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*STUDENT* *STUDENT* Regents Chemistry NOTE PACKET Unit 1: Matter & Measurement Copyright © 2015 Tim Dolgos Copyright © 2015 Tim Dolgos Copyright © 2015 Tim Dolgos *STUDENT* *STUDENT* Unit Vocabulary: 1. 2. 3. 4. 5. 6. 7. 8. 9. S.I. unit Meter Liter Gram Mass Weight Volume Density Intensive 10. 11. 12. 13. 14. 15. 16. 17. 18. Extensive Significant Figures Precision Accuracy Matter Element Compound Mixture Heterogeneous Mixture 19. 20. 21. 22. 23. 24. 25. Homogeneous Mixture Pure Substance Particle Diagram Chromatography Filtration Distillation Scientific Notation Unit Objectives: When you complete this unit you will be able to do the following… 1) Classify types of matter 2) Draw particle diagrams to represent different types of matter 3) Recognize various techniques that can be used to separate matter 4) Convert between units of measurements 5) Differentiate between accuracy and precision 6) Write numbers in scientific notation 7) State rules to determine significant figures 8) Count significant figures 9) Understand the importance of significant figures 10) Calculate the volume and density of an object Copyright © 2015 Tim Dolgos Matter Can NOT be separated by physical means Can NOT be separated by chemical means Particle Diagram CAN be Separated by PHYSICAL means Same composition throughout Separated by chemical means, only Particle Diagram Particle Diagram 4 Different composition throughout Particle Diagram Copyright © 2015 Tim Dolgos Practice Problems: 1. Which particle diagram(s) represent a mixture? 2. Which particle diagram(s) represent a pure substance? 3. Which of the following particle diagrams represents a mixture of one compound and one element? 4. Which particle diagram represents a diatomic element? 5 Copyright © 2015 Tim Dolgos Properties of Matter: Physical properties are the constants about a substance; can use our senses to observe them; do not require chemical analysis Example: o Extensive Property: a property that depends on how much material you are dealing with Ex: o Intensive Property: a property that does not depend on how much material you are dealing with (help identify matter; a constant about that particular type of matter) Ex: Chemical properties include behaviors substances adhere to when they __________ with other substances Examples: Guided Practice: Identify the following as being intensive, extensive, or chemical properties. ____________ 1. The mass of copper wire is 255 g. ____________ 2. The boiling point of ethyl alcohol is 77°C. ___________ 3. Baking soda reacts with vinegar to make carbon dioxide gas. ____________ 4. The density of mercury is 13.6g/mL. ____________ 5. The solubility of sodium chloride in water is 40g/100mL of water. 6 Copyright © 2015 Tim Dolgos Physical vs. Chemical Changes Matter is always changing. Ice in your drink melts. Wood in your fire burns. Physical Change – a change that does NOT alter the chemical properties of a substance (example: ______________, ______________); change in size or shape; ____________________________ *PHYSICAL processes can be reversed ________________________ Example: ice melting to become liquid (its still water!) Chemical Change – a reaction in which the composition of a substance is changed (example: __________); properties _______________________ 1. Signs of a chemical rxn: 1. 2. 3. Example: firewood burning Change of Matter Physical or Chemical? Burning toast Making ice cubes Lighting a candle Spoiling milk Making kool-aid 7 Copyright © 2015 Tim Dolgos Elements vs. Compounds Element = _____________________________________________________ Compound = ____________________________________________________ 1. Circle ( ) all the elements and underline the compounds below. 2. On the line provided, record the number of different symbols within the species. CO ___ C2H5OH Mg ___ H2SO4 ___ O2 ___ C ___ ___ Al(CN)3 ___ He ___ ___ Cl2 ___ NI3 H2O ___ Cu Co ___ ___ NaCl ___ I ___ Questions: 1) Does each compound have the same number of symbols? ____ 2) For each ELEMENT above, how many total symbols are listed? __ 3) What is the minimum number of symbols that must be present in order for a species to be considered a compound? __ Understanding Compound Formulas: Within a compound, you may see subscripts. These subscripts tell you the number of each type of atom that is present. Example: # carbon atoms __ # oxygen atoms __ If there are parentheses present around two or more atoms, the subscript applies to all atoms within the