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Lecture 1 Introduction: Ch 1.1-1.6

Lecture 1 Introduction: Ch 1.1-1.6. Dr. Harris 8/22/12 HW Problems: Ch 1: 5, 7, 8, 11, 15, 27. What is Chemistry?. Chemistry is the study of properties of substances and how they react Chemical substances are composed of matter

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Lecture 1 Introduction: Ch 1.1-1.6

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  1. Lecture 1Introduction: Ch 1.1-1.6 Dr. Harris 8/22/12 HW Problems: Ch 1: 5, 7, 8, 11, 15, 27

  2. What is Chemistry? • Chemistry is the study of properties of substances and how they react • Chemical substances are composed of matter • Matter is the physical material of the universe; anything with mass that occupies space is matter • Matter can take many forms. • Most matter is formed by unique arrangements of elementary substances called elements • Elements are puresubstancesthat and consist of atoms, the building blocks of matter.

  3. Compounds O H H O (Au)n

  4. Scientific Method • In chemistry, the scientific method is used to investigate scientific phenomena & acquire new knowledge • Based on gathering empirical evidence to support or refute a hypothesis • Empirical evidence is either quantitative or qualitative • Quantitative data is numerical. Quantitative results are measured • Qualitative data is NOT numerical, but consists of observations and descriptions

  5. Quantativevs Qualitative Science A + B -----> C Quantitative data • How much C is formed? • How efficient is the reaction? (How much A/B is lost in the process?) • How fast does C form? • How efficient is the reaction Qualitative data • What color is the product? • Is it solid, liquid, gas? • How does it smell? • What is it?

  6. Steps in the Scientific Method What do I want to know? What is known? Has this been done? Similar Gather basic information. I believe that…. Develop a scientific law

  7. Laws and Theories • Once you complete the experiment, determine if the data is consistent. • Look at trends in the data. If the results follow a consistent trend, then you can form a scientific law • Laws interpret data, but do not explain it • Theories are created to explain data

  8. Difference Between Law and Theory • Example • Question: I am curious to see how a student’s grade is related to his/her attendance • Background research: Did someone already do this? I must compare similar students (similar high school background, similar SAT scores, etc.). I would gather info on a pool of students • Hypothesis: Students that attend class regularly will out-perform those students who don’t

  9. Difference Between Law and Theory • Experiment: Monitor the performance of X students with similar credentials in general chemistry. Compare students who: • attend every class • attend 50% • attend 25% • Result: Students who come to class more, on average, earn better grades (This is my law)

  10. Difference Between Law and Theory • So, now I have a law, but I haven’t developed a theory to explain it. • My theoryis: Students who attend class regularly are better prepared for quizzes and exams than those students who don’t. • Theories are not universal truths. They can be challenged, changed, tweaked, or rejected completely. • What is another possible theory for this data?

  11. Units • All quantitative measurements must have units. • A unit is a standard measurement for a physical quantity • units of temperature are Fo, Co; • distance can be reported in units of miles, ft, m; • time is reported in seconds, minutes, hours, etc.;

  12. SI Unit System: The units of Chemistry

  13. Mass and Weight are NOT the same • Mass defines an object’s resistance to movement, and is determined by the composition of the object. Weight is a force. Because of gravity, your body exerts a force on the ground. The unit of force is Joules (J). • An object with the same mass can have different weights. • For example, if you were to travel to the moon, you would weigh only 1/6th of what you weigh on earth • But your mass would not change because you are composed of the same amount of matter

  14. Unit Prefixes • Prefixes indicate powers of 10

  15. Scientific Notation & Exponent Review Scientific notation indicates a factor (F) multiplied by a power (n) of 10 F x 10n (1 < F < 10) • Important: All integers end with a decimal point, even though it is not commonly written (1  1. ) • If no factor is shown, assume there is a 1. in front of powers of 10: 102 = 1. x 102 10-7 = 1. x 10-7 • For every positive power of 10, shift the decimal 1 place to the right, add a zero for each place 102 = 1. x 102 = 100. 105 = 1. x 105 = 100000.

