1. Chapter 2 Classification of Matter

2. Classification of Matter Concept Map

3. Pure Substances Composition Constant Element (one kind of atom)
Compound (two or more atoms chemically combined)

4. Mixtures of Matter
Mixture- a combination of two or more pure substances in which each pure substance retains its individual chemical properties.
Composition is variable - (the number of mixtures that can be created by combining substances is infinite)

5. Types of Mixtures Homogeneous Mixtures Heterogeneous Mixtures
Has a constant composition throughout
Also called a “solution” (i.e., salt water)
Heterogeneous Mixtures
Do not blend smoothly throughout and in which the individual substances remain distinct.
Examples: mixture of sand and water, fresh-squeezed orange juice, pizza

6. Types of Solutions Solid-solid (alloy-steel)
Solid-liquid (sugar in water)
Liquid-liquid (vinegar)
Liquid-gas (water vapor in air)
Gas-liquid (carbonated drinks)
Gas-gas (air is a mixture of nitrogen, oxygen, argon)

7. Types of Heterogeneous Mixtures
Colloid- contains tiny particles that do not settle (i.e., milk, gelatin)
Suspension- contains larger particles that eventually settle out. Particles have to be re-suspended. (i.e., chocolate milk, orange juice)

8. Techniques for Separating Mixtures
Filtration- uses a porous barrier to separate a solid from a liquid
Distillation- A mixture is heated until the substance with the lowest boiling point boils to a vapor that can then be condensed into a liquid and collected.
Crystallization - separation technique that results in the formation of pure solid particles of a substance from a solution containing the dissolved substance.
Chromatography- separates the components of a mixture (called the mobile phase) on the basis of the tendency of each to travel or be drawn across the surface of another material (called the stationary phase).
A magnet can be used to separate magnetic particles from others (Ex. separate sand and iron filings using a magnet to extract the iron based on physical property of magnetism)

9. Techniques for Separating Mixtures
Filtration- uses a porous barrier to separate a solid from a liquid
Distillation- A mixture is heated until the substance with the lowest boiling point boils to a vapor that can then be condensed into a liquid and collected.
Crystallization - separation technique that results in the formation of pure solid particles of a substance from a solution containing the dissolved substance.
Chromatography- separates the components of a mixture (called the mobile phase) on the basis of the tendency of each to travel or be drawn across the surface of another material (called the stationary phase).
A magnet can be used to separate magnetic particles from others (Ex. separate sand and iron filings using a magnet to extract the iron based on physical property of magnetism)

10. Properties & Changes in Matter
Physical Properties
Physical Changes
Chemical Properties
Chemical Changes
Physical and chemical properties depend on temperature and pressure.

11. Physical Properties
Can be observed or measured without changing the substance’s identity.
Ex. Density, color, odor, taste, hardness, melting point, boiling point, solubility, state of matter (s,l,g), temperature

12. Physical Changes
Changes that may dramatically alter the appearance yet leave the composition unchanged
Examples: split, bend, crush, grind, changes in state of matter (boil, freeze, condense, vaporize, melt), sharpening a pencil, cutting a sheet of paper, breaking a crystal, crumpling a piece of paper

13. Chemical Properties
The ability of a substance to combine with or change into one or more other substances having different properties.
Examples:
Flammable
Supports combustion
the ability of iron to rust when exposed to air
Inability of iron to react with nitrogen gas at room temp.

14. Chemical Changes
A process that involves one or more substances changing into new substances
Commonly referred to as a chemical reaction
Examples:
Fermentation of grape juice
Rusting
Exploding
Oxidizing
Corroding
Tarnishing
Burning
Rotting

15. Evidence of a Chemical Reaction
Change in color, odor, appearance, production of a gas
Formation of a solid called a precipitate – a solid that “falls out” of solution)
Changes in energy (production of heat, light, or sound)

16. Law of Conservation of Mass
Mass is neither created nor destroyed during a chemical reaction.
Massreactants = Massproducts

17. Chapter 3 States of Matter & Gas Laws

18. Kinetic Theory of Matter
The idea that all matter is made up of tiny particles in constant motion.

19. Kinetic Theory of Matter – 4 assumptions:
Matter is made of atoms and molecules
These atoms and molecules are always in motion.
The higher the temperature, the faster the particles move.
At the same temperature, more massive particles move slower than less massive ones.

