Popular in Chemistry 1
Popular in Chemistry
This 6 page Study Guide was uploaded by Kacey Lange on Sunday August 21, 2016. The Study Guide belongs to CH 1213 at Mississippi State University taught by Erin Dornshuld in Fall 2015. Since its upload, it has received 47 views. For similar materials see Chemistry 1 in Chemistry at Mississippi State University.
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Date Created: 08/21/16
Chapter 1: Essential Ideas The Scientific Method 1. Observation and curiosity 2. Form hypothesis and make predictions 3. Perform experiment and make more observations If results are not consistent, reverse back to step two and alter hypothesis. If results are consistent, continue to step three. 4. Contributes to body of knowledge If testing does not support the hypothesis, reverse back to step two. If testing does support the hypothesis, continue to step five. 5. Observation becomes law OR hypothesis becomes theory Observation becomes law because testing produces constant observations. Hypothesis becomes theory because testing supports the hypothesis. Domains 1. Macroscopic large enough to be examined by the naked eye 2. Microscopic too small to be visible to the naked eye, but visible when using a microscope. 3. Symbolic diverse organisms that live together. Phases and Classification of Matter Matter is anything that occupies space and has mass. So, everything is matter. YOU DO MATTER Mass is a measure of the amount of matter in an object Weight is the force (mass times acceleration) that gravity exerts on an object. Solid objects have a distinct shape and volume. Liquids do not have a distinct shape, but has a fixed volume. Gases do not have a distinct shape or fixed volume. Law of Conservation The Law of Conservation states that matter is neither created or destroyed. It is only transformed from one type or phase to another. An atom is the smallest particle of an element that has the properties of that element. An atom in indivisible. A molecule consists of two or more atoms that are joined together by forces. A chemical bond is the bond that binds molecules together. Classifying Matter A pure substance is a form of matter that has a definite composition and distinct properties such as color, smell, and taste. Pure substances are homogeneous and have a constant composition. An example of a pure substance is gold, oxygen, or water. Compounds are pure substances that can be broken down via chemical changes. Ex. Mercury (II) oxide decomposes into silvery droplets of mercury and oxygen gas once heated. Elements are substances that are composed of a single type of atom. Elements are pure substances that cannot be broken down into simpler substances by chemical changes Elements are defined by the number of protons in the atom A mixture is composed of two or more types of matter and can be separated by physical changes Ex. Evaporation A heterogeneous mixture is a mixture that varies from point to point (not uniform) and you can see different parts of the mixture. Ex. Cereal and milk, bucket of soil A homogeneous mixture is a mixture that has a uniform composition and appears visually the same throughout. Ex. Gatorade, coffee Physical and Chemical Properties A physical property is a characteristic of matter that is not associated with a change in its chemical composition Ex. Density, color, melting point, boiling point A physical change is a change in the state or properties of matter without any accompanying change in its chemical composition. Ex. Wax melts from solid to a liquid A chemical property is an observable property of matter that undergoes a chemical reaction. Ex. Flammability, acidity, reactivity Ex. If an element rusts or doesn’t rust. A chemical change is a change that always produces one or more types of matter that differ from the matter present before the change. An extensive property depends on the amount of matter present (mass and volume) An intensive property does not depend on the amount of matter present (temperature and density) The periodic table organizes elements based upon certain properties of those elements. Measurements Units are standards of comparison for measurements. Ex. Liters, pounds, and centimeters We usually report the results of scientific measurements in SI Units, which are International System of Units. Scientific notation is used to express large or small quantities conveniently. Positive exponential when moving left. Negative exponential when moving right. Converting between metric units Express 500 m in km. 500x10 3=0.5km5.0x10 km 1 Express 500 m in μm. 500x10 =500,000,000 μm5.0x10 μm 8 Express 52 pm in km. 52pm x (1m/1x10 pm) x (1x10 km/1m) = 5.2x10 km 14 Converting between temperatures The Kelvin Scale is the concept of absolute zero, the lowest temperature theoretical possible. Kelvin = C + 273.15 o o o Convert 450 Kelvin to Fahrenheit. K= C+273.15 C=K273.15 C=450273.15 o C=176.85 oF = 9/5 x ( C) + 32 o o o o Convert 176.85 C to Fahrenheit. F=9/5(176.85)+32 F= 318.33 + 32 F = 350.33 Time Conversions SI base unit for time is the second(s) 12 How many years is 8.523 x 10 s (in scientific notation to one decimal place)? 