Quantitative Qualitative precision accuracy both neither, Slides of Chemistry

Download Quantitative Qualitative precision accuracy both neither and more Slides Chemistry in PDF only on Docsity! Unit 2: Data Collection and Analysis General Chemistry Page 1 of 10 I. Measurement and Observation There are two basic types of data collected in the lab: Quantitative : numerical information (e.g., the mass of the salt was 3.45 g) Qualitative : non-numerical, descriptive data (e.g., the color of the solution is magenta). Uncertainty in Measurement When you carry out an experiment or measurement you need to understand the true quality of your results. The terms scientists typically use are accuracy and precision—they are not the same. 1. Accuracy refers to degree of conformity with a standard (often called true, accepted or theoretical) value. There are times when a calculated value will be used as the standard. 2. Precision refers to how close measurements are to one another. Repeated measurements determine reproducibility or precision. Precision tells you how to report results. precision accuracy both neither Accuracy and Precision Four lab groups performed the same experiment three times to determine the melting point of naphthalene (moth balls). The accepted melting point is 79.0°C. Indicate whether the following sets of data are precise, accurate, both or neither. Precise, Accurate, Both or Neither Reasoning Group Trial 1 Trial 2 Trial 3 accurate Average of the trials is close to the accepted melting point of 79.0 1 76.2°C 79.5°C 81.3°C precise All trials have values that are close to each other 2 76.2°C 76.1°C 76.3°C neither The trials are neither close to each other (precise) or close to the accepted value of 79.0 3 86.4°C 82.8°C 81.2°C both All trials are precise and close to the accepted value of 79.0 4 79.1°C 78.9°C 79.2°C Unit 2: Data Collection and Analysis General Chemistry Page 2 of 10 Glassware Qualitative or Quantitative Function Beaker qualitative Large mouth glass containers used to contain approximate volumes of liquid. Buret quantitative Long tube with a stopcock that opens and closes. It is used to precisely deliver solutions, especially in a titration. Erlenmeyer Flask qualitative Glass container used to contain approximate volumes of liquid. Small mouth accommodates a stopper for storage or shaking. Graduated Cylinder quantitative Used to measure and deliver approximate volumes of liquids. Pipet quantitative Used to precisely deliver variable quantities of liquid. Test Tube qualitative Glass cylinder that holds liquids being tested in an experiment. Volumetric Flask quantitative Designed to precisely contain a specific volume. Commonly used when accurately making aqueous solutions. ***In trying to decide which piece of equipment is the most accurate, always choose the one with the smallest measurement units and smallest diameter. II. Measurement and Significant Figures Results should always be reported to the correct number of significant figures. These will be discussed in more detail in the next unit. When making a measurement in the lab, always report the number of digits necessary to express results of measurement consistent with the measured precision. This means you are to report all certain digits plus one uncertain digit. Every time you take a measurement you should estimate between the lines. If the measurement is on a line, add a zero to show that you are estimating it to be exactly on the line. Always include one estimated digit. Remember that liquids form a curved surface called a meniscus. Measure to the bottom of the meniscus. A buret precisely measures the amount of liquid that is released through the stopcock. This is why a buret is marked “upside-down” compared to a graduated cylinder. The numbers increase going down a buret. Be careful of this when reading burets. Unit 2: Data Collection and Analysis General Chemistry Page 5 of 10 Example 2.5 Perform the following mathematical functions and express the answers with the correct number of significant figures: 0.006 760 ÷ 32 1,234,000 ÷ 0.0000345 278.4 × 25.2 89.554 × 43.1 0.00021 3.58×1010 7020 3860 IV. Scientific Notation Scientific notation is used to represent numbers that are very large or very small. Rules for Scientific Notation To convert from decimal