1 / 55

The Chemistry of Everything Kimberley Waldron

The Chemistry of Everything Kimberley Waldron. Chapter 1 Everything An overview of the composition of matter and the way scientists study it. Richard Jarman, College of DuPage. Chapter Topics. The periodic table, classes of subatomic particles, electron density.

reba
Download Presentation

The Chemistry of Everything Kimberley Waldron

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. The Chemistry of EverythingKimberley Waldron Chapter 1 Everything An overview of the composition of matter and the way scientists study it Richard Jarman, College of DuPage (C)2007 Prentice Hall

  2. Chapter Topics • The periodic table, classes of subatomic particles, electron density. • Elements, compounds, molecules, chemical reactions, balanced chemical equations. • Separation of substances in a mixture, homogeneous and heterogeneous mixtures. • Sizes of atoms, SI units, conversion factors, dimensional analysis. • Common units of measure, measurement of temperature • The scientific method, testing hypotheses, scientific models. (C)2007 Prentice Hall

  3. Science and Change • The twentieth century saw a revolution in technological advances: • Automobiles • Electricity • Computers • Everything has gotten much faster, even if we think things are still slow. (C)2007 Prentice Hall

  4. Technical Advance Builds on Discovery • Technical advances do not just happen. • They depend on scientific discovery. • 19th century was one of discovery: • Electrons • X-rays • Radioactivity (C)2007 Prentice Hall

  5. Changing Our Views of Matter • X-rays and radioactivity changed ideas about atoms. • Atoms can absorb and emit radiation. • X-rays gave us pictures of molecules. (C)2007 Prentice Hall

  6. Radiation for Good and Evil • X-rays gave insights into the nature of penicillin and DNA. • Radioactivity provided the means to split atoms and destroy the world. (C)2007 Prentice Hall

  7. DNA • X-rays provided a picture of the structure of DNA. • The DNA double helix explains transfer of genetic information. • We can modify the helix by genetic engineering. (C)2007 Prentice Hall

  8. For Good and Evil • Science has been used for great good… • …and great evil. • What is good for some is not always good for others. • Arguments for or against are often complex. • Science education helps us make informed decisions. (C)2007 Prentice Hall

  9. Atoms – The Basis of Matter • An atom is the smallest unit of matter. • All matter is made of atoms. • There are different types of atoms. • Each type is an element. (C)2007 Prentice Hall

  10. Elements and the Periodic Table • The elements (about 114) are arranged in the periodic table. • Each row is a period. • Each column is a group. (C)2007 Prentice Hall

  11. Atomic Structure • Atoms contain nuclei. • The nucleus contains protons and neutrons. • The nucleus is very dense and very tiny. (C)2007 Prentice Hall

  12. Atomic Facts and Figures • Atomic number: All atoms of the same element contain the same number of protons. • Each element has a unique atomic number. • Neutral atoms contain same number of electrons and protons. • Electrons live outside the nucleus. (C)2007 Prentice Hall

  13. Particles, Clouds and the Illusive Electron • Electrons were thought to be particles orbiting the nucleus like moons around a planet. • The modern view is that they are “clouds” smeared out in space. (C)2007 Prentice Hall

  14. Chemical Reactions • Any chemical reaction involves electrons. • Electrons are given, taken away or shared. • Chemical reactions involve compounds – the combination of atoms. (C)2007 Prentice Hall

  15. Elements and Compounds • Elements contain only one type of atom: • Neon, Gold, Fluorine • Compounds contain more than one type of atom in simple fixed ratios: • Carbon dioxide, water (C)2007 Prentice Hall

  16. Molecules and Bonds • Molecules are the basic blocks of substances. • Molecules contain atoms bonded together. • Chemical bond is the interaction of electrons to join atoms together. • Without bonds all matter would be a gas. (C)2007 Prentice Hall

  17. Big Bang Theory • Nitroglycerin is a molecule. • It is unstable – decomposition of nitroglycerin produces an explosion. • Dynamite (invented by Alfred Nobel) is a mixture of nitroglycerin with chalky material to make it more stable. (C)2007 Prentice Hall

  18. Describing Reactions • Chemical equations are ways to describe what happens in a chemical reaction. • Symbols are used for the substances. • Reactants are the starting materials. • Products are obtained in the reaction. • Arrows show the direction. Reactants Products (C)2007 Prentice Hall

