Honors Chemistry - PowerPoint PPT Presentation

honors chemistry n.
Download
Skip this Video
Loading SlideShow in 5 Seconds..
Honors Chemistry PowerPoint Presentation
Download Presentation
Honors Chemistry

play fullscreen
1 / 75
Honors Chemistry
229 Views
Download Presentation
ken
Download Presentation

Honors Chemistry

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. Honors Chemistry Chapter 1 Matter and Measurement

  2. 1.1 What is chemistry? • Chemistry is the study of the properties of matter and their changes. • Matter: The physical material of the universe. • Anything that has mass and • Occupies space (volume)

  3. Matter • Matter: Anything that has mass and takes up volume • Mass: the amount of material (stuff) in an object (kg, g, lb, t ) • Volume: the amount of space an object occupies (cm3, mL, L, ci)

  4. Elements • Very basic elementary substances. • Made of unique atoms. • Can not be broken down further without changing the identity. • Named for Latin or Greek characters or places. • Represented by 1 or 2 letter symbols. • First letter CAPITALIZED • Second letter lower case • C, O, Co • CO vs. Co Hydrargyrum Greek for “liquid silver”

  5. Molecules • Combinations of two or more atoms held together in specific shapes. • Properties are based on the structure and arrangement , and the number and types of atoms present. Methane

  6. Classifications of Matter • How is matter classified? (2 ways) • 1. Macroscopic level (more than one molecule) • a. Gas or vapor (g): no volume or shape, highly compressible • gases expands to occupy it’s container. • b. Liquid (l):volume independent of container, no shape, incompressible (relatively) • c. Solid (s):volume and shape independent of container, rigid, incompressible.

  7. 2. Molecular level (single molecules) • a. gas:molecules far apart, high speed, collide often with each other, and container walls. (kinetic molecular [KM] theory) • b. liquid:molecules closer than gas, move rapidly, slide over each other • c. solid:molecules packed closely, definite arrangement, don’t move.

  8. States of matter (4) • 1. Solids • Hold a particular shape. • Have a definite volume. • Particles arranged in an orderly manner. • High density

  9. States of Matter • 2. Liquid • Does not hold its own shape, takes the shape of its container. • Has definite volume. • Particles arranged randomly • High density • Not compressible

  10. States of Matter • 3. Gas • No definite shape or volume. • Expands to occupy the space of its container. • Extremely low density. • Particles far apart. • Highly compressible. • Density depends on pressure. • High pressure forces particles closer.

  11. States of Matter • 4. Plasma • Ionized gasses. • Ions and electrons (charged particles) • Found in the center of stars and space. • Emit light when “excited” • Eg. Lightning, Northern Lignts

  12. Matter

  13. Matter can be classified into four groups. Matter Heterogeneous Homogeneous AKA: Pure Substance Solution Mixtures Element Compound

  14. Pure Substances(2 types) • Pure substance is a substance with constant composition • Pure substances have fixed measurable properties • Pure substances cannot be found in nature. C6H12O6 H2O

  15. Elements (pure substance #1) Symbol Atomic Number • Simplest form of matter. • Can not be broken down by ordinary chemical means. • 92 naturally occurring elements.(118 total) • Arranged in the periodic table. • Solids, Liquids, Gases (at room temp) Mass Number (Atomic Weight) http://www.funbrain.com/periodic/

  16. K Rb Element names and symbols to memorize. (on periodic table)

  17. Cr Mn Fe Co Ni Cu Zn Ag Hg Au Diatomic Elements H2 N2 O2 F2 Cl2 Br2 I2 BrINClHOF Transition Metals to Memorize

  18. Compounds (pure substance #2) Two or more elements combined in a chemical reaction. • Elements are in fixed proportions to each other. • Written using symbols with subscripts to denote the ratio of elements. • Can not be separated by physical means. H2O: 2 Hydrogen: 1 Oxygen

  19. Mixtures • A combination of two or more pure substances. • Two types of mixtures.

  20. Heterogeneous mixture(hetero: different) • A MIXTURE is a combination of • two or more pure substances that are • not chemically united and • do not exist in fixed proportions to each other. • Do not have uniform composition. • Most natural substances are mixtures. • A heterogeneous mixture consists of visibly different substances or phases. • A mixture can be physically separated into pure compounds or elements.

  21. Methods of separating mixtures • 1. Decantation: pouring a less dense liquid off the top of a more dense liquid. • 2. Filtration: Separating particles based on size or phase. Larger particles are trapped in the filter. • Liquid passes through filter, solid is trapped. Separating Funnel

  22. Homogeneous Mixture(Homo: Same) Solution • A homogeneous mixture has the same uniform appearance and composition throughout. Many homogeneous mixtures are commonly referred to as solutions. • All components are all in the same phase. • Particles are uniform in size (atoms or molecules) • Can not be separated by physical means. • Can be separated based on differences in properties of components.

  23. Methods of separating solutions • 3. Crystallization: Evaporating one component to leave another. • 4. Chromatography: Separating components based on specific gravity. • 5. Distillation: Separation of liquids based on differences in boiling point.

  24. Properties of compounds vs. elements. • A compound may have different properties than the elements composing it. • Hydrogen H2: Gas • Oxygen O2: Gas • Water H2O: Liquid

  25. Joseph Proust (1799) • French Chemist • Developed The Law of Definite Proportions • Aka the Law of Constant Composition • Compounds always contain the same elements in the same proportion by mass.

