1 / 50

Chemistry: The Study of Change

Chemistry: The Study of Change. Chapter 1. CHEMISTRY. UNITS. PREFIXES. The number of atoms in 12 g of carbon:. 602,200,000,000,000,000,000,000. The mass of a single carbon atom in grams:. 0.0000000000000000000000199. SCIENTIFIC NOTATION. N is a number between 1 and 10.

mallory-daniels
Télécharger la présentation

Chemistry: The Study of Change

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. Chemistry: The Study of Change Chapter 1

  2. CHEMISTRY

  3. UNITS

  4. PREFIXES

  5. The number of atoms in 12 g of carbon: 602,200,000,000,000,000,000,000 The mass of a single carbon atom in grams: 0.0000000000000000000000199 SCIENTIFIC NOTATION N is a number between 1 and 10 n is a positive or negative integer 6.022 x 1023 1.99 x 10-23

  6. move decimal left move decimal right SCIENTIFIC NOTATION 568.762 0.00000772 n > 0 n < 0 568.762 = __________ 0.00000772 = ___________

  7. SIGNIFICANT FIGURES

  8. Once you start the counting don’t stop!

  9. SIGNIFICANT FIGURES Rule 1: Every nonzero digit in a measurement is significant. Examples: 24.7 0.22 569 Rule 2: Zeros appearing between nonzero digits are significant. Examples: 7003 60.8 0.502 Rule 3 A ZERO is NOT significant when it is a placeholder. A placeholder is used to show the location of the decimal point. Examples: .00099 5280 700

  10. SIG FIGS CONT. Rule 4: Zeros at the end of a number and to the right of a decimal point are always significant. Examples: 86.0 46.00 1.010 Rule 5: When a number is counted or defined within a system of measurement, there is an infinite amount of significant digits. Examples: 11 students 100 cm = 1 m

  11. COUNT THE SIG FIGS: 24 mL 3001 g 0.0320 m3 6.4 x 104 molecules 560 kg

  12. 89.332 + 1.1 one significant figure after decimal point two significant figures after decimal point 90.432 round off to ________ round off to ________ 3.70 -2.9133 0.7867 SIG FIGS – Addition and Subtraction The answer cannot have more digits to the right of the decimal point than any of the original numbers.

  13. 3 sig figs round to 3 sig figs 2 sig figs round to 2 sig figs SIG FIGS – Multiplication and Division The number of significant figures in the result is set by the original number that has the smallest number of significant figures 4.51 x 3.6666 = 16.536366 = _______ 6.8 ÷ 112.04 = 0.0606926 = _______

  14. 1000 mL 1L L2 1.63 L x = 1630 mL mL 1L 1.63 L x = 0.001630 1000 mL FACTOR LABEL METHOD • Determine which unit conversion factor(s) are needed • Carry units through calculation • If all units cancel except for the desired unit(s), then the problem was solved correctly. How many mL are in 1.63 L? 1 L = 1000 mL

  15. The speed of sound in air is about 343 m/s. What is this speed in miles per hour? meters to miles 1 mi = 1609 m seconds to hours 1 min = 60 s 1 hour = 60 min

  16. Important things to consider when solving problems and performing experiments….

  17. Volume – SI derived unit for volume is cubic meter (m3) 1 L = 1000 mL = 1000 cm3 = 1 dm3

  18. weight = c x mass A 1 kg bar will weigh on earth, c = 1.0 1 kg on earth on moon, c ~ 0.1 0.1 kg on moon Matter - anything that occupies space and has mass. mass – measure of the quantity of matter SI unit of mass is the kilogram (kg) weight – force that gravity exerts on an object

  19. mass density = volume A piece of platinum metal with a density of 21.5 g/cm3 has a volume of 4.49 cm3. What is its mass? m D = V Density – SI derived unit for density is kg/m3 1 g/cm3 = 1 g/mL = 1000 kg/m3 For Water: 1g/mL

  20. 0F = x 0C + 32 9 5 K = 0C + 273.15 273 K = 0 0C 373 K = 100 0C 32 0F = 0 0C 212 0F = 100 0C

  21. 0F – 32 = x 0C 0F = x 0C + 32 x (0F – 32) = 0C 5 5 5 0C = x (0F – 32) 0C = x (172.9 – 32) = 78.3 9 9 9 9 9 5 5 Convert 172.9 0F to degrees Celsius.

