1 Chapter 1: Matter and Measurements
2 What is Chemistry? Think, pair, share:What are all the chemicals you use in your daily life?
3 What is Chemistry? Biology vs. Chemistry vs. Physics
4 Biology Physics ChemistryWhat is Chemistry? Biology Physics Chemistry The study of living organisms The study of forces & motion The study of matter and its reactions and properties
5 What is Chemistry? Chemistry is the study of CHANGES in the “stuff” around us. (We formally define “stuff” as matter!)
6 Matter
7 Matter
8 Elements Type of matter that cannot be broken down into simpler, stable substances and is made of only one type of atom
9 Compounds A pure substance that contains two or more elements whose atoms are chemically bonded
10 Compounds Fixed compositionsA given compound contains the same elements in the same percent by mass
11 Compounds The properties of a compound are VERY DIFFERENT from the properties of the elements they contain Ex.) Sodium Chloride (NaCl) vs. Sodium & Chlorine Sodium: https://www.youtube.com/watch?v=bze3hN9j9Cw
12 Electrolysis
13 Mixtures A blend of two or more kinds of matter, each of which retains its own identity and properties Homogeneous Heterogeneous
14 Homogeneous Mixtures Composition is the same throughout the mixtureExamples: salt water, soda water, brass A.k.a. a solution Solute in a solvent (salt dissolved in water)
15 Heterogeneous MixturesNon-uniform; composition varies throughout the mixture
16 Separating Mixtures Filtration
17 Separating Mixtures Distillation
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19 Separating Mixtures Chromatography
20 Scientific MeasurementsChemistry is a quantitative science. This means that experiments and calculations almost always involve measured values. Scientific measurements are expressed in the SI (metric) system. This is a decimal-based system in which all of the units of a particular quantity are related to each other by factors of ten.
21 Metric System The metric system is based on a base unit that corresponds to a certain kind of measurement Length = meter (m) Volume = liter (L) Weight (Mass) = gram (g) Prefixes plus base units make up the metric system Example: Centi + meter = Centimeter Kilo + liter = Kiloliter
22 Prefixes (see handout & reference book)You will need to memorize all of the prefixes (factors, names and abbreviations from 109 (giga-) to (nano)! One example of a memory device:
23 Metric System These prefixes are based on powers of 10. What does this mean? From each prefix every “step” is either: 10 times larger or 10 times smaller For example Centimeters are 10 times larger than millimeters 1 centimeter = 10 millimeters kilo hecto deca Base Units meter gram liter deci centi milli
24 Metric System For each “step” to the right, you are multiplying by 10For example, let’s go from a base unit to centi 1 liter = 10 deciliters = centiliters 2 grams = 20 decigrams = centigrams ( 1 x 10 = 10) = (10 x 10 = 100) (2 x 10 = 20) = (20 x 10 = 200) kilo hecto deca Base Units meter gram liter deci centi milli
25 Metric Prefixes (see handout & reference book)We’ll talk about how to easily convert between the different prefixes later on in the unit! For now, you just need to recognize what the prefixes stand for when you see them in front of a base unit.
26 SI System Definition: modernized version of metric system; uses decimals All units derived from base units; larger and smaller quantities use prefixes with base unit Must memorize prefixes from nano- (10-9) to Giga- (109)
27 INSTRUMENTS & UNITS Use SI units — based on the metric system LengthMass Time Temperature Meter, m Kilogram, kg Seconds, s Celsius degrees, ˚C Kelvins, K
28 Length The standard unit of length in the metric system is the METERwhich is a little larger than a YARD. USING THE PREFIXES WITH LENGTH: cm – often used in lab km – Gm –
29 Units of Length 1 kilometer (km) = 1000 meters (m)10-2 meter (m) = 1 centimeter (cm) 102 meter (m) = 1 hectometer (Hm) 1 nanometer (nm) = 1.0 x 10-9 meter O—H distance = 9.58 x m 9.58 x 10-9 cm nm
30 the liter (milliliter) and cubic centimeter (cm3)Volume THE COMMON UNITS OF VOLUME IN CHEMISTRY ARE: the liter (milliliter) and cubic centimeter (cm3) THE COMMON INSTRUMENTS FOR MEASURING VOLUME IN CHEMISTRY ARE: graduated cylinder & buret Note that 1 cm3 = 1 mL (We will use this exact conversion factor throughout the year, so you will need to memorize it!)
