Chapter 1: Keys to the Study of Chemistry Chemistry is the study of matter, its properties, the changes that matter undergoes, and the energy associated with these changes The topics in this chapter should be review from a previous course. It is expected that you are able to review and master this material quickly and somewhat independently. From this Chapter you should: understand the basics of the Scientific Method. be able to identify physical and chemical properties and changes. be able to apply the kinetic molecular theory to the properties of matter. develop proficiency with metric units and dimensional analysis. know the meaning of the terms precision and accuracy. be able to use significant figures properly, understanding how they relate to uncertainty in measurements. 1 The Scientific Method The scientific method provides guidelines for the practice of science. Collect data (observe, experiment, etc.). Look for patterns. Try to explain data; develop a hypothesis or tentative explanation. Test hypothesis, then refine/revise the hypothesis if experimental results do not support it. Bring information together into a scientific law (natural law), a concise statement or mathematical equation that summarizes or describes the behavior of matter. Bring hypotheses and laws together into a theory. A theory explains what causes certain phenomena and can be used to make predictions. A theory must have considerable evidence or facts to support it. Test predictions based on theory. Modify theory if if experimental results do not support it. Test Theory 2
Definitions (These you should already know!) Matter Composition Properties 3 Physical and Chemical Changes/Properties Each pure substance has a unique set of physical and chemical properties that can be used to identify it. These you should already understand! Give some examples of physical properties: In physical changes the identity of the substance is preserved. Examples? Physical change - Napthalene melts. Give some examples of chemical properties: In chemical changes new substances are produced, the identity of the substance changes. Examples? Chemical change - Formation of water from its elements. 4
Solid The States of Matter (These you should already know!) Liquid Gas 5 Energy in Chemistry Energy is the ability to do work. Potential Energy Kinetic Energy Total Energy = Potential Energy + Kinetic Energy Energy is conserved. (First Law of Thermodynamics) 6
Potential Energy Between Charged Particles (a fundamental concept in understanding chemistry) Lower energy states are more stable and are favored over higher energy states. Energy can be converted from one form to another, but the total energy is conserved. A system of oppositely charged particles. The potential energy gained when the charges are separated is converted to kinetic energy as the attraction pulls these charges together. 7 Chemical Potential Energy Converted to Kinetic Energy (a fundamental concept in understanding chemistry) A system of fuel and exhaust. A fuel is higher in chemical potential energy than the exhaust. As the fuel burns, some of its potential energy is converted to the kinetic energy of the moving car and to heat energy that is lost to the surroundings. 8
Kinetic Molecular Theory Matter consists of small particles; atoms, molecules or ions that are in constant, random motion. As temperature increases, the average kinetic energy of the particles increases and the motion becomes more rapid. HEAT FLOW: Matter has three physical states. Physical state depends upon the nature of the substance (its identity), its temperature and the pressure exerted upon the substance. As chemists we describe these states at either the macroscopic level or microscopic (particulate) level. Conversion between these states is a physical process. Solid Liquid-solid Solid-gas 9 Kinetic Molecular Theory and Physical State 10
Temperature is a measure of average kinetic energy of the particles within matter. Temperature We will use the Kelvin (K) and Celsius scales ( C). Be able to convert between them: K = C + 273.15 C = K - 273.15 11 Intensive vs. Extensive Properties of Matter Extensive properties: Examples? Intensive properties: Examples? 12
Know the SI units of mass, length, time, temperature, and amount of a substance. Table 1. 2 Physical Quantity (Dimension) SI Base Units Unit Name Units of Measurement Unit Abbreviation Mass kilogram kg Length meter m Time second s Temperature kelvin K Electric Current ampere A Amount of substance mole mol Luminous intensity candela cd Know the most common metric prefixes! (mega, kilo, deci, centi, milli, micro, nano and pico) Table 1.3 Common Decimal Prefixes Used with SI Units 13 Derived Units of Measurement Formed using the SI base units. For example, velocity is distance traveled per unit time, thus m/s. Volume Density Table 1.5 Densities of Some Common Substances * Substance Physical State Density (g/cm 3 ) Hydrogen gas 0.0000899 Oxygen gas 0.00133 Grain alcohol liquid 0.789 Water liquid 0.998 Table salt solid 2.16 Aluminum solid 2.70 Lead solid 11.3 Gold solid 19.3 * At room temperature (20 C) and normal atmospheric pressure (1atm). 14
