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Matter and Measurement

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  1. Matter and Measurement Chapter 1

  2. What Is Matter? • Matter is anything that takes up space and has mass. • Mass is the amount of matter in an object. • Mass is resistance to change in motion along a smooth and level surface.

  3. Types of Matter • Substance- a particular kind of matter – pure • Fixed composition • Distinct properties • Ex. Water, Salt • Mixture- more than one kind of matter • Compositions vary • Each substance retains its own chemical identity and properties.

  4. Substances • Elements- simplest kind of matter • Cannot be broken down into simpler • All one kind of atom. • Compounds- are substances that can be broken down by chemical methods • When they are broken down, the pieces have completely different properties than the compound. • Made of molecules- two or more atoms

  5. Mixtures • Heterogeneous- mixture is not the same from place to place. • Chocolate chip cookie, gravel, soil. • Variable Composition • Homogeneous- same composition throughout. • Kool-aid, air. • Every part keeps its properties.

  6. Solutions Homogeneous mixture • Mixed molecule by molecule • Can occur between any state of matter. • Solid in liquid- Kool-aid • Liquid in liquid- antifreeze • Gas in gas- air • Solid in solid - brass • Liquid in gas- water vapor

  7. Solutions • Like all mixtures, they keep the properties of the components. • Can be separated by physical means • Not easily separated- can be separated.

  8. Gold Gold Copper Silver 24 karat gold 18 karat gold 14 karat gold 24/24atoms Au 18/24atoms Au 14/24atoms Au

  9. Solid Brass • An alloy is a mixture of metals. • Brass = Copper + Zinc • Solid brass • homogeneous mixture • a substitutional alloy Copper Zinc

  10. Brass = Copper + Zinc • Brass plated • heterogeneous mixture • Only brass on outside Brass Plated Copper Zinc

  11. Steel Alloys • Stainless steel • Tungsten hardened steel • Vanadium steel • We can engineer properties • Add carbon to increase strength • Too much carbon  too brittle and snaps • Too little carbon  too ductile and iron bends

  12. Nitinol Wire • Alloy of nickel and titanium • Remembers shape when heated Applications: surgery, shirts that do not need to be ironed.

  13. Properties • Physical Properties- A change that changes appearances, without changing the composition, thus identity is preserved. • ex. color, odor, density…. • Chemical Properties- a property that can only be observed by changing the type of substance. • ex. flammability

  14. Physical and Chemical Changes

  15. Intensive Properties Do not depend on on the amount of sample being examined Aid in the identification of a substance. ex. temperature, density, melting point…. Extensive Properties Depend on the quantity of sample ex. mass, volume, area…. Intensive Vs. Extensive Properties

  16. The SI System •The SI system has seven base units from which all others are derived

  17. SI Units (Con’t) • These prefixes indicate decimal fractions or multiples of various units

  18. At home you like to keep the thermostat at 72 F. While traveling in Canada, you find the room thermostat calibrated in degrees Celsius. To what Celsius temperature would you need to set the thermostat to get the same temperature you enjoy at home ? • °C = 5/9 ( °F-32) °F = 9/5 (°C ) +32 K = °C + 273.15 Temperature Conversions

  19. Derived Units • SI units are used to derive the units of other quantities. • These units express speed, velocity, area and volume. • They are either base units squared or cubed, or they define different base units

  20. Volume • Calculated by multiplying L x W x H • Basic SI unit of volume is the cubic meter (m3 ). • Smaller units are sometimes employed ex. cm3, dm3 …. • Volume is more commonly defined by liter (L).

  21. Density Density is an intensive property of all substances; this Means that it is a universal characteristic of the substance and does not change because of extensive or accidental properties such as amount, size or location. Density is the relationship between the mass and volume of an object. It is defined as: • Density = Mass of substance Volume of substance • Density is usually expressed in g/ml

  22. Specific Gravity Specific Gravity is a ratio between the density of a substance and the density of water. It is defined by: Specific Gravity = Density of Sample Density of Water In calculations involving s.g., the units of density must match in order for the units to cancel

  23. Problems Involving Density 1. The mass of 325 mL of the liquid methanol is found to be 257 g. What is the density of methanol? 2. A lead weight used in the belt of a scuba diver has a mass of 226g. When the weight is carefully placed in a graduated cylinder containing 200.0ml of water, the water level rises to 220.0ml. What is the density of the lead weight (g/ml) • A sample of a certain material has a mass of 2.03x 10-3 g. Calculate the • Volume of the sample given that the density is 9.133 x 10-1 g/cm3

  24. Exact numbers • Numbers whose values are known exactly • e.x. 12 eggs in a dozen, 1000g in a kg Inexact numbers • Numbers obtained by measuring a quantity • e.x. height, weight or temperature • There is always a degree of uncertainty in measured values. Why? Uncertainty and Measurement

  25. Precision Vs Accuracy • When comparing sets of data points, scientist want to know two things: which set of data points are precise and which set are accurate.