31 Mass the gram (g) — often used in labTHE COMMON UNIT OF MASS IN CHEMISTRY IS : the gram (g) — often used in lab Mass IS A MEASURE OF THE AMOUNT OF MATTER IN AN OBJECT; Weight IS A MEASURE OF THE GRAVITATIONAL FORCE ACTING ON THE OBJECT. CHEMISTS OFTEN USE THESE TERMS INTERCHANGEABLY. 1000 g= 1 kg Mg = 10 6 g
32 TEMPERATURE IS THE FACTOR THAT DETERMINES the direction of heat flow.Temperature Scales Fahrenheit Celsius Kelvin Anders Celsius Lord Kelvin (William Thomson) TEMPERATURE IS THE FACTOR THAT DETERMINES the direction of heat flow.
33 Freezing point of waterTemperature Scales Fahrenheit Celsius Kelvin Boiling point of water 32 ˚F 212 ˚F 180˚F 100 ˚C 0 ˚C 100˚C 373 K 273 K 100 K Freezing point of water Notice that 1 kelvin degree = 1 degree Celsius
34 Calculations Using TemperatureFahrenheit/Celsius: TF = 1.8 TºC Celsius/Kelvin: TK = TºC + 273
35 Precision & Accuracy in MeasurementsPrecision vs. Accuracy Definitions: Precision—how close answers are to each other (reproducibility) Accuracy—how close answer is to accepted (true) value (agreement to accepted value)
36 Precision and Accuracy in MeasurementsPercent Error - a way to calculate accuracy in the lab Equation: % Error = | Accepted Value – Exp. Value | x 100 Accepted Value
37 Precision and Accuracy in MeasurementsExample #1: A student reports the density of a pure substance to be 2.83 g/mL. The accepted value is 2.70 g/mL. What is the percent error for the student’s results? Equation: % Error = | Accepted Value – Exp. Value | x 100 Accepted Value
38 See pages in homework packetScientific Notation Scientific Notation – See pages in homework packet
39 DETERMINING SIGNIFICANT FIGURESCrunching numbers with accuracy & precision
40 WHAT IS A SIGNIFICANT FIGURE?Significant figures = All the digits in a measurement that are known with certainty plus a last digit that must be estimated.
41 Significant Figures: Why are they Important?Numbers in math: no units, abstract, no context, can read calculator output exactly for answer. vs. Numbers in chemistry: measurements – include units. SIG FIGS WILL BE IMPORTANT THROUGHOUT THIS COURSE!
42 Graduated Cylinder Example
43 WHICH NUMBERS ARE SIGNIFICANT?For the purposes of significant figures there are two major categories: Nonzero digits: 1,2,3,4,5,6,7,8,9 Zero digits:
44 NONZERO DIGITS 257 L – 3 significant figuresAll nonzero digits are significant 3269 cm – 4 significant figures 257 L – 3 significant figures mm – 7 significant figures
45 Zeroes take three forms:Leading zeroes Trapped zeroes Trailing zeroes
46 LEADING ZEROES 0.000012 m – 2 significant digitsLeading zeroes are zeroes that come before the nonzero digits in a number. They are place holders only and are never considered significant. 0.123 L – 3 significant digits m – 2 significant digits mL – 4 significant digits
47 TRAPPED ZEROES Trapped zeroes are zeroes between two nonzero digits. Trapped zeroes are always significant. 101 s – 3 significant figures 20013 m – 5 significant figures cm – 4 significant figures (the leading zero is not significant)
48 TRAILING ZEROES Trailing zeroes are zeroes that follow nonzero digits. They are only significant if there is a decimal point in the number. mm – 4 significant figures 3000 s – 1 significant figure 250. mL – 3 significant figures g – 5 significant figures
49 Exact Numbers Any number which represents a numerical count or is an exact definition has an infinite number of sig figs and is NOT counted in the calculations. 12 inches = 1 foot 25 desks in the room
50 More Examples of Exact Numbers2 in 2r 3 and 4 in ¾ r3 Avogadro’s number is exactly x 1023 One inch is exactly 2.54 cm
51 Example #2: Now you try! 1.034 s 0.0067 g 12 apples 3000 m 72 peopleHow many significant digits are in each of the following: 1.034 s g 12 apples 3000 m 72 people
52 Answers 1.034 s - 4 significant figuresg significant figures 12 apples exact number 3000 m significant figure 69 people exact number
53 Process for Addition/SubtractionStep #1: Determine the number of decimal places in each number to be added/subtracted. Step #2: Calculate the answer, and then round the final number to the least number of decimal places from Step #1.
54 Addition/Subtraction ExamplesRound to tenths place. Example #2: Round to hundredths place. Example #3: Round to ones place. 23.456 24.706 Rounds to: 24.7 3.56 2.0699 2.07 14 26.735 27
55 Process for Multiplication/DivisionStep #1: Determine the number of sig figs in each number to be multiplied/divided. Step #2: Calculate the answer, and then round the final number to the least number of sig figs from Step #1.
