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Chemistry 211

Chemistry 211. General Chemistry Ebbing and Gammon. 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.

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Chemistry 211

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  1. Chemistry 211 General Chemistry Ebbing and Gammon

  2. 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.

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

  4. 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 2002A • Classmates • General Chemistry WEB page: http://osf1.gmu.edu/~jschreif/genchem/ • Instructor

  5. 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)

  6. 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.

  7. Classification of Matter States of Matter • Solid: Rigid StructureLiquid:Less Rigid Structure Gas:Loose StructurePlasma: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, H2. • 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.

  8. 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.

  9. 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?

  10. 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. 2H2(g) + O2(g)  2H2O(l); 2 g of H2 plus 16 g of O2 produce how many grams of water?

  11. 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.

  12. 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?

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

  14. 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: m3; volume often expressed in liter (L) • Density: mass per volume, g/m3 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/mL. What is the mass of 150 mL of bromine?

  15. 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 55x105 kg • ? kg/m3 in 3.45 g/mL • ? pm2 in 6.22x106 cm2 • ? Mm in 256000 m

  16. 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: Ax10a 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.

  17. 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.

  18. 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

  19. 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: • 2.5x1052.0x105 . Powers and roots: For (Ax10n)m raise A to the m power and multiply n by m • Solve: (2.11x105)3 • Log and antilog(natural and base ten):remember the definition of a logarithm: log x = z where x = 10z. The antilog of z in the example would be x . • log xy = log x + log y • Proportionalities: • Linear- y = mx + b P = kT • Squared - y = mx2 + b R = k[A]2 (second order reaction) • Inverse -

  20. 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 2.50x10-2 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?

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