1 / 41

Chapter 1

Chapter 1. Basic Concepts of Matter. Classifying Matter. Matter Anything that has mass and occupies space Mass vs. Weight Kinetic-Molecular Theory All matter consists of extremely tiny particles in constant motion. States of Matter. Solid

sophie
Download Presentation

Chapter 1

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. Chapter 1 Basic Concepts of Matter

  2. Classifying Matter Matter • Anything that has mass and occupies space Mass vs. Weight Kinetic-Molecular Theory All matter consists of extremely tiny particles in constant motion

  3. States of Matter • Solid • -Closely packed together with a definite ridged shape • -Vibrate back and forth in a confined space • -the particles are not able to move past one another • Liquid • -arranged randomly with a definite volume • -“fluid” • -the particles are not confined in space and can move past one another • Gas • -no definite shape or volume • -“fluid” • -the particles are far apart and move very rapidly colliding with other particles and the container walls

  4. Categorizing Matter • Elements • -cannot be decomposed into simpler form via chemical reactions • -found on periodic chart • -atoms are the smallest particle that retains the characteristic properties of the elements • Pure Substance • -consists of all the same substance (pure gold, distilled water, etc) • -have a set of unique properties that identifies it

  5. Categorizing Matter • Chemical Compounds • -two or more elements in a definite ratio by mass with unique properties that separate them from the individual elements • -can be decomposed into the constituent elements by chemical reactions • -chemical compounds are held together by a chemical bound • Water– hydrogen and oxygen • Carbon dioxide – carbon and oxygen

  6. Categorizing Matter • Mixtures • two or more pure substances in the same container homogeneous mixtures (solution) • -uniform composition throughout • -single phase • -cannot be separated easily heterogeneous mixtures • -nonuniform composition thoughout • -easily separated

  7. Physical and Chemical Changes Physical changes • changes in physical properties • -melting, boiling, and cutting Chemical changes changing one or more substances into one or more different substances (chemical reaction) • 2H2 + O2 -> 2H2O

  8. Chemical and Physical Properties Chemical Properties observed during a chemical reaction (change in chemical composition) • -rusting, oxidation, burning… • -chemical reactions Physical Properties observed without changing the substance’s composition • -allow for identification and classification • -density, color, solubility, melting point…

  9. Classification of Physical Properties Extensive Properties depend on the amount of substance present -mass or volume Intensive Properties do not depend on the amount of substance present -melting point, boiling point, density…

  10. Density • Describes how compact a substance is • Who “discovered” density? • Density = mass/volume or D = m/V

  11. Density • Example: Calculate the density of a substance if 742 grams of it occupies 97.3 cm3. • 1cm3 = 1mL => 97.3cm3 = 97.3mL

  12. Density • Example: You need 125 g of a corrosive liquid for a reaction. What volume do you need? • liquid’s density = 1.32 g/mL

  13. Units of Measure • Qualitative measures • Nonnumerical experimental observations describing the identity of a substance in a sample • Quantitative measures • Numerical experimental observations describing how much of a particular substance is in a sample System International d’Unites (SI) measurement system used in the sciences based on the metric system

  14. Math Review and Measurements • We make measurements to understand our environment: • Human senses: sight, taste, smell, hearing… • Our senses have limits and are biased • Instruments: an extension of our senses meter sticks, thermometers, balances • These are more accurate and precise • All measurements have units • METRIC SYSTEM vs. British System

  15. SI units QuantityUnitSymbol • length meter m • mass kilogram kg • time second s • current ampere A • temperature Kelvin K • amt. substance mole mol

  16. Measurements in Chemistry NameSymbolMultiplier • mega- M 106 (1,000,000) • kilo- k 103 (1,000) • deka- da 10 • deci- d 10-1 (0.1) • centi- c 10-2 (0.01)

  17. Measurements in Chemistry NameSymbolMultiplier • Milli- m 10-3(0.001) • Micro-  10-6(0.000001) • Nano- n 10-9 • Pico- p 10-12 • Femto- f 10-15

  18. Units of Measurement Length • Measure of space in any direction • -derived unit cm • -standard length is a meter (m)

  19. Units of Measurement • Volume • Amount of space occupied by matter • -derived unit: mL or cm3 (cc) • -liter (L) is the standard unit

  20. Units of Measurement • Time (t) • Interval or duration of forward events • -standard unit is the second (s) • Mass (m) • measure of the quantity of matter in a body • Weight (W) • measure of the gravitational • attraction (g) for a body (w=m x g) 1 kg = 1000g 1 kg = 2.2 lbs 1 g = 1000mg