parentheses. Example: # aluminum atoms __ # carbon atoms __ # nitrogen atoms __ If one of the atoms within the parentheses has a subscript, you multiply this number by the number outside of the parentheses. Example: # iron atoms __ # sulfur atoms __ 8 # oxygen atoms ___ Copyright © 2015 Tim Dolgos There is no vodcast for this page. Just know that you will be responsible for memorizing the symbol and name for the most commonly used elements, which appear below. You will be quizzed on each of the two sets below at some point next week. I recommend making flash cards for all of them (there’s a flashcard app, and my former students loved it) and spend a couple minutes a night going over them. The Common Elements Rules for writing element symbols: 1) ______________________________________ 2) ______________________________________ * Symbol * Ag Al Ar As Au B Ba Be Br C Ca Cl Co Cr Cs Cu F Fe Fr H He Hg * Name * silver aluminum argon arsenic gold boron barium beryllium bromine carbon calcium chlorine cobalt chromium cesium copper fluorine iron francium hydrogen helium mercury * Symbol * I K Kr Li Mg Mn N Na Ne Ni O P Pb Ra Rb Rn S Si Sn Sr U Xe Zn * Name * iodine potassium krypton lithium magnesium manganese nitrogen sodium neon nickel oxygen phosphorus lead radium rubidium radon sulfur silicon tin strontium uranium xenon zinc MEMORIZE both directions (symbol to name, name to symbol) for Quiz on _____________ 9 Copyright © 2015 Tim Dolgos Separation of Matter Separation Apparatus Type of Separation (Physical or Chemical) Description of Technique What types of matter will it separate? Filtration Watch Glass Evaporation Crucible Evaporation 10 Copyright © 2015 Tim Dolgos Separation of Matter (continued) Distillation Chromatography On the other hand _____________________ requires reacting a sample with something else in order to turn it into a completely different compound 11 Copyright © 2015 Tim Dolgos SCIENTIFIC NOTATION – method for expressing very large or small numbers easily (Example: ___________________) For example, the number 1,000,000 is in standard formation format. The scientific notation of this number is 1.0 x 106 We always move the decimal place to make the ____________(the number out in front) between _______________ We then arrange the ___________ (the number up to the right of the ten) Now, _______________ if you were to take the 1.0 and move the decimal place 6 places to the right (since it is a positive number), you would get the original number (1,000,000) Example: 123000000000 Guided Practice – Write the following numbers in scientific notation (remember the mantissa rule!) 1. 34000000 = 2. 0.0000067 = 3. 25,864 = Now, write the following scientific notations in standard (normal) notation form: 4. 5.7 x 108 = 5. 6.34 x 10-11 = Calculator Practice: First, let’s enter the number 2.3 x 10-5 in scientific notation: 1. Type “2” 2. Type the decimal point 3. Type “3” 4. Then press the “ee” “EXP” or “ ” key(s) 5. Press the “+/-“ key (NOT the “—“ or “subtract” key) 6. Type “5” Next, let’s enter that number by 1 mole, or 6.02 x 1023. What do you get for your answer? _________________ 12 Copyright © 2015 Tim Dolgos Measurements and the Metric System In chemistry we measure matter using ____ units. This is an abbreviation for _________________________________. SI BASE UNITS (AKA Base Units): **If you forget, use Table D in your Reference Tables! 13 Copyright © 2015 Tim Dolgos SI Metric Prefixes Prefix Symbol tera T giga G mega M kilo k hecto h deca da no prefix: deci d centi c milli m micro nano n pico p femto f atto a Numerical (Multiply Root Word by)* 1,000,000,000,000 1,000,000,000 1,000,000 1,000 100 10 1 0.1 0.01 0.001 0.000001 0.000000001 0.000000000001 0.000000000000001 0.000000000000000001 Exponential 1012 109 106 103 102 101 100 10¯1 10¯2 10¯3 10¯6 10¯9 10¯12 10¯15 10¯18 *Example: In the word kilometer, the root word (base unit) is “meter” and the prefix is “kilo.” Kilo means multiply the root word by 1000. Therefore, one kilometer is 1000 meters (1 km = 1000 m). 14 Copyright © 2015 Tim Dolgos Conversion Factors – a mathematical expression that relates two units that measure the same type of quantity Examples: - *Rest Assured! For the Regents, the most you will have to convert will be between the milli-/kilo-/base unit (g, L, etc.). This is always a matter of ___________________. You must also make sure you move the decimal the ___________________ (right or left, which depends on whether you are converting from small to big or vice versa). TRICK: kilo hecto deca base unit deci centi milli k h d base unit d c m Let’s practice! 