  22. Graphing Distance vs. Time Line of Best Fit DependentVariable Independent Variable

  23. Accuracy – how close a measurement is to the true value Precision – how close a set of measurements are to each other

  24. Percents and Percent Error • Percent Error: Measured Value - Accepted Value X 100 Accepted Value Example: The mass of a compound measured in a lab was 25.0 grams. The accepted value for this compound is 24.5 grams. Calculate the percent error. - 24.5 g 25.0 g X 100 = 2.04 % 24.5 g

  25. Scientific Method

  26. SCIENTIFIC METHOD logical approach to solving problems • Observation • Problem • Hypothesis • Experiment • Data • Analysis • Conclusion

  27. Observations • Qualitative: quality, non-numeric terms • Quantitative: quantity, numerical description

  28. MATTER Matter: anythingthat occupies space and has mass

  29. Three States of Matter Plasma

  30. Phase Changes liquid gas solid

  31. phase diagram: summarizes the conditions at which a substance exists as a solid, liquid, or gas. Phase Diagram of Water

  32. Phase Diagram Points • ________________: • above this point a substance becomes a supercritical fluid • critical temperature (Tc): temperature above which the gas cannot be made to liquefy, no matter how great the applied pressure. • critical pressure (Pc): minimum pressure that must be applied to bring about liquefaction at the critical temperature. • ________________: • point at which all three phases coexist

  33. Classification of Matter • Pure Substance: form of matter that has a definite composition and distinct properties 2. Mixture:combination of two or more substances in which the substances retain their own identities. water, ammonia, sucrose, gold, oxygen • mixed together physically • can usually be separated

  34. Examples: Types of Mixtures • Homogenous mixture – composition of the mixture is the same throughout. Examples: • Heterogeneous mixture – composition is not uniform throughout. 1.4

  35. Mixture Pictures

  36. Types of Mixtures • solution: mixture that remains uniformly mixed • solute: • solvent: • suspension:mixture where visible particles settle • colloid:mixture where particles are unevenly distributed but do not separate, positive Tyndall Effect

  37. Hom/Het? Soln/Susp/Coll? • Fog ____________ ___________ • Paint ____________ ___________ • Syrup ____________ ___________

  38. magnet distillation Physical means can be used to separate a mixture into its pure components. 1.4

  39. Methods of Separation PASTA/WATER SAND/IRON FILINGS SAND/WATER SALT /WATER FOOD COLORING/WATER BLOOD Strainer Filtration Physical Evaporation Centrifuge Distillation

  40. Element: • all atoms are the same • cannot be broken down by physical or chemical means • 114 elements named on the Periodic Table • 83 elements occur naturally on Earth gold, aluminum, lead, oxygen, carbon • many elements have been created by scientists technetium, americium, seaborgium

  41. Water (H2O) Glucose (C6H12O6) Ammonia (NH3) Compound: • 2 or more elements combined • Cannot be broken down by physical means • Can be broken down by chemical means • Appears different from original elements • Fixed ratios in definite proportions

  42. PHYSICAL AND CHEMICAL CHANGES

  43. Physical Properties and Physical Changes physical property: characteristic that can be observed or measured without changing the identity of the substance. physical change: change in a substance that does not involve a change in the identity of the substance.

  44. Physical Properties • Intensive: INDEPENDENT of amount of matter present (sample size) • Example: • Extensive: DEPENDENT on the amount of matter present (sample size) • Example:

  45. Chemical Properties and Chemical Changes • chemical property: a substance’s ability to undergo changes that transform it into different substances • Example: • chemical change: change in which one or more substances are converted into different substances • Example:

  46. Evidence of a Chemical Change

  47. sugar dissolving in water ice melting Physical or Chemical? hydrogen gas burns in oxygen gas to form water Remember: physical change does not alter the composition or identity of a substance. chemical change alters the composition or identity of the substance(s) involved.

More Related