  21. Heat and Temperature • Heat (q) vs. Temperature (T) • 3 common temperature scales: • all use water as a reference • -Fahrenheit (F) • -Celsius (C) • -Kelvin (K)

  22. Temperature Reference Points Melting Point Boiling Point of waterof water • 32 oF 212 oF • 0.0 oC 100 cC • 273 K 373 K • Body temperature 37.0 oC or 98.6 oF • 37.2 oC and greater—sick • 41 oC and greater, convulsions • <28.5 oC hypothermia Fahrenheit Celsius Kelvin

  23. Temperature Scales

  24. Temperature Scales Fahrenheit to Centigrade Relationships • Example: Convert 211 oF to degrees Celsius. Example: Express 548 K in Celsius degrees.

  25. Accuracy • how closely measured values agree with the correct value • Precision • how closely individual measurements agree with each other Precision and Accuracy Precise Accurate Both Neither

  26. Mathematics in Chemistry • Exact numbers (counted numbers) • 1 dozen = 12 things • Measured Numbers • Use rules for significant figures • Use scientific notation when possible • Significant figures • digits in a measured quantity that reflect the accuracy of the measurement • -in other words, digits believed to be correct by the person making the measurement • Exact numbers have an infinite number of significant figures 12.000000000000000 = 1 dozen

  27. Significant figures (numbers/digits) Why use significant numbers? • -Calculators give 8+ numbers • -People estimate numbers differently • -Dictated by the precision (graduation) on your measuring device • -In the lab, the last significant digit is the digit you (the scientist) estimate Scientists have develop rules to help determine which digits are “significant”

  28. Rules for Significant Figures 1. All Nonzero numbers are significant!!! 2. Leading zeroes are never significant • 0.000357 3. Imbedded zeroes are always significant • 3.0604 4. Trailing zeroes may be significant - You must specify significance by how the number is determined or even written • 1300 nails - counted or weighed? • 1.30000 –How many significant figures?

  29. Significant Figures • Multiplication & Division rule: • The product retains the number of significant figures that corresponds to the multiplier with the smallest number of significant figure (sig. fig.)

  30. Significant Figures • Addition & Subtraction rule: • Answer retains the smallest decimal place value of the addends.

  31. Scientific Notation Express answers as powers of 10 by moving the decimal place right (-) or left (+) • Use of scientific notation is to remove doubt in the Significant Figures: 2000  2 x 103 15000  1.5 x 10? 0.004  4 x 10-3 0.000053  __.__ x 10? In scientific notation, zeros are given if they are significant!!! 1.000 x 103 has 4 significant figures 2.40 x 103 has ? significant figures Key to Sig. Figs… Locating the decimal and deciding when to count the zeros!!!

  32. Review #2 • Units of Measure • -length • -volume • -time • -mass • -weight • Heat vs. Temperature • -three temperature scales • -temperature conversions • Precision vs. Accuracy • Significant Figures • Scientific Notation

  33. Conversion Factors • Length • 1 m = 39.37 inches • 2.54 cm = 1 inch • Volume • 1 liter = 1.06 qt • 1 qt = 0.946 liter • See Text for more conversion factors

  34. Conversion Factors Why do conversions? • -Scientists often must convert between units Conversion factors can be made for any relationship of units -Use known equivalence to make a fraction that can be used to “convert” from one unit to the other

  35. Dimensional Analysis • 1 inch = 2.54 cm • Use the ratio to perform a calculation so the units will “divide out” Example: Convert 60 inches to centimeters

  36. Dimensional Analysis • Example: Express 9.32 yards in millimeters. 3 ft = 1 yard 1 ft = 12 in or 1 in = 2.54 cm 100 cm = 1 m 1000 mm= 1 m

  37. Dimensional Analysis • Example: Express 627 milliliters in gallons. • 1 liter = 1.06 qt • 1 qt = 0.946 liter

  38. Practice on your Own 1kg = 2.20 lbs • Convert 25 g to lbs • Convert 1 mL to Liters • Convert 20 meters to cm

  39. Dimensional Analysis Area = length x width Area is two dimensional thus units must be in squared terms: • Express: 2.61 x 104 cm2 in ft2

  40. Dimensional Analysis Volume =length x width x height • Volume is three dimensional thus units must be in cubic terms Express: 2.61 ft3 in cm3 • this volume is used in medical measurements--cc

  41. Percentage • Percentage is parts per hundred of a sample • % = x100 • Example: A 335 g sample of ore yields 29.5 g of iron. What is the percent of iron in the ore? g of substance total g of sample

More Related