Chemistry and Chemical Reactivity Chapter 1 Notes PDF

1 Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M...

1 Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 1 Matter and Measurement Lectures written by John Kotz ©2006 © 2006 Brooks/Cole Brooks/Cole - - Thomson Thomson  Welcome to the 2 World of Chemistry © 2006 Brooks/Cole - Thomson  3 Chemistry & Matter We can explore the MACROSCOPIC world — what we can see — to understand the PARTICULATE worlds we cannot see. We write SYMBOLS to describe these worlds. © 2006 Brooks/Cole - Thomson  4 A Chemist’s View of Matter © 2006 Brooks/Cole - Thomson  5 A Chemist’s View Macroscopic 2 H2(g) + O2 (g) --> 2 H2O(g) Particulate Symbolic © 2006 Brooks/Cole - Thomson  6 Kinetic Nature of Matter Matter consists of atoms and molecules in motion. © 2006 Brooks/Cole - Thomson  7 STATES OF MATTER © 2006 Brooks/Cole - Thomson  8 STATES OF MATTER SOLIDS — have rigid shape, fixed volume. External shape can reflect the atomic and molecular arrangement. – Reasonably well understood. LIQUIDS — have no fixed shape and may not fill a container completely. – Not well understood. GASES — expand to fill their container. – Good theoretical understanding. © 2006 Brooks/Cole - Thomson  9 Classifying Matter © 2006 Brooks/Cole - Thomson  10 Chicken Noodle Soup Blood Salt Water © 2006 Brooks/Cole - Thomson Fig. 1.4, p.16  11 The Language of Chemistry CHEMICAL ELEMENTS - – pure substances that cannot be decomposed by ordinary means to other substances. Aluminum Bromine Sodium © 2006 Brooks/Cole - Thomson  12 The Language of Chemistry The elements, their names, and symbols are given on the PERIODIC TABLE How many elements are there? 116 elements © 2006 Brooks/Cole - Thomson  13 The Periodic Table Dmitri Mendeleev (1834 - 1907) © 2006 Brooks/Cole - Thomson  14 Glenn Seaborg (1912-1999 ) Discovered 8 new elements. Only living person for whom an element was named. © 2006 Brooks/Cole - Thomson  15 An atom is the smallest particle of an element that has the chemical properties of the element. Copper atoms on silica surface. See CD- ROM Screen 1.4 Distance across = 1.8 nanometer (1.8 x 10-9 m) © 2006 Brooks/Cole - Thomson  The Atom 16 An atom consists of a nucleus –(of protons and neutrons) electrons in space about the nucleus. Electron cloud Nucleus © 2006 Brooks/Cole - Thomson  CHEMICAL COMPOUNDS are 17 composed of atoms and so can be decomposed to those atoms. The red compound is composed of nickel (Ni) (silver) carbon (C) (black) hydrogen (H) (white) oxygen (O) (red) nitrogen (N) (blue) © 2006 Brooks/Cole - Thomson  18 A MOLECULE is the smallest unit of a compound that retains the chemical characteristics of the compound. Composition of molecules is given by a MOLECULAR FORMULA H2O C8H10N4O2 - caffeine © 2006 Brooks/Cole - Thomson  19 Elements form Compounds © 2006 Brooks/Cole - Thomson  20 The Nature of Matter Gold Mercury Chemists are interested in the nature of matter and how this is related to its atoms and molecules. © 2006 Brooks/Cole - Thomson  21 Physical Properties What are some physical properties? color melting and boiling point odor © 2006 Brooks/Cole - Thomson  22 Graphite — layer structure of carbon atoms reflects physical properties. © 2006 Brooks/Cole - Thomson  23 Physical Changes Some physical changes would be boiling of a liquid melting of a solid dissolving a solid in a liquid to give a homogeneous mixture — a SOLUTION. © 2006 Brooks/Cole - Thomson  24 DENSITY - an important and useful physical property Density = mass (g) volume (cm3) Platinum Mercury Aluminum 13.6 g/cm3 21.5 g/cm3 2.7 g/cm3 © 2006 Brooks/Cole - Thomson  25 Relative Densities of the Elements © 2006 Brooks/Cole - Thomson  26 Problem A piece of copper has a mass of 57.54 g. It is 9.36 cm long, 7.23 cm wide, and 0.95 mm thick. Calculate density (g/cm3). Density = mass (g) volume (cm3) © 2006 Brooks/Cole - Thomson  27 SOLUTION 1. Get dimensions in common units. 