  19. Anatomizing the Reaction • Existing chemical bonds are broken. • New chemical bonds are made. • The number of atoms remains the same. • They just change places. (C)2007 Prentice Hall

  20. The Chemical Formula • Subscripts tell us how many atoms of each element are in the molecule. • Each molecule of TNT contains: • 7 atoms of carbon • 5 atoms of hydrogen • 3 atoms of nitrogen • 6 atoms of oxygen (C)2007 Prentice Hall

  21. Keeping in Balance (C)2007 Prentice Hall

  22. Models and Molecules: Visualizing the Unseen • We use models to show the atoms in molecules – balls and sticks. • Computers use molecular models to design new drugs. (C)2007 Prentice Hall

  23. Clean or Pure? • A pure substance is one that contains only one element or compound. • Something that is “clean’ is probably not strictly pure – a mountain stream contains salt. (C)2007 Prentice Hall

  24. Mixtures are Impure • Nature tends to mix things – more than one pure substance mixed together. • Homogeneous mixture is uniform on a molecular level (a solution). • Heterogeneous mixture is non-uniform – a bottle of Dulcolax. (C)2007 Prentice Hall

  25. YES NO YES NO YES NO (C)2007 Prentice Hall

  26. Let’s Get Physical • Components of a mixture can be separated (purified) by physical means. • Physical is altering a substance without any changes to the chemical composition. • A saline solution (mixture) is still salt and water. • Water can be removed by evaporation. (C)2007 Prentice Hall

  27. Chemical Changes • Chemical change involves breaking down a substance into other substances. • Boiling (physical) converts water into steam. • Electrolysis (chemical) converts water into hydrogen and oxygen – new substances: (C)2007 Prentice Hall

  28. All Matter is Ultimately Made from Elements (C)2007 Prentice Hall

  29. Questions of Science • Scientific discoveries create complex questions: • Discovery of the atom led to potential for source of power combined with radioactive waste. • Should scientists be free to pursue any research? • Should government control scientific progress? • What is the progression of science? • Research normally follows a logical progression. • Interest focuses on problems that impact humanity. (C)2007 Prentice Hall

  30. The Importance of Publication • Results of chemical research are submitted for publication. • Publication involves peer review by other scientists. • Publication and peer review establish the legitimacy of scientific progress. (C)2007 Prentice Hall

  31. What is Science? • Science conforms to rules and conventions. • Without rules it is meaningless. • Hypothesis is the basis of the scientific method. • A proposal about how something behaves which is testable by observation. • Hypotheses that cannot be tested are not scientific (the supernatural). (C)2007 Prentice Hall

  32. SCIENTIFIC METHOD 1. FACT : An observable event; indisputable evidence which does not explain but simply is. 2. HYPOTHESIS: A guess to try to explain an observation. 3. EXPERIMENT: A systematic exploration of an observation or concept. 4. THEORY: An explanation of the facts; it can be proven by experiment and it confirms an hypothesis. 5. LAW: A theory which has undergone rigorous experimentation and no contradiction can be found. Note: MODEL: A visual or mathematical device or method used to demonstrate a theory or concept.

  33. Theory and the Method • A theory is an explanation of the observed phenomenon using experimental data. (C)2007 Prentice Hall

  34. Which Side is Your Bread Buttered? • Scientific hypothesis must be subject to experimental test. Hypothesis: The buttered side of bread always lands face down on the floor. • Test by dropping bread on the floor many times. (C)2007 Prentice Hall

  35. Results and Theories • Buttered side lands face down 90 % of the time. • Formulate a theory: the height of the table determines what side the bread lands. • Experiments test a theory but never prove it. • The next experiments may falsify it. (C)2007 Prentice Hall

  36. Numbers and Units in Science • Scientists measure very large numbers: • 160,000,000,000,000,000,000,000 atoms in a coin. • And very small numbers: • The diameter of the atom is 0.000 000 000 153 meter. (C)2007 Prentice Hall

  37. MEASUREMENTSScientific Notation Many measurements in science involve either very large numbers or very small numbers (#). Scientific notation is one method for communicating these types of numbers with minimal writing. GENERIC FORMAT: # . # #… x 10# A negative exponent represents a number less than 1 and a positive exponent represents a number greater than 1. 6.75 x 10-3 is the same as 0.00675 6.75 x 103 is the same as 6750

