____________. Conversion Factor A ________ of units that are _________________ to one another. *Examples: 1 min/ ___ sec (or ___ sec/ 1 min) ___ days/ 1 week (or 1 week/ ___ days) 1000 m/ ___ km (or ___ km/ 1000 m) Conversion factors need to be set up so that when multiplied, the unit of the “Given” cancel out and you are left with the “Unknown” unit. In other words, the “Unknown” unit will go on _____ and the “Given” unit will go on the ___________ of the ratio. conversion units ratio equivalent 60 60 7 7 1 1 top bottom
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If your units did not ________ ______ correctly, you’ve messed up!
How to Use Dimensional Analysis to Solve Conversion Problems Step 1: Identify the “________”. This is typically the only number given in the problem. This is your starting point. Write it down! Then write “x _________”. This will be the first conversion factor ratio. Step 2: Identify the “____________”. This is what are you trying to figure out. Step 3: Identify the ____________ _________. Sometimes you will simply be given them in the problem ahead of time. Step 4: By using these conversion factors, begin planning a solution to convert from the given to the unknown. Step 5: When your conversion factors are set up, __________ all the numbers on top of your ratios, and ____________ by all the numbers on bottom. Given Unknown conversion factors multiply divide If your units did not ________ ______ correctly, you’ve messed up! cancel out
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How many hours are there in 3.25 days?
Practice Problems: How many hours are there in 3.25 days? (2) How many yards are there in 504 inches? (3) How many days are there in 26,748 seconds? 24 hrs 3.25 days 78 hrs x = 1 day 1 ft 1 yard 504 in. 14 yards x x = 12 in. 3 ft 1 min 1 hr 1 day 26,748 sec days x x x = 60 sec 60 min 24 hrs
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Converting Complex Units
A complex unit is a measurement with a unit in the _____________ and ______________. *Example: 55 miles/hour meters/sec g/mL To convert complex units, simply follow the same procedure as before by converting the units on ______ first. Then convert the units on __________ next. Practice Problems: (1) The speed of sound is about 330 meters/sec. What is the speed of sound in units of miles/hour? (1609 m = 1 mile) (2) The density of water is 1.0 g/mL. What is the density of water in units of lbs/gallon? (2.2 lbs = 1 kg) (3.78 L = 1 gal) numerator denominator top bottom 330m 1 mile 3600 sec 738 miles/hr x x = sec 1609 m 1 hr 1.0 g 1 kg 2.2 lbs 1000 mL 3.78 L 8.3 lbs/gal x x x x = mL 1000 g 1 kg 1 L 1 gal
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Ch. 6 Notes -- Chemical Composition
What is a mole?
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Ch 6 – Chemical Quantities
The Mole!!! A counting unit Similar to a dozen, except instead of 12, it’s 602 billion trillion… (602,000,000,000,000,000,000,000) ___________ (in scientific notation) This number is named in honor of Amedeo _________ (1776 – 1856), who studied quantities of gases and discovered that no matter what the gas was, there were the same number of molecules present…6.02 x 1023 6.02 x 10 23 Avogadro
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Just How Big is a Mole? Enough soft drink cans to cover the surface of the earth to a depth of over 200 miles. If you had Avogadro's number of un-popped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles. If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole.
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The Mole 1 dozen cookies = ___ cookies
12 1 dozen cookies = ___ cookies 1 mole of cookies = ___________ cookies 1 dozen cars = ___ cars 1 mole of cars = __________ cars 1 dozen Al atoms = ___ Al atoms 1 mole of Al atoms = __________ atoms Note that the NUMBER is always the same, but the ______ is very different! Mole is abbreviated ______ . 6.02 X 1023 12 6.02 X 1023 12 6.02 X 1023 MASS mol
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The Mole and Mass Mass in grams of 1 mole equal to __________ of the atomic masses. Practice problem: Calculate the mass of 1 mole of CaCl2. Ca = 1 x ________ g/mol = Cl = 2 x ________ g/mol = = __________ g/mol CaCl2 1 mole of CaCl2 = the sum 40.1 40.1 g/mol 35.5 71.0 g/mol 40.1 g/mol 71.0 g/mol 111.1 111.1 g/mol
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Mole Conversion Factors that you will need to know!