1cm 0.95 mm = 0.095 cm 10 mm 2. Calculate volume in cubic centimeters. (9.36 cm)(7.23 cm)(0.095 cm) = 6.4 cm3 Note only 2 significant figures in the answer! 3. Calculate the density. 57.54 g = 9.0 g / cm3 6.4 cm3 © 2006 Brooks/Cole - Thomson  28 DENSITY Density is an INTENSIVE property of matter. Styrofoam Brick –does NOT depend on quantity of matter. –temperature Contrast with EXTENSIVE –depends on quantity of matter. –mass and volume. © 2006 Brooks/Cole - Thomson  29 PROBLEM: Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in grams? In pounds? Solve the problem using DIMENSIONAL ANALYSIS. © 2006 Brooks/Cole - Thomson  30 PROBLEM: Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg? 1. Convert volume to mass ( 95 cm3 )(13.6 g/cm3) = 1.3 x 103 g 2. Convert mass (g) to mass (lb) 1 lb 1.3 x 103 g = 2.8 lb 454 g © 2006 Brooks/Cole - Thomson  Chemical Properties and 31 Chemical Change Chemical change or chemical reaction — transformation of one or more atoms or molecules into one or more different molecules. © 2006 Brooks/Cole - Thomson  32 Types of Observations and Measurements We make QUALITATIVE observations of reactions — changes in color and physical state. We also make QUANTITATIVE MEASUREMENTS, which involve numbers. Use SI units — based on the metric system © 2006 Brooks/Cole - Thomson  33 UNITS OF MEASUREMENT Use SI units — based on the metric system Length Meter, m Mass Kilogram, kg Time Seconds, s Temperature Celsius degrees, ˚C kelvins, K © 2006 Brooks/Cole - Thomson  34 Units of Length 1 kilometer (km) = ? meters (m) 1 meter (m) = ? centimeters (cm) 1 centimeter (cm) = ? millimeter (mm) 1 nanometer (nm) = 1.0 x 10-9 meter O—H distance = 9.58 x 10-11 m 9.58 x 10-9 cm 0.0958 nm © 2006 Brooks/Cole - Thomson  35 Temperature Scales Fahrenheit Celsius Kelvin Anders Celsius 1701-1744 Lord Kelvin (William Thomson) 1824-1907 © 2006 Brooks/Cole - Thomson  36 Temperature Scales Fahrenheit Celsius Kelvin Boiling point of water 212 ˚F 100 ˚C 373 K 180˚F 100˚C 100 K Freezing point 32 ˚F 0 ˚C 273 K of water Notice that 1 kelvin degree = 1 degree Celsius © 2006 Brooks/Cole - Thomson  37 Temperature Scales 100 oF 38 oC 311 K oF oC K © 2006 Brooks/Cole - Thomson  38 Calculations Using Temperature Generally require temp’s in kelvins T (K) = t (˚C) + 273.15 Body temp = 37 ˚C + 273 = 310 K Liquid nitrogen = -196 ˚C + 273 = 77 K © 2006 Brooks/Cole - Thomson  39 Precision, Accuracy, Error Precision – how well measurements agree – Measure by calculating the range Accuracy – How close measurements are to the accepted value. – Measure by calculating Percent Error » % error = (Your value – Accepted value) x 100 Accepted Value Error – Experimental—error you cause – Laboratory—error inherent in an experiment © 2006 Brooks/Cole - Thomson  40 © 2006 Brooks/Cole - Thomson Fig. 1.19a, p.33  41 © 2006 Brooks/Cole - Thomson Fig. 1.19b, p.33  42 © 2006 Brooks/Cole - Thomson Fig. 1.19c, p.33  43 Significant Figures numbers which reflect accuracy in measurement Number Rule Example Non zero All are significant 1.23 : 3 1.23 x 10-3 : 3 Leading zeros ALL zeros before nonzero 0.000123 : 3 numbers are NOT significant Trailing Zeros # w/ decimal: all zeros 123000 : 3 following nonzero numbers 12300.000 : 8 are significant # w/o decimal: all zeros following nonzero numbers are NOT significant © 2006 Brooks/Cole - Thomson