1. Matter and Measurement http://www.skanschools.org/webpages/rallen/

2. Matter Anything that has mass and volume (takes up space)

3. Matter Pure substances (can NOT be separated by physical means) Each piece looks the same (PURE!) Each piece has the exact same composition Mixtures (CAN be separated by physical means Each piece is different (not pure)

4. Pure Substances Can NOT be separated by chemical means Can be separated by chemical means, ONLY Element (simplest form of matter) Example: Monoatomic= Na (One atom) Diatomic= O 2 (Two atoms) Compound or Molecule (2 or more different elements chemically combined) Example: NaCl H 2 O (Table Salt) (Water)

5. Pure Substances Particle Diagrams ElementsCompound

6. Mixtures Same Composition Throughout Different Composition Throughout Homogenous Mixture Uniform throughout-distinct pattern Example: salt water, iced tea Homo= same Heterogeneous Mixture Not uniform throughout- no pattern Example: Italian dressing, concrete, soil, chocolate chip cookies Hetero= different

7. Mixtures Particle Diagrams HeterogeneousHomogeneous

8. Properties of Matter Physical properties are the constants about a substance Can use our senses to observe them Do not require chemical analysis Ex: melting point, color, texture

9. Properties of Matter Extensive Property  a property that depends on how much material you are dealing with Energy, mass, heat Intensive Property  a property that does not depend on how much material you are dealing with (helps identity matter; a constant about a particular type of matter) Melting point, boiling point, color, density, hardness, solubility

10. Properties of Matter Chemical properties include behaviors substances adhere to when they react with other substances Examples reacting hydrogen gas with oxygen gas results in a combustion reaction Very reactive when in the presence of nonmetals pH

11. Physical vs. Chemical Changes Matter is always changing Physical Change  a change that does NOT alter the chemical properties of a substance Ex. Cutting paper, phase change Change in size or shape (same composition) Ex. ice melting to become liquid

12. Physical vs. Chemical Changes Chemical Change  a reaction in which the composition of a substance is changed Ex. rusting Properties different composition 1. Signs of a chemical reaction Color change Bubbling/fizzing Energy produce or consumed Ex. firewood burning

13. Do Now Change of MatterPhysical or Chemical Burning Toast Making Ice Cubes Lighting a Candle Spoiling Milk Making Kool Aid Physical Chemical Physical

14. Elements vs. Compounds Element= formula that contains only one symbol Compound = formula which contains 2 or more different symbols/elements

15. Separation of Matter Separation Apparatus Type of Separation (physical or chemical) Description of technique What types of matter will it separate Filtration PHYSICALFiltrate flows through filter paper, undissolved particles (solids) remain on the filter paper Heterogeneous mixtures or mixtures involving more than one phase (ex. Sand and water)

16. Separation of Matter Separation Apparatus Type of Separation (physical or chemical) Description of technique What types of matter will it separate Evaporation PHYSICALSeparate solute (dissolved solid) from solvent (liquid) by boiling solution Solute escapes Very limited precision Homogeneous Mixture (solution)

17. Separation Apparatus Type of Separation (physical or chemical) Description of technique What types of matter will it separate Distillation PHYSICALSeparate solute from solvent by boiling solution and recondensing in receiving flask (both solute and solvent captured) Separate 2 or more liquids w/different boiling points Homogeneous (can use to remove impurities from water)

18. Distillation

19. Separation Apparatus Type of Separation (physical or chemical) Description of technique What types of matter will it separate Chromatograph y PHYSICALSeparates particles based on 1)Size 2)Solubility Homogeneous

20. Chemical separation  requires reacting a sample with something else in order to turn it into a completely different compound

21. Scientific Notation Method for expressing very large or small numbers easily Ex. 6.02 x 10 23 atoms = 1 mole

22. Practice Write the following numbers in scientific notation 1. 34000000 = 2. 0.0000067 = 3. 25,864 = 3.4 x 10 7 6.7 x 10 -6 2.5864 x 10 4

23. Measurements and the Metric System In chemistry we measure matter using SI units SI  System International SI Unit  Base Units

25. SI Metric Prefixes Table C

26. SI Metric Prefixes Ex. In the word kilometer, the root word (or base unit) is “meter” and the prefix is kilo. Kilo = 1000 1 km = 1000 m