  16. Scientific Notation & Exponent Review • For all non integers, simply shift the decimal 2.5 x 105 = 250000. 1.8773 x 108 = 187730000. • For negative exponents, shift the decimal left 7.141 x 10-2 = .07141 3.867 x 10-7 = .0000003867 • Examples: Convert the following to standard notation • 3.4912 x 102 • 9.1001 x 106 • 8.971 x 10-3 • 6.50 x 100

  17. Converting to Standard Scientific Notation • Remember : 1<F < 10 • For factors greater than 10, you must “pull out” powers of 10 from F until an allowed value of F is obtained, then add those powers to the exponential term • For factors less than 1, you must “pull out” powers of 10 from the exponential term, then add those powers to F until an allowed value of F is obtained. 11.60 x 104 = 1.160x 105 3217.4 x 102 = 3.2174 x 105 11834.1 x 10-7 = 1.18341 x 10-3 Converted to proper scientific notation 0.000185 x 104 = 1.85 x 100 = 1.85 0.000007 x 1010 = 7 x 104 0.0003840 x 106 = 3.840 x 102 Converted to proper scientific notation

  18. Multiplying and Dividing Exponents • When multiplying powers of 10, the product is the sum of the powers • 102 x 105 = 10 2+5 = 107 • (2.5 x 103) x (4 x 10-6) = (2.5 x 4) x (103+(-6)) = 10 x 10-3 = 1 x 10-2 • When dividing powers of 10, subtract 102 / 105 = 10 (2-5) = 10-3 (6.6 x 1010)/ (2.2 x 10-6) = 3.3 x (10 10-(-6)) = 3.3 x 1016

  19. Group Problems Convert the following values to grams in proper scientific notation. • 421.4 kg • 110.1 x 10-6 mg • 18.9 Mg • 481 µg Express each value in standard form.

  20. Derived SI Units • Volume: defines the quantity of space an object contains or occupies; or the amount of fluid a container can hold • expressed in units of Liters (L) or length3 • 1 L is equal to the volume of fluid that a cube which is 10 cm on each side can hold V = (10 cm) x (10 cm) x (10 cm) = (10 cm)3= 1000 cm3 10 cm 10 cm 1 L = 1000 cm3 10 cm 1000 mL = 1000 cm3 mL = cm3

  21. Derived SI Units • Density: mass per unit volume (mass/volume). Different materials have different densities. Example: Does a 20-gallon trash filled with bricks weigh the same as one filled with feathers? Of course not!! Volume of trash can= 20 gallons THE DENSITY OF WATER IS

  22. Group Problem • A cubic container that is 100 cm on each side is filled with water. • What is the volume of water in the cube in cm3? • What is the volume in mL? • What is the volume in L? • What is the mass of water in the container in g? Give answers in scientific notation!!

  23. Temperature • Temperature: a measure of the tendency of a substance to lose or absorb heat. • Heat always flows from bodies of higher temperature to those of lower temperature • The stove top is ‘hot’ because the surface is at a much higher temperature than your hand, so heat flows rapidly from the stove to your hand, causing burns • Ice feels ‘cold’ because it is at a lower temperature than your body, so heat flows from your body to the ice, causing it to melt

  24. Temperature • When performing calculations in chemistry, temperature must always be converted to Kelvin (oK) units. The lowest possible temperature that can ever be reached is 0oK, or absolute zero. At this temperature, all molecular motion stops. • To convert temperatures to the Kelvin scale: oK : oC + 273.15  If given oC, convert oC to oK oK : (oF + 459.67)  If given oF, convert oF to oK

  25. Intensive vs. Extensive Properties • Density and temperature are examples of intensive properties; meaning that they do NOT depend on the amount. • The density of water is 1g/cm3 no matter how much water you have. • Mass and volume, however, are examples of extensive properties, and do depend on the amount. • If you double the amount of a substance, you double its mass, and it takes up twice as much space, so its volume is doubled as well.

  26. Accuracy and Precision • Accuracy defines how close to the correct answer you are. Precision defines how reproducible (repeatable) your result is. Accurate and precise. Reached the target and the data was reproduced. Accurate, but not precise. Reached the target, but could not reproduce the result. Precise, but not accurate. Did not reach the target, but result was reproduced.

  27. Measuring Accuracy: Percentage Error • Accuracy is calculated by percentage error (%E) • We take the absolute value because you can’t have negative error. • GROUP PROBLEM • - The boiling temperature of water is known to be exactly 100oC. You bring a pot of distilled water to a boil and measure the temperature 4 times. The thermometer reads: 100.6o, 100.4o, 99.4o, 101.0o. What is the percentage error?

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