1 / 26

Lecture 1 Introduction: Ch 1.1-1.6

Lecture 1 Introduction: Ch 1.1-1.6. Dr. Harris Suggested HW: ( Ch 1) 5, 7, 8, 11, 13, 15, 16, 27. What is Chemistry?. Chemistry is the study of properties of substances and how they react Chemical substances are composed of matter

iona-norman
Download Presentation

Lecture 1 Introduction: Ch 1.1-1.6

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Lecture 1Introduction: Ch 1.1-1.6 Dr. Harris Suggested HW: (Ch 1) 5, 7, 8, 11, 13, 15, 16, 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

  3. Elements, Molecules and Compounds • Any pure substance that can not be broken down into simpler substances is an element. • Millions of different materials in the world, all comprised of some combination of only 118 elements • Similar to how the alphabet combines 26 letters to yield hundreds of thousands of words, elements bond in unique arrangements to give different molecules • Molecules agglomerate to yield compounds

  4. Compounds O H H O (Au)n

  5. 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

  6. 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? Qualitative data • What color is the product? • Is it solid, liquid, gas? • How does it smell? • What is it?

  7. Quantitative Measurements and Units • Internationally accepted system of measurements is called the SIunit system • Quantitative measurements are represented by a: NUMBER and a UNIT • A unit is a standard against which a physical quantity is compared physical quantity • Temperature is measured in Co, Ko,orFo • Currency is measured in $USD • Distance is measured in meters, miles, ft, etc. • Time is reported in seconds, minutes, hr, etc.

  8. SI Unit System: The units of Chemistry

  9. Scientific Notation & Exponent Review Scientific notation indicates a factor (F) multiplied by a power (n) of 10 F x 10n (1 < F < 10) scientific • 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. standard

  10. Scientific Notation & Exponent Review • Convert to standard notation • 3.4912 x 104 • 8.971 x 10-3 • 6.50 x 100 • Convert to scientific notation • 15 • 125.3 • 0.003003 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. All values less than 1 have negative exponents. 7.141 x 10-2 = .07141 3.867 x 10-7 = .0000003867

  11. Converting to Proper 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, n • For factors less than 1, you must “pull out” powers of 10 from n, then add those powers to Funtil 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 103= 3.840 x 10-1 Converted to proper scientific notation

  12. 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) = 3.0 x (10 10-(-6)) = 3.3 x 1016

  13. Unit Prefixes • Prefixes indicate powers of 10 • ex. k= 103; 5 kg = 5 x (103)g

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

  15. Derived SI Units (Volume) • Many measured properties have units that are combinations of the fundamental 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 (length)3 • 1 L is equal to the volume of fluid that a cube which is 10 cm on each side can hold 10 cm V = (10 cm) x (10 cm) x (10 cm) = (10 cm)3= 1000 cm3 10 cm 1 L = 1000 cm3 10 cm 1000 mL = 1000 cm3 mL = cm3

  16. Derived SI Units (Density) Does a 20-gallon trash filled with bricks weigh the same as one filled with feathers? NO! All matter has mass, and must therefore occupy space. Densitycorrelates the mass of a substance to the volume of space it occupies. Density = mass per unit volume (mass/volume). Different materials have different densities. THE DENSITY OF WATER IS

  17. KNOW THIS !!! THE DENSITY OF WATER IS

  18. 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!!

  19. Derived Units (Energy) • What is Energy? • Energy is defined as the capacity to perform “work” • How do we define work? • Work is defined as the action of applying a force acting over some distance • Ex. Pushing an object along a rough surface • In SI units, we use the unit Joule(J) to represent energy.

  20. Kinetic Energy • In this class, you will perform several calculations involving kinetic energy. • Energy of motion (e.g. a moving car). An object with mass m, moving at a velocity V (meters/sec) has kinetic energy equal to:

  21. Temperature • Temperature: a measure of the tendency of a substance to lose or absorb heat. Temperature and heat are not the same. • 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 • 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

  22. 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

  23. Intensive vs. Extensive Properties • Density and temperature are examples of intensive properties; meaning that they do NOT depend on the amount of substance. • Ex. 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 of substance. • 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.

  24. Accuracy and Precision • Accuracy defines how close to the correct answer you are. Precision defines how repeatable your result is. Ideally, data should be both accurate and precise, but it may be one or the other, or neither. Precise, but not accurate. Did not reach the target, but result was reproduced. Accurate, but not precise. Reached the target, but could not reproduce the result. Accurate and precise. Reached the target and the data was reproduced.

  25. 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 • - A certain brand of thermometer is considered to be accurate if the %E is < 0.8%. The thermometer is tested using water (BP = 100oC). You bring a pot of distilled water to a boil and measure the temperature 5 times. The thermometer reads: 100.6o, 100.4o, 99.4o, 101.0o, and 100.4o. Is it accurate?

More Related