Mr. Dolgos Regents Chemistry NOTE PACKET. Unit 1: Matter & Measurement

Transcription

1 Mr. Dolgos Regents Chemistry NOTE PACKET Unit 1: Matter & Measurement 1

2 Unit Vocabulary: 1. S.I. unit 2. Meter 3. Liter 4. Gram 5. Mass 6. Weight 7. Volume 8. Density 9. Intensive 10. Extensive 11. Significant Figures 12. Precision 13. Accuracy 14. Matter 15. Element 16. Compound 17. Mixture 18. Heterogeneous Mixture 19. Homogeneous Mixture 20. Pure Substance 21. Particle Diagram 22. Chromatography 23. Filtration 24. Distillation 25. 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 2

3 Date: Matter and Measurement Warm-up Questions Question Answer Date: Date: Date: Date: 3

4 Matter Can NOT be separated by physical means CAN be Separated by PHYSICAL means Can NOT be separated by chemical means Separated by chemical means, only Same composition throughout Different composition throughout Particle Diagram Particle Diagram Particle Diagram Particle Diagram 4

5 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

6 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

7 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: Example: firewood burning Change of Matter Burning toast Making ice cubes Lighting a candle Spoiling milk Making kool-aid Physical or Chemical? 7

8 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 Mg Co C 2 H 5 OH Al(CN) 3 Cl 2 H 2 SO 4 He NI 3 O 2 H 2 O NaCl C Cu 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 # oxygen atoms 8

9 * The Common Elements * Rules for writing element symbols: 1) 2) * Symbol * * Name * * Symbol * * Name * Ag silver I iodine Al aluminum K potassium Ar argon Kr krypton As arsenic Li lithium Au gold Mg magnesium B boron Mn manganese Ba barium N nitrogen Be beryllium Na sodium Br bromine Ne neon C carbon Ni nickel Ca calcium O oxygen Cl chlorine P phosphorus Co cobalt Pb lead Cr chromium Ra radium Cs cesium Rb rubidium Cu copper Rn radon F fluorine S sulfur Fe iron Si silicon Fr francium Sn tin H hydrogen Sr strontium He helium U uranium Hg mercury Xe xenon Zn zinc MEMORIZE both directions (symbol to name, name to symbol) for Quiz on 9

10 Separation of Matter Filtration Separation Apparatus Type of Separation (Physical or Chemical) Description of Technique What types of matter will it separate? Watch Glass Evaporation Crucible Evaporation 10

11 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

12 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 10 6 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: Guided Practice Write the following numbers in scientific notation (remember the mantissa rule!) = = 3. 25,864 = Now, write the following scientific notations in standard (normal) notation form: x 10 8 = x = 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 What do you get for your answer? 12

13 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

14 SI Metric Prefixes Numerical (Multiply Root Word Prefix Symbol by)* Exponential tera T 1,000,000,000, giga G 1,000,000, mega M 1,000, kilo k 1, hecto h deca da no prefix: deci d centi c milli m micro nano n pico p femto f atto a *Example: In the word kilometer, the root word (base unit) is meter and the prefix is kilo. Kilo means multiply the root word by Therefore, one kilometer is 1000 meters (1 km = 1000 m). 14

15 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 L to milliliters. 3. Convert 2500mL to liters g = kg 5. 1 km = m 6. 1 kg = g 7. 1 L = ml 8. 7 m = mm ml = L Compare by placing a <, >, or = on the line provided: cm 6 m g 698 mg Once you get your answer, check it! Does it make sense? 15

16 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 Density of Cube g/ml g/ml g/ml 17

17 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

18 Rules for Determining Number of Significant Figures in a Given Number Rule 1. All nonzero numbers (ex: 1 9) are always significant Example m 1.23 x 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 1 st nonzero number are NOT significant m kg 4. Zeros at the end of a number and to the right of a decimal point are significant g L 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 m 1 ft = 12 inch PRACTICE: Measurement Number of Significant Figures Rule(s) Applied 1020 ml 1200 m L mm km L m cl mm 101,000 km 20

19 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 Examples Multiplication/Division Perform operation as normal & express answer in least # sig figs that were given to you x = Addition/Subtraction Line decimal points up; round final answer to lowest decimal place (least accurate) value given Examples: = 69.25/45.8 = Rules for Calculations with Numbers in Scientific Notation: Rule Example Addition/Subtraction All values must 4.5 x x 10 5 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 10 3 ) (5.01 x 10 4 ) added Division mantissas are divided and exponents are subtracted 7.63 x 10 3 / x

20 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

21 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 cm 3? 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/ cm 3 and a mass of 120 g, what is its volume? 4) What is the mass of an object that has a volume of 34 cm 3 and a density of 6.0 g/cm 3? 5) Which is heavier, a ton of feathers or a ton of bowling balls? 23

22 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