Units of Measurement CHEM 1211 PDF

Units of Measurement CHEM 1211 Tools How to record measurements (SI units) How to record measurements (significant figures) Calculating with significant figures Scientific Notation Mass Mass: A measure of how much “stuff” in an object. Units – official SI unit is kilogram (kg)....

Units of Measurement CHEM 1211 Tools How to record measurements (SI units) How to record measurements (significant figures) Calculating with significant figures Scientific Notation Mass Mass: A measure of how much “stuff” in an object. Units – official SI unit is kilogram (kg). It is defined by the mass of a cylinder of platinum-indium alloy in France. 1 kilogram = 2.2 pounds on Earth Mass of a nickel is 5.000 g Length The standard SI unit of length is the meter (m). It is defined as the distance light in a vacuum travels in 1/299,792,458 of a second.  Volume A measure of how much space the object occupies. – SI unit is m3. Lab units include mL and cm3 – A die is about 1 cm3. Common lab unit is the Liter (L) or mL. – 1 L = 1000 mL = 103 mL – 1 L = 103 cm3 – 1 mL = 1 cm3. Volume Water displacement: When an object enters water, it pushes out water to make room for itself. The object pushes out a volume of water that is equal to its own volume. This is called displacement. Density Density = mass / volume D = m/V Units: any mass unit over any volume unit Examples: g/mL, g/cm3, kg/L, kg/m3 Temperature The measurement of the average kinetic energy of the particles in a substance. – Higher temperature= higher energy= more movement. Celsius & Fahrenheit Based on boiling and freezing temperatures of water. Kelvin – K= °C + 273.15 Precision vs Accuracy Uncertainty – How precision applies to chemistry devices (e.g., beaker, Erlenmeyer flask, graduate cylinder) – In chemistry each device will give you a different accuracy and precision. – Example: The uncertainty in a measure is the degree or error in a measurement. Scientific Notation A number written in scientific notation contains a coefficient and a power of ten. To write a number in scientific notation – Decimal point is moved after the first nonzero digit – Space moved are shown as a power of ten Scientific Notation To add or subtract using scientific notation, each number must conform to the same exponent. Example: (5.78 x 105 ) + (3.1 x 104 ) = (5.78 x 105 ) + (0.31 x 105 ) = (6.09 x 105 ) Scientific Notation To multiply or divide using scientific notation, we multiply the numbers together as usual and then add or subtract the exponents. Example: (6.0 x 104) x (5.0 x 103) = (6.0 x 5.0) x (104+3) = (30 x 107) = (3.0 x 108)  Significant Figures Precision of a measurement is indicated by the number of digits reported. – These reported digits are called significant figures. – Look to the right of the last significant figure for rounding. Significant Figure Examples Measurement Number of Significant Figures 42.38 cm 4 6.2 ft 2 32.9 kg 3 Leading zeroes are not significant Measurement Number of Significant Figures 0.007 mm 1 0.0152 g 3 0.00369 mL 3 Zeroes between nonzero numbers are significant Measurement Number of Significant Figures 40.7 mm 3 2002 mL 4 0.50505 g 5 Significant Figure Examples Zeroes that follow nonzero numbers in measurements with no decimal point are not significant Measurement Number of Significant Figures 200 g 1 35000 mm 2 250500 mL 4 All digits including zeroes in the coefficient are significant Measurement Number of Significant Figures 9 x 104 g 1 9.0 x 104 g 2 9.00 x 104+ g 3 Calculations and Significant Figures Multiplication and division. Do the calculation. Round the number on your calculator to the number of digits in the measurement with the fewest. Don’t pay attention to where the decimal is. Just count total significant digits no matter which side of the decimal. Rounding – the usual rounding rules apply. Addition and subtraction: Don’t pay attention to total number of digits. Just count places to the right. You are limited by the measurement which has the fewest significant digits to the right.