Matter & Measurement. Chapter 1 Chemistry 2A

Transcription

1 Matter & Measurement Chapter 1 Chemistry 2A

2 Chemistry: the branch of science concerned with the characteristics, composition, and transformations of matter Matter: anything that has mass and occupies space Living and non-living Macroscopic and microscopic

3 States of Matter Three physical states

4 Solids Definite shape Definite volume Atoms packed tightly together May be crystalline or amorphous Very low compressibility

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6 Liquids No definite shape Definite volume Atoms close together, but not locked into place Low compressibility

7 Gases No definite shape No definite volume Atoms/Molecules far apart Compressible

8 Classification of Matter Pure Substances or mixture Pure Substances: single component which cannot be broken down Mixture : composed of more than one substance, can be separated into its components

9 Pure Substances Element or compound 1) Element: a substance that cannot be decomposed or transformed into other chemical substances by ordinary chemical processes Atom: smallest particle of an element that can exist and still have the properties of that element Aluminum (Al), Carbon (C), Neon (Ne), Potassium (K)

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12 2) Compound: A substance consisting of two or more different atoms chemically bonded in a fixed ratio NaCl, CO2 Ionic compounds and molecules

13 Mixtures Mixtures: combinations of two or more pure substances in which each substance retains its own identity 1) Heterogeneous Mixture: a substance in which elements and/or compounds are blended together in such a way that there is no uniform composition or fixed ratio of the components of the mixture Examples: Oil and water, mixed nuts

14 2) Homogeneous Mixture: A substance in which the different elements/compounds being mixed exist in definite ratios, but are not chemically bonded Consists of two or more substances in the same phase No amount of magnification will reveal an interface Called a solution Examples: Salt water, sugar water, O2 dissolved in water

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16 Problems Decide whether the following mixtures are heterogeneous or homogenous 1. Chocolate chip cookie dough 2. Wine 3. Milk 4. O2 in water 5. Chicken noodle soup

17 Physical Properties Properties of an object or substance that can be measured or perceived without changing the identity of the substance

18 Classification of Physical Properties Extensive Properties: properties that depend on the amount of substance present Mass Volume Intensive Properties: properties that do not depend on the amount of substance present M.P. B.P.

19 Physical Change A change in a physical property of a substance Same substance before and after the change

20 Changes in State

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22 Chemical Property Any property of a material that becomes evident during a chemical reaction Qualities that become evident by changing a substance s identity Capability to undergo chemical reactions Flammability Acidity Corrosiveness Toxicity

23 Chemical Change A process in which reactants are changed into one or more different products Have breaking and making of chemical bonds chemical reaction

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26 Problems Decide whether the following are chemical or physical changes 1) Sawing a log in half 2) Melting chocolate in a pot on your stove 3) Burning your chocolate 4) Dissolving a nickel in acid 5) Cutting your hair

27 Measurements in Chemistry

28 Data Qualitative Data obtained from one s opinion Does not involve numbers Quantitative Data obtained from measurements Involves numbers

29 U.S. Customary System Also called: American System English System Inch Gallon Pound Teaspoon

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31 Metric System Système International (SI) International decimalized system of measurement First adopted by France in 1791 Meter Gram Liter

32 Length How long something is, SI unit = meter (m)

33 Mass Measure of the quantity of matter (stuff) in an object SI unit = Kilogram (kg)

34 Volume The amount of space that an object or substance occupies. SI unit = Cubic meter (m3) 1 L = m3 1 L = 1000 ml 1 ml = 1 cm3 = 1 cc

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37 Time Duration of event SI unit = Second (s)

38 French Revolutionary Clock

39 System International (SI) Units

40 Prefix Giga Mega Kilo Hecto Deka No prefix (Unity) Deci Centi Milli Micro Nano Pico Symbol G M k h da d c m μ n p Multiple = = = = = = Example Gigabyte = Gbyte Kilogram = kg Meter, liter, gram = m, L, g Milliliter = ml Nanometer = nm

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43 Common Units and Their Equivalents Length 1 kilometer (km) 1 meter (m) 1 meter (m) 1 foot (ft) 1 inch (in.) = = = = = mile (mi) inches (in.) yards (yd) centimeters (cm) 2.54 centimeters (cm) exactly 1 kilometer (km) = 1000meter 1 meter = 100 centimeter 1 meter = 1000 millimeter

44 Common Units and Their Equivalents Mass 1 kilogram (km) = pounds (lb) 1 pound (lb) = grams (g) 1 ounce (oz) = grams (g) Volume 1 liter (L) 1 liter (L) 1 liter (L) 1 U.S. gallon (gal) = = = = 1000 milliliters (ml) 1000 cubic centimeters (cm3) quarts (qt) liters (L)

45 1Kg ( kilogram) = 1000g ( gram) 1g (gram) = 1000mg (milligram)

46 Problems 1) Green light has a wavelength of approximately 550 nm. What is this value in meters? Picometers? Kilometers?

