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Unit 1

Unit 1 . Properties and Changes in Matter. Physical Properties. can be observed or measured without changing the composition of matter. are used to observe and describe matter. describe bulk quantities of atoms, not single atoms Ex. density, color, malleability.

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Unit 1

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  1. Unit 1 Properties and Changes in Matter

  2. Physical Properties • can be observed or measured without changing the composition of matter. • are used to observe and describe matter. • describe bulk quantities of atoms, not single atoms • Ex. density, color, malleability

  3. Intensive: Size doesn’t matter Properties that are INdependent of the amount of matter present (sample size) Examples: Color, odor, melting point, boiling point, density, malleability, ductile, temperature, hardness, luster Extensive: Size matters Properties that do depend on the amount of matter present (sample size) Examples: Mass, weight, volume, length Intensive vs. Extensive physical properties

  4. Chemical Properties • a material's properties that become evident during a chemical reaction. • any quality that can be established only by changing a substance's chemical identity. • Examples: rust, bonding potential, flammability

  5. Practice quiz • http://antoine.frostburg.edu/chem/senese/101/matter/classify-properties-quiz.shtml

  6. Measurement • a quantitative observation consisting of 2 parts: a number and a unit • Examples: 20 grams 6.63 9seconds

  7. Metric to metric conversions • King Henry Died By Drinking Chocolate Milk King - Kilo (K + base unit) Henry – Hecta (H + base unit) Died- Deka (Da + base unit) By – Base – meter, liter, gram Drinking – Deci (d + base unit) Chocolate – Centi ( c + base unit) Milk – Milli (m + base unit)

  8. 1 2 3 MetersLitersGrams How do you use the “ladder” method? 1st – Determine your starting point. 2nd – Count the “jumps” to your ending point. 3rd – Move the decimal the same number of jumps in the same direction. Starting Point Ending Point __. __. __. 2 3 1 Ladder Method KILO1000Units HECTO100Units DEKA10Units DECI0.1Unit CENTI0.01Unit MILLI0.001Unit 4 km = _________ m How many jumps does it take? 4. = 4000 m

  9. Compare using <, >, or =. 56 cm 6 m 7 g 698 mg Conversion Practice Try these conversions using the ladder method. 1000 mg = _______ g 1 L = _______ mL 160 cm = _______ mm 14 km = _______ m 109 g = _______ kg 250 m = _______ km

  10. Nano- milli divided by 1000000 Nano 1 X 10-9 units Micro

  11. Accuracy and Precision • Accuracy refers to the agreement of a particular value with the truevalue. • Precisionrefers to the degree of agreement among several measurements of the same quantity.

  12. Label each experiment

  13. A graduated cylinder is filled 5 times from the more accurate buret? Describe the accuracy and precision of the graduated cylinder. Trial Volume Shown Volume Shown By Buret By G. Cylinder 1 25 mL 26.54 mL 2 25 mL 26.51 mL 3 25 mL 26.60 mL 4 25 mL 26.49 mL 5 25 mL 26.57 mL Average 25 mL 26.54 mL

  14. Uncertainty in Measurement A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty.

  15. Sig digs and instruments • Include all known digits (marks on the instrument) PLUS one estimated digit in your measurement.

  16. Reading measuring devices You know 50 ml You know 2 ml You estimate .8ml The measurement is 52.8 ml

  17. How long is the green line?

  18. Rules for Counting Significant Figures (can be found on EOC chart) • Non-zero digits and zeros between non-zero digits are always significant. • Leading zeros are not significant. • Zeros to the right of all non-zero digits are only significant if a decimal point is shown. • For values written in scientific notation, the digits in the coefficient are significant. Sig Fig Song

  19. Rules for Counting Significant Figures - example Nonzero integersand zeros between non-zero digits are always significant. 3456 has 4 sig figs. 30506 has 5 sig figs.

  20. Rules for Counting Significant Figures - example Leading zeros are not significant. 0.0486 has 3 sig figs.

  21. Rules for Counting Significant Figures - example Zeros to the right of all non-zero digits are only significant if a decimal point is shown. 9.300 has 4 sig figs. 9,300 has 2 sig figs.

