Chapter 1 Keys to the Study of Chemistry

Transcription

1 1-1 Chapter 1 Keys to the Study of Chemistry

2 1-2 Chapter 1 : Keys to the Study of Chemistry 1.1 Some Fundamental Definitions 1.2 Chemical Arts and the Origins of Modern Chemistry 1.3 The Scientific Approach: Developing a Model 1.4 Chemical Problem Solving 1.5 Measurement in Scientific Study 1.6 Uncertainty in Measurement: Significant Figures

3 1-3 Is the study of matter, its properties, the changes that matter undergoes, and the energy associated with these changes.

4 1-4 Definitions Matter anything that has mass and volume -the stuff of the universe: books, planets, trees, professors, students Composition the types and amounts of simpler substances that make up a sample of matter Properties the characteristics that give each substance a unique identity Physical Properties those which the substance shows by itself without interacting with another substance such as color, melting point, boiling point, density Chemical Properties those which the substance shows as it interacts with, or transforms into, other substances such as flammability, corrosiveness

5 1-5 Figure 1.1 The distinction between physical and chemical change. A Physical change B Chemical change

6 1-6 Table 1.1 Some Characteristic Properties of Copper Physical Properties reddish brown, metallic luster easily shaped into sheets (malleable) and wires (ductile) good conductor of heat and electricity density = 8.95 g/cm 3 melting point = C Chemical Properties slowly forms a basic blue-green sulfate in moist air reacts with nitric acid and sulfuric acid slowly form a deep-blue solution in aqueous ammonia boiling point = C

7 1-7 Figure 1.2 The physical states of matter.

8 1-8 Energy is the capacity to do work. Potential Energy energy due to the position of the object or energy from a chemical reaction Kinetic Energy energy due to the motion of the object Potential and kinetic energy can be interconverted.

9 1-9 Energy is the capacity to do work. Figure 1.3A less stable change in potential energy EQUALS kinetic energy more stable A gravitational system. The potential energy gained when a lifted weight is converted to kinetic energy as the weight falls.

10 1-10 Energy is the capacity to do work. Figure 1.3B less stable change in potential energy EQUALS kinetic energy more stable A system of two balls attached by a spring. The potential energy gained by a stretched spring is converted to kinetic energy when the moving balls are released.

11 1-11 Energy is the capacity to do work. Figure 1.3C less stable change in potential energy EQUALS kinetic energy more stable 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.

12 1-12 Energy is the capacity to do work. Figure 1.3D less stable change in potential energy EQUALS kinetic energy more stable 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.

13 1-13 Table 1. 2 SI Base Units Physical Quantity (Dimension) mass length Unit Name kilogram meter Unit Abbreviation kg m time second s temperature kelvin K electric current ampere A amount of substance mole mol luminous intensity candela cd

14 1-14 Table 1.3 Common Decimal Prefixes Used with SI Units Prefix Prefix Symbol Word Conventional Notation Exponential Notation tera T trillion 1,000,000,000,000 1x10 12 giga G billion 1,000,000,000 1x10 9 mega M million 1,000,000 1x10 6 kilo k thousand 1,000 1x10 3 hecto h hundred 100 1x10 2 deka da ten 10 1x one 1 1x10 0 deci d tenth 0.1 1x10-1 centi c hundredth x10-2 milli m thousandth x10-3 micro μ millionth x10-6 nano n billionth x10-9 pico p trillionth x10-12 femto f quadrillionth x10-15

15 1-15 Sample Problem 1.1 Unit Conversion PROBLEM: The average speed of a nitrogen molecule in air at 25 ºC is 515 m/s. Given that 1 mi = km, convert this speed to miles per hour?

16 Figure Some interesting quantities A Length B Volume C Mass

17 1-17 Table 1.5 Densities of Some Common Substances * Substance Physical State Density (g/cm 3 ) Hydrogen Gas Oxygen Gas Grain alcohol Liquid Water Liquid Table salt Solid 2.16 Aluminum Solid 2.70 Lead Solid 11.3 Gold Solid 19.3 * At room temperature(20 º C) and normal atmospheric pressure(1atm).

18 1-18 Sample Problem 1.2 Calculating Density from Mass and Length PROBLEM: (a) Calculate the density in g/cm 3 of mercury if kg occupies a volume L (b) calculate the mass of 65.0 cm 3 of mercury.

19 1-19 Figure 1.12 The freezing and boiling points of water.

20 1-20 Temperature Scales and Interconversions Kelvin ( K ) - The Absolute temperature scale begins at absolute zero and only has positive values. Celsius ( o C)- The temperature scale used by science, formally called centigrade, most commonly used scale around the world; water freezes at 0 o C, and boils at 100 o C. Fahrenheit ( o F)- Commonly used scale in the U.S. for our weather reports; water freezes at 32 o F and boils at 212 o F. T (in K) = T (in o C) T (in o C) = T (in K) T (in o F) = 9/5 T (in o C) + 32 T (in o C) = [ T (in o F) - 32 ] 5/9

21 Rules for Determining Which Digits are Significant All digits are significant except zeros that are used only to position the decimal point. Make sure that the measured quantity has a decimal point. Start at the left of the number and move right until you reach the first nonzero digit. Count that digit and every digit to it s right as significant Zeros that end a number and lie either after or before the decimal point are significant; thus ml has four significant figures, and L has four significant figures also. Numbers such as 5300 L are assumed to only have 2 significant figures. A terminal decimal point is often used to clarify the situation, but scientific notation is the best! i.e L = L vs = L

22 1-22 Rules for Significant Figures in Answers 1. For addition and subtraction. The answer has the same number of decimal places as there are in the measurement with the fewest decimal places. Example: adding two volumes 83.5 ml ml ml = ml Example: subtracting two volumes ml ml ml = ml

23 1-23 Rules for Significant Figures in Answers 2. For multiplication and division. The number with the least certainty limits the certainty of the result. Therefore, the answer contains the same number of significant figures as there are in the measurement with the fewest significant figures. Multiply the following numbers: 9.2 cm x 6.8 cm x cm = cm 3 =23cm 3

24 1-24 Rules for Rounding Off Numbers 1. If the digit removed is equal to, or more than 5, the preceding number increases by rounds to 5.38 if three significant figures are retained and to 5.4 if two significant figures are retained. 2. If the digit removed is less than 5, the preceding number is unchanged rounds to if three significant figures are retained and to 0.24 if two significant figures are retained. 3. Be sure to carry two or more additional significant figures through a multistep calculation and round off only the final answer.

25 1-25 Issues Concerning Significant Figures Electronic Calculators be sure to correlate with the problem FIX function on some calculators Choice of Measuring Device graduated cylinder < buret pipet Exact Numbers numbers with no uncertainty 60 min = 1 hr 1000 mg = 1 g These have as many significant digits as the calculation requires.

26 1-26 Precision and Accuracy Errors in Scientific Measurements Precision - Refers to reproducibility or how close the measurements are to each other. Accuracy - Refers to how close a measurement is to the real value. Systematic error - Values that are either all higher or all lower than the actual value. Random Error - In the absence of systematic error, some values that are higher and some that are lower than the actual value.

27 1-27 Figure 1.16 Precision and accuracy in the laboratory. precise and accurate precise but not accurate

28 1-28 Figure 1.16 continued Precision and accuracy in the laboratory. random error systematic error