27. Conversion Factors A mathematical expression that relates two units that measure the same type of quantity 1 min = 60 sec 1000 g = 1 kg 1 L = 1000 mL

28. Conversion Factors Kilo Hecta Deca base unit deci centi milli g = gram m = meter L = liter Practice: 3 g = kg30007 m = mm0.007

29. Dimensional Analysis When you are required to solve a problem with mixed units, or to convert from one set of units to another Ex. How many minutes are there in the month of October? 60 minutes x 24 hours x 31 days = 1 hour1 day 44, 640 minutes

30. Accuracy vs. Precision Accuracy  how close your results are to the desired value Ex. Hitting bulls eye when you’re aiming for it For most experiments, ACCURATE means +/- 5% from the expected value Precision  how close your results are to one another; how repeatable your results are; consistency/grouping

31. There are two kinds of numbers in the world: exact: example: There are exactly 12 eggs in a dozen. example: Most people have exactly 10 fingers and 10 toes. inexact numbers: example: any measurement. If I quickly measure the width of a piece of notebook paper, I might get 220 mm (2 significant figures). If I am more precise, I might get 216 mm (3 significant figures). An even more precise measurement would be 215.6 mm (4 significant figures).

32. Significant Figures Aka Sig Figs A method for handling UNCERTAINTY in all measurements This arises due to the fact that we have different equipment with different degrees of ACCURACY Sig figs are associated with MEASURED VALUES EXACT NUMBERS do NOT COUNT when determining sig figs Ex. Atomic masses on the periodic table Conversions 1 in = 2.54 cm

33. Significant Figures The Atlantic/Pacific Method Determine if a decimal point is present If yes think “P” for present  P = Pacific Coast 1. Start at the first nonzero number 2. Count all the way to the Atlantic-NO EXCEPTIONS If no think “A” for absent  A = Atlantic Coast 1. Start at first nonzero number Count all the way to the Pacific-NO EXCEPTIONS

34. Significant Figures Rules 1) ALL non-zero numbers (1,2,3,4,5,6,7,8,9) are ALWAYS significant. 2) ALL zeroes between non-zero numbers are ALWAYS significant (CAPTIVE Zeros) Ex. 40.7 L or 87,009 km 3) For numbers less than one, all zeros to the left of the 1 st nonzero number are NOT SIGNIFICANT (Leading Zeros) 0.009587 m or 0.0009 kg

35. Significant Figures Rules 4) Zeros at the end of a number and to the right of a decimal point are SIGNIFICANT (Trailing Zeros) 85.00 g or 9.07000000 L 5) Zeros at the end of a whole number may be significant or not. If there is a decimal after the last zero, they are significant. If there is no decimal after the end zeros, they are NOT significant 2000. m or 2000 m

36. Significant Figure Rules 6) Exponential digits in scientific notation are not significant 1.12x10 6 has three significant digits (1, 1, and 2.)

37. Significant Figures Number# Sig Figs 48,9235 3.9674 900.065 0.0004 ( 4 x 10 -4 )1 8.10005 501.0406 3,000,000 ( 3 x 10 6 )1 10.0 ( 1.00 x 10 1 )3

38. Using Sig Figs in Calculations General Rule  your final answer must be expressed in the lowest amount of significant figures that were originally given to you

39. OperationRuleExamples Multiplication/Divis ion Perform operation as normal & express answer in the least # of sig figs that were given to you 12.257 x 1.162 = 14.242634 Answer 14.24 Addition/Subtractio n Line decimal points up; round final answer to lowest decimal place 3.9 5 2.8 79 + 213.6 220.4 29

40. Measuring Matter Mass vs. Weight MassRelationshipWeight How much matter something has Directly proportional: As mass increases weight increases Depends on gravity (force pulling an object toward earth)

41. Measuring Matter Chemistry deals mainly with mass We have the same mass on earth and the moon. The different forces of gravity on each cause us to weigh more on earth than on the moon

42. Measuring Matter Volume  the amount of space an object takes up Techniques: Liquids  use graduated cylinder, burette (beaker, flask) Regular Solids  measure dimensions and use ( l x w x h) Irregular Solids  displacement method

43. Measuring Matter Density  amount of mass in a given space; RATIO of mass to volume Formula: D = mass  Volume mVmV