56 Multiplication/Division ExamplesRound to 1 sig fig. Example #2: 2 sig figs. Example #3: 3 sig figs. 23.456 x x Rounds to: 1 3.56 x x 0.67 14.0/ 11.73 1.19
57 Example #3 Write the answers to the following computations using the correct number of sig figsa g g g = b m m = c × 4.20 × = d. 17 ÷ =
58 Important Rounding RuleWhen you are doing several calculations, carry out all the calculations to at LEAST one more sig fig than you need (I carry all digits in my calculator memory) and only round off the FINAL result.
59 Use of Conversion FactorsAlso known as dimensional analysis, factor-label method, or unit conversions Dimensional analysis/ Use of conversion factors Definition: technique to change one unit to another using a conversion factor Ex.) # in original unit x new unit = New # in new unit original unit
60 Using Dimensional AnalysisExample #4: Calculate the following single step conversions: a. How many Joules are equivalent to 25.5 calories if cal = joules? b. How many liters gasoline can be contained in a 22.0 gallon gas tank if L = 1 gal?
61 Example #5: The following multiple step conversions can be solved, knowing that 1 in = 2.54 cm. Convert the length of ft to millimeters.
62 Example #6: The average velocity of hydrogen molecules at 0oC is x 105 cm/s. Convert this to miles per hour.
63 Example #7: A piece of iron with a volume of 2.56 gal weighs lbs. Convert this density to scruples per drachm with the following conversion factors: 1.00 L = gal, kg = lb, scruple = g, mL = drachm.
64 Using Dimensional Analysis: Area ConversionsExample #8: Express the area of a 27.0 sq yd carpet in square meters. (1 yard = 3 feet, 1 foot = meters)
65 Volume Conversions Example #9: Convert 17.5 quarts to cubic meters.(1 L = qt, 1 ft3 = L)
66 Properties of Substances1. Every pure substance has its own unique set of properties that serve to distinguish it from all other substances. 2. Properties used to identify a substance must be intensive; that is, they must be independent of amount. Extensive properties depend on the amount. Classify the following as either intensive (I) or extensive (E): a. density b. mass c. melting point d. volume
67 Properties of SubstancesDensity is an INTENSIVE property of matter, which does NOT depend on quantity of matter. Contrast with EXTENSIVE properties of matter, which do depend on quantity of matter. Examples of extensive properties: mass and volume. Styrofoam Brick
68 Chemical and Physical PropertiesChemical property – Observed when the substance changes to a new one. Example of a chemical property: Copper reacts with air to form copper (II) oxide. Physical property – Observed without changing the substance to a new one. Example of a physical property: Water boils at 100oC.
69 Physical Changes Physical changes do not result in a new substance:boiling of a liquid melting of a solid dissolving a solid in a liquid to give a SOLUTION.
70 Physical vs. Chemical ChangeAnother name for a Chemical change is a chemical reaction — change that results in a new substance.
71 Example #10: Classify the following as either physical (P) or chemical (C) changes: a. ice melting b. gasoline burning c. food spoiling d. log of wood sawed in half
72 DENSITY : Review Definition: ratio of mass to volume for an object= Mass (g) volume (cm3) Aluminum Platinum Mercury 13.6 g/cm3 21.5 g/cm3 2.7 g/cm3
73 Density of Water Density of water changes with temperature(As water temperature changes, volume changes) Maximum density of water is at 4oC = g/cm3 (often rounded to 1.00 g/cm3)
74 Example #11: A piece of copper has a mass of g. It is 9.36 cm long, 7.23 cm wide, and 0.95 mm thick. Calculate the density (g/cm3).
75 Density as a Conversion FactorDensity is a “bridge” between mass and volume, or vice versa Volume (cm3) x density g = mass (g) cm3 Mass (g) density cm3 = Volume (cm3) g
76 Solve the problem using DENSITY AS A CONVERSION FACTOR.Example #12: Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in grams? In pounds? (Need to know conversion factor; 454 g = 1 lb) Solve the problem using DENSITY AS A CONVERSION FACTOR.
77 Example #13: What is the density of Hg if g occupy a volume of 12.1cm3?
78 Example #14: What is the mass of 2.15 cm3 of Hg?
79 Ex. #15: What is the volume of 94.2 g of Hg?
80 Example #16: Given the following densities: chloroform 1.48 g/cm3, mercury 13.6 g/cm3, and copper 8.94 g/cm3, calculate if a 50.0 mL container will be large enough to hold a mixture of 50.0 g of mercury, 50.0 g of chloroform and a 10.0 g chunk of copper.
81 Example #17: How many kilograms of methanol (d = g/mL) does it take to fill the 15.5-gal fuel tank of an automobile modified to run on methanol?
82 Derived Units Definition: derived from base unitsExample: m/sec (unit of speed) Divide meters by seconds Volume examples m3 (m x m x m) or cm3 (cm x cm x cm)