Chemistry 211 General Chemistry Ebbing and Gammon John

Chemistry 211 General Chemistry Ebbing and Gammon John A. Schreifels Chemistry 211 1

WHY STUDY CHEMISTRY? • • Chemistry is important to us all. Chemistry major is reasonably likely to find a job after graduation! Career advisors suggest a broad background. Many students end up doing chemistry after graduation because of the overcrowding in their field. John A. Schreifels Chemistry 211 2

SUCCESSFUL STRATEGIES • • Memorize strategies not equations! Study a lot! Self-evaluate after quiz results. Old examinations on the WEB. John A. Schreifels Chemistry 211 3

LEARNING TOOLS • Textbook: read chapter; work lots of problems. • Other textbooks (check the library as well as used bookstores. • Tutoring center – Student Union Building II Room 2002 A • Classmates • General Chemistry WEB page: http: //osf 1. gmu. edu/~jschreif/genchem/ • Instructor John A. Schreifels Chemistry 211 4

Chapter 1: Matter and Measurement Overview: • The Study of Chemistry • Classifications of Matter • Properties of Matter • Units of Measurement • Uncertainty in Measurement • Dimensional Analysis • Basic Math Concepts (see Appendix A) John A. Schreifels Chemistry 211 5

The Study of Chemistry • Chemistry = the study of the composition, properties and transformations of matter. • Matter = physical material of the universe. • Elements = basic building blocks of all other forms of matter. • Atoms = small particles derived from one the elements. All matter can be described in terms of the interactions of atoms with each other. • Molecules (compounds) = combination of two or more atoms. Most common form for atoms. John A. Schreifels Chemistry 211 6

Classification of Matter States of Matter • Solid: Rigid Structure. Liquid: Less Rigid Structure Gas: Loose Structure. Plasma: Gaseous atoms or small molecules in an ionized state. • Substances and mixtures: • Substance: matter having a fixed composition and distinct properties. Can be either an element or a compound. – Element: a substance containing only one type of atom. E. g. Na, H 2. – Compound: a substance composed of atoms from two or more elements chemically combined. • Mixture: matter composed of two or more substances. Air composed of hydrogen, oxygen, nitrogen, etc. John A. Schreifels Chemistry 211 7

Classifications of Matter (cont’d) • Mixtures: can be separated by physical means into two or more substances. – Homogeneous: Composition constant in all parts of sample – Heterogeneous: Composition not-constant. • Mixtures separated by: – Filtration: Mixture consists of a solid and liquid; liquid separated by filtration. – Chromatography: Separates mixtures by distributing components between a mobile and stationary phase. – Distillation: Liquid mixture is boiled; components in the mixture boil off at different temperatures. John A. Schreifels Chemistry 211 8

LAW OF DEFINITE PROPORTIONS • Law of definite proportions: a pure compound always contains definite (constant) proportions by mass of the elements in the compound. E. g. Data from analysis of cyclopropane was found to contain 6. 00 g of carbon and 1. 00 g of hydrogen. If another sample was analyzed and found to contain 24. 0 g of hydrogen, how many grams of carbon would it contain? John A. Schreifels Chemistry 211 9

CONSERVATION OF MASS • Law of conservation of mass: mass is neither created or destroyed during a reaction. • The atoms form new bonds and thus are present after reaction only bound to some other atoms. • E. g. 2 H 2(g) + O 2(g) 2 H 2 O(l); 2 g of H 2 plus 16 g of O 2 produce how many grams of water? John A. Schreifels Chemistry 211 10

Properties of Matter • Physical: properties that can be measured without changing the chemical composition of the substance – E. g. melting point, smell, density. • Chemical: properties that described a substance’s reactivity. E. g. Alkali metals react to form positively charge substances; halogens form negatively charged substances • Intensive: physical or chemical property that does not depend upon the amount of the substance. Temperature, density, etc. • Extensive: physical or chemical property that depends upon the amount of material. E. g. two tanks of propane produce twice as much heat when burned as one tank. John A. Schreifels Chemistry 211 11

Units of Measurement • Mass, temperature and volume are commonly measured in the lab. • SI (System International) internationally accepted measurement system for measuring: Mass, length, temperature, etc. • Basic Units: – Mass: measured in grams; tells how much of an object there is; related to weight, which is the gravitational pull on the object. – Length: measured in meters – Temperature: measured in K or °C. Each is based upon the same reference temperatures: freezing and boiling point of water. • Conversion from one scale to the other based upon change in height of mercury between two temperatures. i. e. h = ki Ti E. g. what is the temperature in the centigrade scale if it is 44°F? John A. Schreifels Chemistry 211 12