  26. Law of Definite Proportions • H20 (by mass is always) • 88.9% Oxygen, 11.1% Hydrogen • If we had an 80g sample of H20 how much is O? • .889 x 80 = 71g • How much is H? • .111 x 80 = 9g

  27. 1.2 Properties of Matter • Physical Properties: • Characteristics that can be observed and measured without altering the identity of the substance. • Density, color, melting point, odor, boiling point, etc. Ice, water and steam.

  28. Physical Properties: 2 Types • 1:Intensive Properties: • Do not depend on the amount of matter present. • Examples: • Density • Conductivity • Melting point • Boiling point • Temperature • Pressure

  29. Physical Properties: 2 Types • 2: Extensive Properties • Depends on how much material is on hand. • Changes with the amount of material present. • Examples: • Mass • Volume • Weight

  30. Chemical Properties • The way matter behaves when brought into contact with other substances, or a source of energy. • Describes how substances react with other substances. • Example: • Flammability • Inert (nonreactive)

  31. “Oh the humanity”

  32. Changes in Matter • Physical Changes: • Alter the form, but do not alter the identity of the substance. • Change in physical appearance. • Examples: • Crushing • Tearing • Changes of state (phase)

  33. Chemical Changes(Reactions) • Chemical Changes (reactions) • Changes a substance into chemically different substances. • Irreversible • Alters the identity of the substance being changed. • Examples: • Wood burning, • Food cooking, • Iron rusting

  34. Homework Practice • Exercises #s 2, 3, 5, 10, 12, 15 • On page 29 in Textbook

  35. 1.4 Units of Measurement & the Metric System A. Why Measure? In order to get complete observations, quantifying is essential.

  36. B. Consists of: 1. Quantity a. Anything that can be measured b. Involves a number followed by a unit c. Examples: Volume, Mass, Length, Time 2. Unit (can be English (customary) or metric) a. Is what unit the quantity is measured in b. Examples: Liter, Gallon, Gram, Foot, Meter, Second C. Why Use the Metric System? 1. It’s easier because it is based on units of 10 2. Everyone uses it so it gives consistency 3. There are standards that everyone can compare to

  37. D. International System of Units – called SI 1. Composed of fundamental or base units. Chart 1-1

  38. Prefixes used in the Metric System Greek mu Page 14 in your book

  39. Length and Mass • SI unit of length is the meter (m) (ca. 39 inches) • Mass: The measure of the amount of material (matter, stuff) in an object. • SI base unit of mass is the kilogram (kg) (ca. 2.2 lbs) • This is odd because it uses a prefix instead of gram alone.

  40. Temperature • The measure of how fast molecules in a substance are moving. • Or • The measure of the average kinetic energy of particles in a substance.

  41. Temperature scales Say aaah. • Fahrenheit • Named after Gabriel Fahrenheit (1686-1736) • Thermometer maker who devised his own temperature scale. • Based on the temperature of an equal ice-salt mixture. (0ºF) • Average temperature of a healthy horse. (100 º F) • H2O freezes at: 32º F • H20 boils at 212º F Why me?

  42. Temperature Scales • Celsius (Centigrade) • Anders Celsius • Developed a temperature scale more compatible with the metric system. • Based on H20 • H2O freezes at: 0º C • H20 boils at 100º C (1701- 1744) Thanks Mr. Celsius!!

  43. Fahrenheit vs Celsius • Fahrenheit • 212º - 32º = 180º from boiling to freezing. • Celsius • 100º - 0º = 100º from boiling to freezing. • 180/100 = 9/5 • Celsius degrees are 9/5 bigger than Fahrenheit • The Fahrenheit scale also begins 32º above the Celsius scale.

  44. Converting between Fahrenheit and Celsius • Taking the degree size, and starting point in consideration, we get: • ºF = 9/5ºC + 32 and/ or • ºC = 5/9(ºF – 32) *watch order of ops*

  45. Conversion Practice • The temperature of a healthy human is 98.6ºF. Convert this to Celsius. • ºC = 5/9 (98.6ºF – 32) • 37ºC • A baby has a fever of 39.5ºC. What is her temperature in Fahrenheit? • ºF = 9/5 x 39.5ºC + 32 • 103.1 ºF

  46. The Kelvin Temperature Scale • The SI unit used to measure temperature. • Named for William Thomson, Lord Kelvin • Not Melvin Kelvin • The º symbol is not used for kelvin (K) • 1ºC = 1K 1824-1907

  47. The Kelvin Temperature Scale(aka Absolute Scale) • The zero point on the Kelvin scale is absolute zero • 0K = -273ºC • Absolute zero is the point at which all molecular motion ceases. Lowest temperature possible. • Can absolute zero ever be reached?

  48. Converting between Celsius and Kelvin • ºC = K – 273 • At 50K air will freeze to a solid. Convert this to ºC. • ºC = 50 – 273 = -223 ºC • K = ºC + 273 • Antifreeze (ethylene glycol) boils at 199ºC. Convert this to K. • K = 199 + 273 = 472 K

  49. Comparing temperature scales • Note that in ºC, and K, there are 100 degrees between the freezing and boiling points of H20. • Note that in ºF there are 180º between the freezing and boiling points of H2O. • The Celsius scale begins 273.15º higher than Kelvin. • The Fahrenheit scale begins 32º higher than Celsius. Page 15