1. A car travels 845 km. How many meters is this? 2. Convert 0.0290 L to milliliters. 3. Convert 2500mL to liters. 9. 12 mL = ______ L 4. 3 g = _______ kg 5. 1 km = ______ m Compare by placing a <, >, or = on the line provided: 6. 1 kg = _______ g 10. 56 cm 11. 7g 7. 1 L = ________ mL 8. 7 m = _______ mm __ 6 m __ 698 mg Once you get your answer, check it! Does it make sense? 15 Copyright © 2015 Tim Dolgos Dimensional Analysis Often you will be required to solve a problem with mixed units, or to convert from one set of units to another. Dimensional analysis is a simple method to accomplish this task. Ex: How many minutes are there in 15 days? Solution A: STEP 1: Figure out the units that you have and the steps to get to the units that you need. HAVE (What’s missing?) NEED Days (d) Minutes (min) STEP 2: Make a “grid” and plug in the numbers to make your first conversion. The number/units you HAVE goes in the top left, the number/units you NEED go in the top right, and the conversion factor goes in the bottom right. Need Have 15 d 24 h 1d = Conversion Factor STEP 3: Cancel “like terms.” Then, multiply the top numbers (the numerators) together and divide the result by the bottom number (the denominator). 15 d 24 h 1d = 360 h Since 24 hours and 1 day are equivalent, you are actually multiplying 15 days by a factor of 1. This means that the magnitude of your number stays the same and only the units change. In other words, 15 days = 360 hours STEP 4: Now, use your answer from Step 3 as the new “HAVE” and repeat the process using the conversion factor 60 minutes = 1 hour 360 h 60 min 1h = 21,600 min Now you try on: How many minutes are there in the month of October? 16 Copyright © 2015 Tim Dolgos ACCURACY VS. PRECISION Accuracy – ____________________________________________ Ex: Hitting bulls eye when you are aiming for it *For most experiments, ______________ means _____ from the expected value Precision – _____________________________________________ _____________________________________________________ *For an experiment with +/- 5% as the margin for accuracy, that means the difference between the highest and lowest percent error cannot exceed a ____________________ Ex: If the highest percent error for an experiment is +7.6%, and the lowest is -5.4% that range is 13.0%, which means that experiment was not precise Practice: Cheryl, Cynthia, Carmen, and Casey take target practice in PE. Assuming that they were all aiming at the bulls eye, match each target with the proper description. (a) Accurate and precise (b) Not precise, but one piece of data accurate (c) Precise but not accurate (d) Neither precise nor accurate Practice: The following data was collected during a lab experiment. The density of the cube should be 10.8 g/mL. Is this data is accurate, precise, both, or neither? Justify your answer. ______________________________________________ ______________________________________________________________ Trial Number 1 2 3 Density of Cube 6.2 g/mL 6.3 g/mL 6.5 g/mL 17 Copyright © 2015 Tim Dolgos SIGNIFICANT FIGURES - also known as Sig Figs (SF) A method for handling ____________________ in all measurements This arises due to the fact that we have different equipment with different degrees of ___________________ Significant figures are associated with ____________________________ _______________________ do _______________ when determining sig figs o Ex: Atomic masses on periodic table Conversions (1in = 2.54 cm) Examples: 1. Reading a ruler We know for sure that the object is more than _____, but less than _____ We know for sure that the object is more than _____, but less than _____ This ruler allows us to estimate the length to ________ 2. Reading a graduated cylinder: ► Measurements are read from the bottom of the _________ ►Which gives a volume reading of _______ 18 Copyright © 2015 Tim Dolgos The Atlantic/Pacific Method - another way to determine the # sig figs in a number 1) 2) Determine if a decimal point is present. If a decimal is present, think of “P” for “present.” If there is no decimal, think of “A” for “Absent.” P stands for the Pacific coast and A stands for the Atlantic Coast. Imagine the number you are looking at is a map of the USA. Begin counting from the correct side of the number (Atlantic/right side or Pacific/left side) based on what you determined in step 1. Consider the first nonzero number you land on the start of your count. Consider each