  26. Precision • How many times a given measurement can be repeated with results close in value to each other. • Which of the following data is more precise: 119g, 120g, 128g or 101g, 100g, 99g Good precision poor accuracy

  27. Accuracy • How close an experimental measurement is to the true, actual, or book value • Usually the more accurate a measurement the more precise it will be. Good Accuracy and Precision

  28. Precision and Accuracy • Accuracyrefers to the agreement of a particular value with the truevalue. • Precisionrefers to the degree of agreement among several measurements made in the same manner. Precise but not accurate Precise AND accurate Neither accurate nor precise

  29. Significant Figures • When using a measuring device to measure anything the last number of the measurement is always estimated • Measured quantities are usually reported in such a way that only the last digit is uncertain. All digits including the uncertain digit are called significant figures

  30. A number is not significant if it is: • A zero at the beginning of a decimal number ex. 0.0004lb, 0.075m • A zero used as a placeholder in a number without a decimal point ex. 992,000,or 450,000,000 A number is a S.F. if it is: •Any real number ( 1 thru 9) • A zero between nonzero digits ex. 2002g or 1.809g • A zero at the end of a number or decimal point ex. 602.00ml or 0.0400g • Any digit in the coefficient of a number written in scientific notation ex. 4.0 x 105 m, 5.70 x 10-3 Significant Figures

  31. Sig Fig Practice #1 How many significant figures in each of the following? 1.0070 m  5 sig figs 17.10 kg  4 sig figs 100,890 L  5 sig figs 3.29 x 103 s  3 sig figs 0.0054 cm  2 sig figs 3,200,000  2 sig figs

  32. When adding or subtracting the answer is rounded so it has the same number of decimal places as the number with the least number of decimal places 6.23 cm 39.24 cm + 677.1 cm 722.6 cm Addition and Subtraction

  33. When multiplying or dividing the answer has the same number of significant digits as the number with the least number of S.F. 2.85ml x 67.4ml = 192ml ml 49.618g = Multiplication and Division 43.8ml

  34. Sig Fig Practice #2 Calculation Calculator says: Answer 10.24 m 3.24 m + 7.0 m 10.2 m 100.0 g - 23.73 g 76.3 g 76.27 g 0.02 cm + 2.371 cm 2.39 cm 2.391 cm 713.1 L - 3.872 L 709.228 L 709.2 L 1821.6 lb 1818.2 lb + 3.37 lb 1821.57 lb 0.160 mL 0.16 mL 2.030 mL - 1.870 mL

  35. Odd Numbers • Odd numbers are rounded up, even numbers are left alone when the remainder is 5 ex. 22.15 g if rounded to 3 SF = 22.2g 22.25 g if rounded to 3 SF = 22.2

  36. Scientific Notation In science, we deal with some very LARGE numbers: 1 mole = 602000000000000000000000 In science, we deal with some very SMALL numbers: Mass of an electron = 0.000000000000000000000000000000091 kg

  37. . 2 500 000 000 9 7 6 5 4 3 2 1 8 Step #1: Insert an understood decimal point Step #2: Decide where the decimal must end up so that one number is to its left Step #3: Count how many places you bounce the decimal point Step #4: Re-write in the form M x 10n

  38. 2.5 x 109 The exponent is the number of places we moved the decimal.

  39. 0.0000579 1 2 3 4 5 Step #2: Decide where the decimal must end up so that one number is to its left Step #3: Count how many places you bounce the decimal point Step #4: Re-write in the form M x 10n

  40. 5.79 x 10-5 The exponent is negative because the number we started with was less than 1.

  41. Review: Scientific notation expresses a number in the form: M x 10n n is an integer 1  M  10

  42. Conversion Factors Fractions in which the numerator and denominator are EQUAL quantities expressed in different units Example: 1 in. = 2.54 cm Factors: 1 in. and 2.54 cm 2.54 cm 1 in.

  43. Learning Check Write conversion factors that relate each of the following pairs of units: 1. Liters and mL 2. Hours and minutes 3. Meters and kilometers

  44. How many minutes are in 2.5 hours? Conversion factor 2.5 hr x 60 min = 150 min 1 hr cancel By using dimensional analysis / factor-label method, the UNITS ensure that you have the conversion right side up, and the UNITS are calculated as well as the numbers!

  45. Dimensional Analysis • Many problems in chemistry and health sciences require a change in units, thus dimensional analysis is used to make these conversions • The key is the correct use of a conversion factor to change 1 unit to another ex. 1hr = 60 min this allows us to write the relationship: 60 min and 1hr 1hr 60min

  46. Problems Involving Dimensional Analysis • On a bicycle trip, Maria averaged 35 miles per day. How many days did it take her to cover 175 miles? • To prevent bacterial infection, a doctor orders 4 tablets of amoxicillin per day for 10 days. If each tablet contains 250 mg of amoxicillin, how many ounces of the medication are given in 10 days?

  47. Solution to Problems Step 1: Given 175 miles 35 miles = 1 day Step 2: Unit plan MilesDays 175 miles x 1 day = 5.0 days 35 miles

  48. Solution to Problem Step 1: Given 10 days 4 tablets/day 250 mg each Step 2: Unit Plan mgglboz 1000mg x 1g x 1lb x 16 oz = 0.35 oz 1000mg 453.59g 1lb

  49. Dealing with Two Units If your pace on a treadmill is 65 meters per minute, how many seconds will it take for you to walk a distance of 8450 feet?

  50. Learning Check • A Nalgene water bottle holds 1000 cm3 of dihydrogen monoxide (DHMO). How many cubic decimeters is that?