form to scientific notation: Move the decimal point to the left or the right so that only one nonzero digit remains to the left of the decimal point. The exponent is the number of places that you moved the decimal point. If you moved the decimal to the left, the exponent is positive. If you moved it to the right, the exponent is negative. To convert from scientific notation to decimal form: Move the decimal point to the right if the exponent is positive (add zeroes if needed). Move the decimal to the left if the exponent is negative (add zeroes if needed). A calculator can automatically show numbers in scientific notation if it is in scientific mode: It can automatically show numbers in decimal form if it is in floating point mode: Regardless of the mode in which the calculator is set, numbers in scientific notation should be entered using the “EE” button. Do NOT enter scientific notation using “× 10” or the “^” or “10x ” buttons. These will make it more difficult to get the correct order of operations during calculations. To enter 1.0×10-14 in scientific notation: Example 2.6 Convert the following numbers from decimal form to scientific notation: 75,100,000 -234,900 0.000 002 31 -0.000 035 49 7.51×107 -2.349×105 2.31×10-6 -3.549×10-5 ◄2nd SCI/ENG DRG select SCI ENTER ═ ►2nd SCI/ENG DRG select FLO ENTER ═ 1 . 0 2nd EE x-1 ( – ) 1 4 ENTER ═ Unit 2: Data Collection and Analysis General Chemistry Page 6 of 10 Example 2.7 Convert the following numbers from scientific notation to decimal form: 1.12×103 -2.35×105 1.12×10-3 -2.35×10-5 1,120 -235,000 0.001 12 -0.000 023 5 To correct INCORRECT scientific notation: Move the decimal point to the left or the right so that only one nonzero digit remains to the left of the decimal point. Increase the exponent if you moved the decimal to the left. Decrease the exponent if you moved it to the right. Example 2.8 Correct the following incorrect scientific notation: 36.7×101 -0.015×10-3 0.123×104 851.6×10-3 3.67×102 -1.5×10-5 1.23×103 8.516×10-1 Calculations in scientific notation: (Your calculator takes care of this for you.)  Addition and Subtraction: Exponents must be the same.  Multiplication: Multiply the coefficients and add the exponents.  Division: Divide the coefficients and subtract the exponents. Example 2.9 Perform the following mathematical functions and express the answers in correct scientific notation: 3.20×10 3 + 9.77×10 2 3.20×10 3 - 9.77×10 2 3.20×10 3 × 9.77×10 2 3.20×10 3 ÷ 9.77×10 2 4.18×103 2.22×103 3.13×106 3.28 X. Algebraic Manipulation Example 2.10 Rearrange the following equations to solve for the variable that is in bold/italics: V D m  PV = nRT K = °C + 273 m = DV D m V  nT PV R  °C = K - 273 V m D  Unit 2: Data Collection and Analysis General Chemistry Page 7 of 10 XI. Density Density is the mass of a substance per unit volume or how much it weighs per given volume. It is an intensive physical property. V mass D  The units for mass are grams. For liquids, the units for volume are milliliters and the units for density are grams/milliliter. For gases, the units for volume are liters and the units for density are grams/liter. Remember: 1 cm 3 = 1 mL. Water has a density of about 1.0 g/mL. Substances with densities less than 1.0 g/mL float on water. Substances with densities greater than 1.0 g/mL sink in water. Example 2.11 Is ice more or less dense than liquid water? Ice floats on water, therefore it is less dense. Example 2.12 A certain solid has a volume of 35.7 cm 3 and a mass of 85 grams. What is its density? 3cm g 3 2.4 cm35.7 g 85 D  Example 2.13 The density of liquid mercury is 13.6 g/mL. What is the mass of 35.0 mL of mercury?    g 476mL 35.0 13.6mass mL g  Example 2.14 If the density of gold is 19.3 g/cm 3 , what is the volume of 200 g of gold? 3 3cm g cm 10 19.3 g 200 V  Example 2.15 Find the density of a 500. g rectangular solid whose dimensions are 3.4 cm by 1.2 cm by 1.7 cm. V = (3.4 cm)(1.2 cm)(1.7 cm) = 6.936 cm3 (Don’t round significant digits until the end.) 3cm g 3 72 cm 6.936 g 500. D  Example 2.16 An empty graduated cylinder weighs 26.5 grams. When it is filled with an unknown liquid up to the 45.8 mL mark, the cylinder and the liquid together weigh 70.0 grams. What is the density of the unknown liquid? mass = 70.0 g – 26.5 g = 43.5 g mL g 0.950 mL 45.8 g 43.5 D 