digit from here on out significant as well until you reach the other end of the number. Pacific Coast 3. Atlantic Coast Decimal is Present Decimal is Absent 1. Start @ 1st NONZERO 1. Start @ 1st NONZERO 2. Count all the way to the Atlantic—NO EXCEPTIONS 2. Count all the way to the Pacific—NO EXCEPTIONS Determine the number of significant numbers in each of the following: 1) 357 _______ 5) 0.0357 2) 3570 _______ 6) 3.570 x 103 _______ 3) 3570. _______ 7) 0.3570 _______ 4) 0.357 _______ 19 _______ Copyright © 2015 Tim Dolgos Rules for Determining Number of Significant Figures in a Given Number Rule Example 1. All nonzero numbers (ex: 1 – 9) are always 123456789 m significant 1.23 x 102 2. Zeros located between nonzero numbers are significant 40.7 L 87,009 km 3. For numbers less than one, all zeros to the left of the 1st nonzero number are NOT significant 0.009587 m 4. Zeros at the end of a number and to the right of a decimal point are significant 85.00 g 5. Zeros at the end of a whole number may be significant or not. If there is a decimal after the last zero, they are significant. If there is not a decimal point after the end zeros, they are NOT significant 6. Exact numbers have an infinite number of significant figures 2000 m PRACTICE: Measurement 1020 mL 1200 m 1200. L 1200.00 mm 0.001 km 10.00 L 12000 m 00.100 cL 22.101 mm 101,000 km 0.0009 kg 9.070000000 L 2000. m 1 ft = 12 inch Number of Significant Figures 20 Rule(s) Applied Copyright © 2015 Tim Dolgos Rules for Using Sig Figs in Calculations General Rule Final answer must be expressed in the lowest amount of significant figures that were originally given to you (you can’t create accuracy when you didn’t have it to start with!) Operation Rule Multiplication/Division Perform operation as normal & express answer in least # sig figs that were given to you Addition/Subtraction Examples: Examples Line decimal points up; round final answer to lowest decimal place (least accurate) value given 12.257 x 1.162 = + 3.95 2.879 213.6____ 5.1456 – 2.31 = _______ 69.25/45.8 = _________ Rules for Calculations with Numbers in Scientific Notation: Rule Example Addition/Subtraction All values must 4.5 x 106 - 2.3 x 105 have the same exponent. Result is the sum or difference of the mantissas, multiplied by the same exponent of 10 Multiplication mantissas are multiplied and exponents of 10 are (3.1 x 103) (5.01 x 104) added Division mantissas are divided and exponents are subtracted 7.63 x 103 / 8.6203 x 104 21 Copyright © 2015 Tim Dolgos MEASURING MATTER 1. Mass vs. Weight MASS WEIGHT *We really only work with ________ in chemistry class! ** We have the same _________ whether we are on earth or on the moon. The different forces of gravity on each cause us to weigh more on earth than on the moon though (this is why we float on the moon!) 2. Volume - amount of _____________ an object takes up Techniques: Liquids Regular Solids Irregular Solids 3. Density: amount of mass in a given space; _________ of mass to volume Formula (Table T): BOX A BOX B Which box has a higher density? Explain your answer. ____________________________________________________ ____________________________________________________ 22 Copyright © 2015 Tim Dolgos Density Problems – Show all work! *Note: the density of water is ______________ 1) What is the density of an object with a mass of 60 g and a volume of 2 cm3? 2) If you have a gold brick that is 2.0 cm x 3.0 cm x 4.0 cm and has a mass of 48.0 g, what is its density? 3) If a block of wood has a density of 0.6 g/ cm3 and a mass of 120 g, what is its volume? 4) What is the mass of an object that has a volume of 34 cm3 and a density of 6.0 g/cm3? 5) Which is heavier, a ton of feathers or a ton of bowling balls? 23 Copyright © 2015 Tim Dolgos Percent Error Measurement of the % that the measured value is “off” from accepted value Measured value = Accepted value = Formula is found in Table T (back page 12) of your Reference Tables: If negative, your measured value is ________________ the accepted value If positive, your measured value is ________________ the accepted value *It is very important that you put the given values into the proper place in the formula! Sample Problem: In a lab experiment, you are told by your teacher that the actual (or accepted) amount of sugar in a can of Coke is 39 g. You experimentally determine it to be 40 g based on your own data and calculations. What is your percent error? Express answer in the proper amount of significant figures. 24 Copyright © 2015 Tim Dolgos