Temperature Conversions • Conversion between temperature scales use: Where Ta, ref = Tb, a – Tf, a and Ta = Ta – Tf, a E. g. Determine the temperature in the centigrade scale corresponding to 76°F. John A. Schreifels Chemistry 211 13

Units of Measurement(cont’d) • Prefixes make it possible to express measurements in a more convenient manner. • Derived Units: • Speed: distance per elapsed time: m/s • Volume: m 3; volume often expressed in liter (L) • Density: mass per volume, g/m 3 E. g. 100 g of table salt occupies 46. 2 g. What is its density? E. g. 2: The density of liquid bromine is 3. 12 g/m. L. What is the mass of 150 m. L of bromine? John A. Schreifels Chemistry 211 14

Dimensional Analysis Often necessary to convert from one type of unit to another. • The method of dimensional analysis is used: – Multiply original number by conversion factors which change from one unit to another. – Conversion factor is the relationship between two units. • Can involve derived units. • Determine: – – – ? micrometers in 100 pm; ? ng in 55 x 10 5 kg ? kg/m 3 in 3. 45 g/m. L ? pm 2 in 6. 22 x 10 6 cm 2 ? Mm in 256000 m John A. Schreifels Chemistry 211 15

SIGNIFICANT FIGURES • All measurement made at a certain level of uncertainty and are often made several times to reduce it. • Precision: closeness of measured values. • Accuracy: closeness of measured value to the correct answer. • Significant figures in reported measurement indicate its precision. Rules: – Numbers (except zeroes) at beginning of the number sequence are significant. – Zeroes at the end of the number and to the right of the decimal point are significant. – Zeroes at the end of the number and to the left of the decimal point are not necessarily significant. • Avoid ambiguities by expressing in scientific notation: Ax 10 a where A is a number between 1 and 9. 999 and a is an integer. • Exact numbers are known to an infinite number of significant figures. John A. Schreifels Chemistry 211 16

SIG FIG AND CALCULATIONS • • Results of measurements often used to calculate some number. How many significant figures (sig fig) should the resulting number have? Rules: – – Addition and subtraction of numbers: Add all numbers; express to same number decimal places as the original number with the least number of sig. figs. Multiplication and division of numbers: Do calculation and then round to same number of sig figs as the number with least number of sig fig. John A. Schreifels Chemistry 211 17

ROUNDING • Rounding up or down is required to obtain correct number of significant figures. • Rules: If the number immediately after the last digit to be saved is: – >5 round up (add 1 to previous digit). – <5 round down (drop number). – = 5 round up if the last digit to be saved is an odd number; round down if even. E. g. 1 Round each to three significant figures: 0. 2226, 0. 22225, 5555, 554523 E. g. 2 Express the following in the correct number of significant figures: • 10000. 0 3. 14159 • 142. 7 + 0. 081 • 6. 246 + 8. 139 12. 75 • 19. 69 0. 041 + 1. 27 John A. Schreifels Chemistry 211 18

Basic Math (see appendix) • • Essential to know HS Algebra II well! Multiplication of two numbers: add the exponents for the powers of ten and multiply the two numbers together. Solve: 2. 5 x 105 2. 0 x 10 5. Powers and roots: For (Ax 10 n)m raise A to the m power and multiply n by m Solve: (2. 11 x 105)3 Log and antilog(natural and base ten): remember the definition of a logarithm: log x = z where x = 10 z. The antilog of z in the example would be x. log xy = log x + log y – – • Proportionalities: – Linear– Squared – Inverse - John A. Schreifels Chemistry 211 y = mx + b y = mx 2 + b P = k. T R = k[A]2 (second order reaction) 19

Basic Math Problems • Eg. The pressure of a gas was 1. 00 atm. at 273 K. What was the pressure at 373 K? • E. g. 2 The rate of a second order reaction was 2. 50 x 10 2 s 1 when the concentration was 0. 100 M. What would the rate be when the concentration is 0. 250 M? • E. g. 3 A gas occupied a volume of 12. 5 L when the pressure was 2. 00 atm. What volume would it occupy if the pressure increased to 3. 75 atm? John A. Schreifels Chemistry 211 20