1 mol = __________ atoms/molecules/etc. 1 mol = _____________ grams 1 mol = ________ Liters of gas at STP (STP is Standard Temp. and Pressure… we will talk about what this means later!) 6.02 x 1023 ? ( molar mass) 22.4
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Ch. 6 Notes -- Chemical Composition
Practice Problems: (1) How many atoms of hydrogen are there in each compound? a) Ca(OH)2 ___ b) C3H8O___ c) (NH4)2HPO4 ___ d) HC2H3O2 ___ (2) Calculate the formula mass of each compound. (Add up all the atomic masses for each atom from the Periodic Table.) a) CaCO3 b) (NH4)2SO4 c) C3H6O d) Br2 e) H3PO4 f) N2O5 2 8 9 4 2 N’s = 2 x 14.0 = 28.0 8 H’s = 8 x 1.0 = 8.0 S = 32.1 4 O’s = 4 x 16.0 = 64.0 Ca = 40.1 C = 12.0 3 O’s =3 x 16.0 = 48.0 Add them up! 132.1 g/mol Add them up! 100.1 g/mol 159.8 g/mol C = 3 x 12.0 = 36.0 H = 6 x 1.0 = 6.0 O =16.0 2 Br’s = 2 x 79.9 = Add them up! 58.0 g/mol 2 N’s = 2 x 14.0 = 28.0 5 O’s = 5 x 16.0 = 80.0 3 H’s = 3 x 1.0 = 3.0 P = 31.0 4 O’s = 4 x 16.0 =64.0 Add them up! 98.0 g/mol Add them up! 108.0 g/mol
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3) Convert 835 grams of SO3 to moles.
4) How many molecules of CH4 are there in 18 moles? 5) How many grams of helium are there in 5.6 x 1023 atoms of helium? 6) How many molecules are there in 3.7 grams of H2O? 1 mole SO3 835 g SO3 10.4 moles of SO3 x = 80.1 g SO3 6.02 x 1023 molecules CH4 18 moles CH4 x = 1.08 x 1025 molecules CH4 1 mole CH4 4.0 grams He 5.6 x 1023 atoms He x 3.72 grams He = 6.02 x 1023 atoms He 6.02 x 1023 molecules H2O 3.7 grams H2O 1.24 x 1023 molecules H2O x = 18.0 grams H2O
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Calculating Percent Composition by Mass
Step 1: Find the formula mass of the compound by adding the individual masses of the elements together. Step 2: Divide each of the individual masses of the elements by the formula mass of the compound. Step 3: Convert the decimal to a % by multiplying by 100. Practice Problems: (1) Find the % composition of the elements in each compound. a) Na3PO b) SnCl4 = = 42.1% 3 Na’s = 3 x 23.0 = 69.0 P = 31.0 4 O’s = 4 x 16.0 = ÷ 164 Sn = 118.7 4 Cl’s = 4 x 35.5 = ÷ 260.7 = 45.5% = = 18.9% ÷ 164 + ÷ 260.7 = 54.5% 260.7 = = 39.0% + ÷ 164 164
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Elements in the Universe: % Composition by Mass
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Earth’s Crust: % Composition by Mass
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Entire Earth (Including Atmosphere): % Composition by Mass
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Human Body: % Composition by Mass
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Determining the Empirical Formula for a Compound
The empirical formula for a compound is the simplest __________ number __________ of the atoms in the compound. Examples: H2O is the empirical formula for water. _______ is the empirical formula for glucose, C6H12O6. Practice Problems: What is the empirical formula for the following compounds? a) C6H6= ________ b) C8H14O2 = ________ c) C10H14O2 = _________ d) Ca5Br10 = ________ e) N3O9 = ________ whole ratio C1H2O1 CH C4H7O C5H7O CaBr2 NO3
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Determining the Molecular Formula for a Compound
The molecular formula for a compound is either the same as the empirical formula ratio or it is a “_________ _________ of this ratio. It represents the true # of atoms in the molecule. Examples: 1) H2O is the empirical & molecular formula for water ) CH2O is the empirical formula for sugar, ethanoic acid, and methanol. The molecular formula for glucose is C6H12O6, (___times the empirical ratio!) Practice Problems: (1) If the empirical formula for a compound is CH2, which of the following is a possible molecular formula for the compound? a) C8H16 b) C8H8 c) C4H2 d) C3H9 (2) If the empirical formula for a compound is C2H3, which of the following is a possible molecular formula for the compound? a) C2H6 b) C10H c) C6H12 d) C8H14 whole # multiple 6
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Determining the Molecular Formula for a Compound
Find the molecular formula for C2H7 if the molecular mass of the compound is 93.0 g/mol. Find the molecular formula for P2O5 if the molecular mass of the compound is g/mol. C2H7 = 31.0 g/mol 93.0 g/mol = 3 31.0 g/mol C2H7 x 3 = C6H21 g/mol P2O5 = g/mol = 2 g/mol P2O5 x 2 = P4O10
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