20. States of Matter
Solid
Liquid
Gas
Plasma

22. Solid Has a definite shape and volume
Particles are packed very closely together
Particles are incompressible
Expand when heated

23. Types of Solids
Crystalline solids-
Amorphous solids

24. Crystalline solids have a definite shape or pattern
a solid having a distinct shape because its atoms are arranged in orderly, repeating, three dimensional geometric patterns.
Examples: sodium chloride (table salt), sucrose (table sugar), diamonds

25. Amorphous solids
a solid having no form or shape; lacks an ordered internal structure
sometimes called slow-moving liquids: (i.e. wax, plastics, glass)
Glasses are supercooled liquids.(No regular pattern when shattered.)

26. Liquids
Have a constant volume, but take the shape of the container they are in
Particles are not rigidly held in place and are less closely packed than a solid
Thus, a liquid “flows”

27. Gases Have no definite shape or volume
Flow to conform to the shape of its container and fill the entire volume of the container

28. Gases (cont’d.) Particles can be very far apart, or close together
Gases are easily compressed
Examples: neon, methane, air

29. Temperature
A measure of the average kinetic energy of the particles.

30. Standard Temperature and Pressure
Is defined as 25˚C and
1 atmosphere of pressure

31. Temperature Scales Celsius- most desirable for laboratory work (°C)
Kelvin- the absolute scale (also the SI unit of temperature) (K)
Fahrenheit- the old English scale (never used in the lab) (°F)

32. Kelvin SI base unit of temperature Absolute temperature scale
Absolute zero = zero degrees Kelvin, is the theoretical point at which all particle motion stops.
Absolute zero is -273˚C
K = ˚C + 273
Do not use degree (°) symbol, just K

33. Celsius Water freezes at 0˚C and boils at 100˚C ˚C = K- 273
˚C = 5/9 (˚F-32)

34. Fahrenheit Will not be used in the science lab
Water freezes at 32˚F and boils at 212˚F
˚F = 9/5˚C

36. Will change with temperature
Properties of Water
Will change with temperature

37. WATER (at standard temperature and pressure--STP)
Liquid
Density is 1.00 g/cm3
Chemically unreactive

38. WATER (at 0°C) Solid Density is 0.9167 g/cm³
The crystalline structure of solid water makes it less dense than it is in its liquid form, therefore, it floats.
Water is unusual because its solid form is less dense than its liquid form. For mmost matter, this is not true

39. Water (at 100˚C and up) Is a gas Has a density of 0.0006 g/cm3
Reacts readily with nitrogen gas

40. Kinetic Energy and Temperature
The average kinetic energy of the particles of a substance is proportional to the temperature of the substance.

41. Kinetic Theory of Gases
A gas is composed of particles, usually molecules or atoms
Gas particles move rapidly in constant random motion.

42. Changes in the State of Matter
Are a result of changes in temperature (the average kinetic energy of the particles) or changes in pressure.

43. Changes in the State of Matter
Melting-Solid to liquid
Freezing-Liquid to solid
Vaporization-Liquid to gas
Sublimation-Solid to gas
Condensation-Gas to liquid
Deposition-Gas to solid

44. Changes in the State of Matter are caused by:
Changes in temperature
(average kinetic energy of the particles that make up matter)
Changes in pressure
(i.e., decrease the volume at constant temperature causes an increase in pressure. Higher pressure = more particle collisions with the side of the container, and thus, temperature increases, causing a change in physical state)

45. Vapor
Refers to the gaseous state of a substance that is solid or liquid at room temperature
Example: steam

46. Heat of Vaporization
The amount of energy needed to change a material from a liquid to a gas.
For water, it is 2260 kJ/kg.

47. Condensation
The change of a substance from a gaseous state to a liquid state due to a loss of heat.
Ex: Water vapor condensing on a cold surface.

48. Sublimation
A change of a solid directly to a gas without going through the liquid state.
Ex: dry ice forming “smoke”, solid iodine

49. Heat of Fusion
The amount of energy needed to change a material from the solid state to the liquid state.
For water, the heat of fusion us 334 kJ/kg.

50. Evaporation or Vaporization
The gradual change of a substance from a liquid to a gaseous state at temperatures below the boiling point.

51. GAS LAWS

52. Biography of Robert Boyle
Boyle’s Law
P1V1 = P2V2
When temperature is held constant, volume and pressure are inversely proportional.
Biography of Robert Boyle
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Image URL--

53. Boyle’s Law

54. Gay-Lussac’s Law
The pressure of a gas increases as the temperature increases, if the volume of the gas does not change. The pressure decreases as temperature decreases.
Biography of Gay-Lussac

55. Gay-Lussac’s Law

56. Charles’ Law
V1 = V2
T T2
When pressure is held constant, volume and temperature and are directly proportional.
Charles' Biography
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57. Charles’ Law