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48 2) Your neighbor lost 50 pounds. How many kg did she lose? How many micrograms? 3) How many milliseconds in a year?

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51 Dimensional Analysis Using units as a guide to problem solving is called dimensional analysis Figure out which unit you want to start with and which one you want to get to Use conversion factors to get there Relationship between two units May be exact or measured Generated from equivalence statements Always include units in your calculations!

52 Conversion Factors To convert one unit to another we use one or more conversion factor original quantity X conversion factor = desired quantity

53 12 eggs = 1 dozen desired unit given unit = desired unit given unit

54 Choose correct conversion factor How many grams are there in 5Kg Convert 20 Km to meter How many dozen have 36 eggs How many grams of aspirin are contained in a 325-mg tablet

55 Temperature A measure of the average kinetic energy of the particles in a sample of matter, expressed in terms of units or degrees designated on a standard scale. A physical property that determines the direction of heat flow in an object upon contact with another object. Fahrenheit ( F), Celsius ( C), Kelvin (K)

56 Fahrenheit (ºF), Celsius (ºC), Kelvin (K) ºF = ºC(1.8) + 32 ºC = (ºF 32)/1.8 K = ºC ºC = K 273

57 Lord William Thomas Kelvin

58 Problems 1) If it s 35ºC in London, would you say that it s probably winter or summer? What is this temperature in Kelvin? 2) You are feeling sick and decide to take your temperature. Your thermometer, which only reads temps in Kelvin, says that you are at approximately 312 K. Do you have a fever?

59 Density The ratio of the mass of an object to its volume Mercury Water 13.6 g/cm3 1.0 g/ml 8.94 g/cc

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61 Problems 1) Calculate the density of the rock in the picture to the right. The rock has a mass of 29.5 g.

62 2) What is the mass of 5.5mL of mercury if Hg has a density of g/ml? 3) Calculate the width of the piece of wood to the right. Oregon Pine d = 0.53 g/ml

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66 Scientific Notation 1) Locate the decimal point 2) Move the decimal so that there is only one number to the left of it 3) Write x 10 behind you new number 4) Count the number of places you ve moved your decimal point and make this number the exponent on your 10 5) Assign a + or sign to your exponent a) b) If your original # is larger than your SN #, the exponent is + If your original # is smaller than your SN #, the exponent is

67 Problems Write the following standard numbers in scientific notation and write the numbers in scientific notation in standard form. 1) 252 2) ) ) ) 3.33 x 102

68 6) 4 x ) ) 8 9) 80

69 Significant Figures

70 Scientific measurements are reported so that every digit is certain except the last, which is estimated Certain Uncertain

71 Rules for Significant Figures Numbers up to and including the uncertain number are significant 2) All non-zero numbers are significant 3) Zeros may or may not be significant 4) Zeros are significant if a) They are between two non-zero digits b) They are at the end of a decimal number 1)

72 5) Zeros are not significant if a) They are used as place holders in large numbers without a decimal point b) They are at the beginning of decimal numbers 6) All numbers displayed in a number written in scientific notation are significant

73 Problems Identify the correct number of significant digits in the figures below. 1) 2) 3) 4) 5) ) ) ) ) 4.0 x ) 3 x 108

74 Mt. Everest ft, x 104 ft., or ft?

75 Calculation With Significant Digits Multiplication and Division The final answer has the same number of sig figs as the measurement with the fewest sig figs Example 1: 22.2 cm x cm =? Example 2: mm / mm =?

76 Addition and Subtraction The final answer is written so that it has the same number of decimal places as the measurement having the fewest decimal places Example 1: 44.4 L L L =? Example 2: 4107 in in =?

77 Problems 1) 2) 3) 4) / 4.4 x

78 Precision and Accuracy Precision: how well several determinations of the same measurement agree Reproducibility/repeatability Accuracy: agreement of a measurement with the accepted value

79 Determine whether the following students exhibit good or poor accuracy and precision Exam 1 Exam 2 Student A 99% 100% Student B 100% 89% Student C 59% 59% Student D 25% 49% Accuracy & Precision

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