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CHEM 101 LECTURE NOTES Tuesday, August 30, 2001 Fall 2001 Chapter 1 Dr. Joy Heising S 572-580 I. scientific notation. – a shorthand that scientists use when dealing with very large or very small numbers. a. 108 g silver = 602,000,000,000,000,000,000,000 silver atoms. To simplify, we write 6.02 x 1023. b. 1 silver atom = 0.000 000 000 000 000 000 000 179 gram To simplify, we write 1.79 x 10-22. c. 7,400,000 = 7.4 x 106 d. 0.000246 = 2.46 x 10-4 II. significant figures. a. Exact number - no estimation is involved. Most numbers are not exact. b. Accuracy – describes how closely measured values agree with correct values. c. Precision – describes how closely individual measurements agree with each other. d. Significant figures – the digits believed to be correct by the person making the measurements. Ex. 10 mL cylinder is graduated in tenths. 4.73 mL – how many significant figures? (3) Ex. 2. population of a city is 2,678,342 people. How many ‘significant’ figures? (7) Do you believe this number? 1 CHEM 101 LECTURE NOTES Tuesday, August 30, 2001 Fall 2001 Chapter 1 Dr. Joy Heising S 572-580 III. some rules for significant figures. a. Zeroes i. ‘leading’ zeroes are never significant ii. ‘trailing’ zeroes are sometimes significant – definitely, if there is a decimal point in the number iii. zeroes in the middle of a number (4,309) are significant b. math involving significant figures. i. Addition/subtraction 3.6923 + 1.234 + 2.02 = 6.95 ii. multiplication/division 2.7832 x 1.4 = 3.9 iii. rounding – in a multi-step calculation, carry all digits (even non-significant ones) until the end of the calculation. However, if you are reporting the results of each step of the calculation, report the correct number of significant figures. 2 CHEM 101 LECTURE NOTES Tuesday, August 30, 2001 Fall 2001 Chapter 1 Dr. Joy Heising S 572-580 IV. Unit Factor method (also called dimensional analysis). a. 1 foot = 12 inches. This relationship can be converted into two unit factors. The numerator and denominator describe the same distance: ex. 1 foot 12 inches 12 inches = 1 foot = 1 9.32 yards = ? millimeters? 1 meter = 39.37 inches 9.32 yards x 3 feet 1 yard x 12 inches 1 foot x 1 meter 39.37 inches x 1000 mm 1 meter = 8520 mm 627 mL = ? gallons? 1 gallon = 0.946 L 627 mL x 1L 1000 mL x 1 gallon 0.946 L = 0.663 gallons V. Density density – mass per unit volume. D = M/V (g/mL) 742 g / 97.3 mL = 7.62 g/mL 3 CHEM 101 LECTURE NOTES Tuesday, August 30, 2001 Fall 2001 Chapter 1 Dr. Joy Heising S 572-580 VI. Heat. a. Temperature – the intensity of heat in a body, that is, how hot or how cold the body is. Three different temperature scales represent same temp: i.Fahrenheit – bp H2O 212ºF ii.Celsius – bp H2O 100ºC 5 x (ºF – 32)/9 iii.Kelvin - bp H2O 373 K (Celsius + 273) b. the SI unit of heat and energy is the joule (J). [another useful unit is the calorie – the amount of heat required to raise the temperature of 1 g of H2O from 14.5ºC to 15.5ºC. 1 cal = 4.184 J.] c. specific heat – the amount of heat (usually expressed in joules) required to raise the temperature of one gram of the substance one degree Celsius with no change in state. d. Heat capacity – the amount of heat required to raise the temperature of one mole of a substance one degree Celsius with no change in state. e. Specific heats of a few common substances Substance specific heat Ice Liquid water Steam 2.02 J/gºC 4.18 J/gºC 2.03 J/gºC Ex. calculate the amount of heat required to raise the temperature of 200 g H2O from 10.0ºC to 55.0ºC. 55.0ºC – 10.0ºC = 200.0g x 45.0ºC x 45.0ºC 4.18 J gºC = 3.76 x 104 J 4 CHEM 101 LECTURE NOTES Tuesday, August 30, 2001 Fall 2001 Chapter 2 Dr. Joy Heising S 572-580 Chapter 2 Stoichiometry–describes the quantitative relationships among elements 1. in compounds (composition stoichiometry) 2. in chemical changes (reaction stoichiometry) Dalton’s Atomic Theory – summarized experimental observations and interpretations in the nature of atoms: 1. an element is composed of extremely small, indivisible particles called atoms 2. Atoms cannot be created, destroyed, or transformed into atoms of another element 3. compounds are formed when atoms of different elements combine with each other in small, wholenumber ratios. 4. the relative numbers and kinds of atoms are constant in a given compound. Atoms – (review, chapter 1) Molecules - (review, chapter 1) 5 CHEM 101 LECTURE NOTES Tuesday, August 30, 2001 Fall 2001 Chapter 2 Dr. Joy Heising S 572-580 Chemical formula – the chemical composition of a substance. Some examples of chemical formulas for elements: Monoatomic elements – sodium (Na), copper (Cu), barium (Ba) Diatomic elements – O2, N2, Br2 More complex molecules – S8, P 4 For compounds, the formula indicates 1. elements present 2. the ratio in which the atoms of the elements occur Some examples: HCl – 1 H atom, 1 Cl atom H2O -2 H atoms, 1 O atom NH3 -1 N atoms, 3 H atoms C3H8 – 3 C atoms, 8 H atoms 1 chemical formula 1 chemical formula 1 chemical formula 1 chemical formula (Fig. 2.6) law of constant composition – (review, chapter 1) 6 CHEM 101 LECTURE NOTES Tuesday, August 30, 2001 Fall 2001 Chapter 2 Dr. Joy Heising S 572-580 molecular compounds: (Table 2.2) H2O2 CH3CH2OH Methane NH3 CO2 benzene hydrogen peroxide methanol CH4 ammonia carbon dioxide C6 H6 Ion – an atom or group of atoms that carries an electrical charge Anion – negatively charged ions (Cl-) Cation – positively charged ions (Na +) Ionic compounds – extended array of ions in which the total positive and negative charges are equal (Fig 2.7) Formula unit – the simplest whole number ratio of ions in the compound (NaCl) Is PO43- a molecule? No. polyatomic anion. Compound names and formulas (table 2-3) NaCl Sulfuric acid Molecule sodium chloride H2 SO4 ? NH4NO3 Na2SO4 ion 7 ammonium nitrate sodium sulfate CHEM 101 LECTURE NOTES Tuesday, August 30, 2001 Fall 2001 Chapter 2 Dr. Joy Heising S 572-580 Atomic weight – the relative masses of atoms of different elements that are proportional to the actual masses of atoms. amu – units of atomic weight. (atomic mass unit) 1 amu – 1/12 the mass of carbon-12. H 1 amu Mg 24.3 amu N Ag 14 amu 107.8 amu amu is not terribly practical ‘in real life’ for experimental chemists, as the quantities are too small (most scientists aren’t weighing individual atoms to do their experiments). As a result, chemists use another larger unit called the mole. Mole – 6.022 x 1023 atoms, molecules or formula units Avogadro’s number **one mole of atoms of an element has a mass in grams numerically equal to the atomic weight of the element** H 6.022 x 1023 atoms N 6.022 x 1023 atoms Mg 6.022 x 1023 atoms H2 5.022 x 1023 molecules, 1.204 x 1024 atoms How many atoms are contained in 1.67 moles of Mg? How many grams does 1.67 moles of Mg weigh? = 1.00 x 1024 atoms = 40.6 grams 8