8.523 x 10 s (1min/60s)(1hr/60min)(1day/24hr)(1yr/365days)= 2.7 x 10 years 5 QUALITATIVE – DESCRPITION OF SOMETHING THAT INCLUDES HOW THEY LOOK/FEEL/SMELL QUANTITATIVE A DESCRIPTION OF SOMETHING THAT CONTAINS AN AMOUNT OR NUMBER Derived SI Units Volume is the measure of the amount of space occupied by an object The standard volume is given as the cubic meter (m ) 3 The SI unit for volume is liter (L) Volume = (Length) 3 Derived SI Units are derived from base SI units Density is the ratio of the mass of a sample of a substance and the volume it occupies (mass per unit volume) SI units for density is kg m (derived from the SI unknit for mass and volume) A common unit for density is g cm 3 Density = mass/volume Volume and density conversions 8 Convert 1.85 x 10 L to Nl 1.85 x 10 L (1x10 Nl/ 1L) = 1.85x10 Nl OR 18.5 Nl Convert 0.85 g Ml to kg Cl 1 4 0.85g > 0.85g x (1kg/1000g) = 8.5x10 kg 1ml 1 ml (1L/1000ml) = 0.001L(100cl/1L) = 0.1cl Density = 8.5x10 kg/01.cl = 8.5x10 kg cl 1 Density Example What mass of gold (in kg) must occupy a cube of gold that has an edge length of 2.00 3 cm if the density of gold is 19.3 g cm ? Mass = ? Length = 2.00 Density = 19.3 g cm 3 3 3 3 Volume = (length) = (2.00) = 8.00 cm 19.3 g x (1kg/1000g) = 0.0193 kg Density = mass/volume 0.0193 = mass/8.00cm mass = 0.1544 kg Measurement Uncertainty, Accuracy and Precision Exact numbers are obtained only through the measurement of counting; exact quantities that have zero uncertainty; includes defined quantities. Number of shoes you are wearing, number of eggs in a carton of eggs, weight of a 1 lb bag of rice Inexact Numbers are obtained by any measurement other than counting; inexact quantities that have a level of uncertainty Your weight, the number of grains of rice in a 1 lb bag Significant Figures All digits in a measurement (including the uncertain last digit) are called significant figures or significant digits Any digit that in not zero is significant Zeroes located between nonzero digits are significant Zeroes to the left of the first nonzero digit are not significant Zeroes to the right of the last nonzero digit are significant is the number contains a decimal point Trailing zeroes in a number that does not contain an explicit decimal point may or may not be significant. Exact number have an infinite number of significant figures For a number in scientific notation, all digits comprising N are significant by the first 6 rules Calculations with Inexact Number When adding or subtracting, round the result to the same number of decimal places as the number with the least number of decimal places (i.e. the least precise value in terms of addition and subtraction) When multiplying or dividing, round the result to the same number of digits as the number with the least number of significant figures (the least precise value in terms of multiplication and division) Exact numbers can be considered to have an infinite number of significant figures and do not limit the number of significant figures in a calculated result. In calculations with multiple steps, rounding the result of each step can result in “rounding error.” Retain at least one extra digit until the end of a multistep calculation to minimize rounding error. Don’t round until very end. Examples of Rounding Number Rounding Rules >5 rounds up <5 rounds down An exact 5 rounds to the nearest even number For multistep calculations, do not round intermediate steps. Keep track of sig. figs. by underlining the last significant digit in each step. Round at the very end of the calculation (rule #4). Accuracy and Precision Accuracy tells how close a measurement is to true value Precision tells how close a series of replicate measurements are one to another. Dimensional Analysis A conversion factor is a ratio of two equivalent quantities expressed with different measurement units You should be able to implement any conversion factor (even fictional ones). Simply track your units and determine how to write out the fraction 1 H = 627.510 kcal mol 1 Ex.
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