Other Derived Units Give the derived SI units for each of the following quantities in base SI units: a) acceleration = distance/time 2 b) force = mass acceleration c) work = force distance d) pressure = force/area e) energy = mass x (velocity) 2 15 Uncertainty and Errors in Measurements: Accuracy and precision are two very different concepts. You are expected to have a complete understanding of these! Accuracy: Systematic Error: Percent Error (MEMORIZE FORMULA!): Precision: Random Error: Significant Figures: Range (MEMORIZE FORMULA): 16
Significant Figures, Precision and Device Every measurement includes some uncertainty. The rightmost digit of 10-mL Graduated Cylinders any measured quantity always contains some random error and is estimated. For graduated glassware, the estimated digit is dependent on the scale. For digital devices (thermometers and balances), the estimated digit is the last digit shown on the read out and ALL digits should be recorded, unless directed otherwise. The number of significant figures in a measurement, such as 52.8 ml for the 100-mL graduated cylinder shown below, is equal to the number of digits that are known with confidence (5, and 2) plus the last digit (8), which is an estimate or approximation. The measuring device used determines which digit is estimated, and thus the precision of the measurement. Analytical Balance 100-mL Graduated Cylinder Estimate the level of the liquid between 52 and 53 ml. This estimated digit is the last digit that should be recorded. 200-mL Beaker 17 Example: Sig Figs and Precision vs Accuracy Three students each weigh a dry 50-mL graduated cylinder, add 25.0 ml of water to the cylinder, and then weigh the cylinder plus water. The difference in the masses is calculated to give a measured mass of the water in the cylinder. All students use the same cylinder and same balance. Each student performs this operation four times. Student A s Trial 1 data is as follows: Mass of dry cylinder: Mass of cylinder plus water: 110.0 g 135.2 g What is the measured mass of the water in the cylinder for this trial? Report answer with the correct sig. figs. If the density of the water is 1.000 g/ml at the temperature of the experiment, what is the actual mass of 25.0 ml of water? (Use correct sig. figs. in your answer.) The tables and graphs that follow on the next page summarize the results of the measurements for each student. 18
Example: Sig Figs and Precision vs Accuracy Trial Number Water Mass (g) Trial Number Water Mass (g) Trial Number Water Mass (g) 1 25.2 2 25.1 3 24.8 4 24.9 Average 25.0 Range % Error 1 27.0 2 26.8 3 26.8 4 27.1 Average 26.9 Range % Error 1 23.9 2 26.5 3 25.7 4 24.0 Average 25.0 g Range % Error 19 Unit Conversions- Dimensional Analysis English < > Metric (Conversion factors provided) Metric < > Metric (Must know pico through Mega) Ratios Sequential Converting squared and cubed units 20
Questions and Problems Small spheres of equal mass are made of lead (density = 11.3 g/cm 3 ), silver (10.5 g/cm 3 ), and aluminum (2.70 g/cm 3 ). Which sphere has the largest diameter, and which has the smallest? Clearly explain how you came to your conclusion. The anesthetic procaine hydrochloride is used during dental surgery. It is packaged as a 10.0% solution by mass in water. The density of the solution is 1.1 g/ml. If a dentist injects 0.50 ml of the solution into a patient, what mass of procaine hydrochloride is injected? 21 Questions and Problems Suppose you build a planter box for growing vegetables in your back yard and you want to determine the amount of soil needed to fill the box to within 2.0 cm from the top. You measure the box to be 3.10 m long, 1.52 m wide and 50.0 cm high. When you go to Orchard Supply to order the soil, you discover that garden soil is sold by the cubic yard at a cost of $35.00 per cubic yard. Being highly skilled in unit conversions, you confidently calculate the amount of soil needed in cubic yards. How many cubic yards do you decide to order and what will be the total cost? In the 2012 London Olympics, Michael Phelps won a gold medal in the 100 m butterfly with a time of 51.21 seconds. What was his average speed in miles per hour? (Assume that the uncertainty in the distance is ±0.1 m.) 22
Questions and Problems Gold can be hammered into extremely thin sheets called gold leaf. If a 200-mg piece of gold (density = 19.32 g/cm 3 ) is hammered into a sheet measuring 2.4 ft 1.0 ft, what is the average thickness of the sheet in meters? How might the thickness be expressed without exponential notation, using an appropriate metric prefix? 23