  38. MEASUREMENTSScientific Notation 5.289003 x 1011 3.400 x 10-12 Give the following in scientific notation (or write it out) with the appropriate significant figures. 1. 528900300000 = 2. 0.000000000003400 = 3. 0.23 = 4. 5.678 x 10-7 = 5. 9.8 x 104 = 2.3 x 10-1 0.0000005678 98000

  39. GROUP STUDY PROBLEM #1 Show all work for the following questions on the back page. Always give the correct significant figures. 1. What is the scientific method? 2. Explain why tasting chemicals during an experiment is not a safe practice? 3. Express each of the following numbers in scientific notation. a) 0.0404 _______ b) 0.0081 _______ c) 40.0 _______ d) 2900000 _______ e) 0.0000055 _______ f) 40300 _______ 4. (452 x 6.2) / 14.30 = ______________________ 5. 98.78564 - 97.9299 = ________________ 6. 6.022 x 1023 / 1.9 x 10-19 = ________________ 7. ( 0.3 - 0.09) / (4.3 x 10-3 + 7.232 x 10-4) = ____________ 8. [(2.4 x 1012) (5.78 x 10-31)]3 / (2.965 x 1014)1/2 = __________

  40. MEASUREMENTSSignificant Figures All nonzero numbers are significant figures. 2. Zero’s follow the rules below. Zero’s between numbers are significant. 30.09 has 4SF Zero’s that precede are NOT significant. 0.000034 has 2SF Zero’s at the end of decimals are significant. 0.00900 has 3 SF Zero’s at the end without decimals are ambiguous. 4050 has either 4SFor 3SF

  41. MEASUREMENTSSignificant Figures & Calculations 358.35 0.05655 Adding/subtracting 345.678 0.07283 + 12.67 -0.0162789 Multiplying/dividing 12.0340 x 3.98 = 98.657 / 43 = 47.9 2.3

  42. Metric System – a Unit for All Occasions • Systeme Internationale (SI) is standard units for science. • Prefixes adjust the scale to suit the object: • 1 m = .001 km • 1 m = 1,000 mm • 1 m = 1,000,000,000 nm • System is based on powers of ten for easy use. (C)2007 Prentice Hall

  43. MEASUREMENTS • There are different types of measurements that can be made in the laboratory like mass, time, volume, and length. • These measurements can be made using either the metric system or the English system. The metric system is based on increments of 10. 1 base = 100 centibases “c” = centi 1 base = 1000 millibases “m” = milli 1 kbase = 1000 bases 1 base = 106 microbases “m” = micro k = kilo 1 base = 109 nanobases “n” = nano • The first step to understanding measurements is to learn the types, symbols, & units associated with these measurements. (C)2007 Prentice Hall

  44. The Base Units • The most important and frequently used units in chemistry: (C)2007 Prentice Hall

  45. MEASUREMENTS • There are different types of measurements that can be made in the lab for length, mass, volume, temperature, area, time, heat and pressure. (C)2007 Prentice Hall

  46. Using Conversion Factors 1 m = 1,000 mm • Divide one number by the other and we have a conversion factor. (C)2007 Prentice Hall

  47. The sizes of atoms • Atoms are measured in picometers: 1,000,000,000,000 pm = 1m • Convert picometers to meters • Rationale for using prefixes: size of a Kr atom is 110 pm: (C)2007 Prentice Hall

  48. Importance of Dimensional Analysis • Height of each letter = 0.000 000 000 00130 km • Radius one atom = • Convert to m • Convert to pm • Atoms in figure are Xe (C)2007 Prentice Hall

  49. MEASUREMENTS Since two different measuring systems exist, a scientist must be able to convert from one system to the other. CONVERSIONS Length  1 in = 2.54 cm  1 mi = 1.61 km Mass  1 lb.... = 454 g  1 kg = 2.2 lb.... Volume  1 qt = 946 mL  1 L = 1.057 qt  4 qt = 1 gal  1 mL = 1 cm3 (C)2007 Prentice Hall

  50. Units of Volume • Standard metric unit for volume in chemistry is milliliter mL. • Equivalent in cubic centimeter cm3 or cc. 1 mL = 1 cm3 = 1 cc • Other volume units are nanoliter (nL), picoliter (pL) and femtoliter (fL). (C)2007 Prentice Hall

More Related