  22. Rules for Counting Significant Figures - example For values written in scientific notation, the digits in the coefficient are significant. 9.35 x 1023 has 3 sig figs.

  23. Rules for Significant Figures in Mathematical Operations Your answer can’t be more accurate than your least accurate measurement.

  24. Rounding If you have too many sig figs, you will need to round. Make sure the rounded answer is still close to the original answer. You may need placeholder zeroes to do this.

  25. Density • physical property of a substance • ratio of mass to volume • Independent of sample size • Density = mass = mvolume v • high density: a lot of mass in a small area • increasing the temperature decreases the density • increasing the pressure increases the density • **the density of water is 1 g/cm3 or 1 g/ml

  26. Determining mass and volume • To determine mass (a measure of the amount of matter in an object): use an electronic balance • To determine volume (a measure of the amount of space occupied): • Use a graduated cylinder for liquids • use v = l x w x h for regular solids • Use water displacement for irregular solids

  27. Density example Compare the mass, volume and density of sample A and B below. What would happen to the density of B if the sample were cut in half? A B Mass Volume Density

  28. Density example Compare the mass, volume, density of sample A, B and C below. C A B

  29. Density example Find the density of a material that has a mass of 100 grams and takes up a space of 25 cubic centimeters. Density = mass volume = m V = 100 g = 4 g/cm325 cm3 D= ? M= 100 grams V = 25 cm3

  30. Density example Find the volume of a 15 g rock whose density is 4.5 g/cm3. m = 15 g D = 4.5 g/cm3 v = ? D = m 1 v 4.5 g/cm3= 15 g v (4.5g/cm3 )v=15g v = 3.3 cm3

  31. Density example The picture shows a cube that contains 20 mL of a solution. The solution has a mass of 40 grams. What is the density in g/mL of this solution? A 60 g/mL B 20 g/mL C 2 g/mL D 0.5 g/mL

  32. Density example What is the mass of a 500.00 mL sample of seawater with a density of 1.025 g/mL? A 487.8 g B 500.0 g C 512.5 g D 625.0 g

  33. Density example This pipette is filled with a 20% NaOH solution. The solution is at 20°C and has a density of 1.23 g/mL. According to this information, what is the mass of this NaOH solution? A 3.88 g B 15.7 g C 23.9 g D 24.6 g

  34. Shown are six blocks of various solids. The blocks are the same size but have different masses. A B C D E M= 40 g M = 15 g M = 10 g M = 120 g M = 5 g V= 50 cm3 V= 50 cm3 V= 50 cm3 V= 50 cm3 V= 50 cm3 Rank these situations from greatest to least on the basis of the density of the blocks.

  35. Density and Slope • The slope of a mass versus volume line is density; don’t forget units Density = m/v = ∆y/∆x = y2 – y1 x2 – x1 = 11g-3g 11ml-3ml = 8g/8ml = 1 g/ml

  36. SCIENTIFIC NOTATION • Expresses numbers as multiples of two factors: • a number between 1 and 10 (only 1 digit to the left of the decimal!) • 10 raised to a power, or exponent. The exponent tells how many times the number must be multiplied by 10.

  37. From sci. not. to non-sci. not. • 1.23 X 102 = 123 (move decimal 2 places to right) • 1.23 x 10-2 = .0123 (move decimal 2 places to left)

  38. From non-sci. not. to sci. not. When a number is greater than 1… • the exponent is positive. • (decimal moves to the left) • 1392000 = 1.392 x 106 When a number is less than 1… • the exponent is negative. • (decimal moves to the right) • 0.000000028 = 2.8 x 10-8

  39. The Mole • represents a counted number of things; the number of particles in a substance. • One mole represents 6.02 x 1023 particles. This is called Avogadro’s number. • 1 mole of atoms = 6.02 X 1023 atoms • 1 mole of molecules = 6.02 X1023 molecules

  40. Just how big is a mole? • If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole. • 1 mole of marbles would cover the surface of Earth to a depth of more than 6 kilometers – mole song

  41. Dimensional Analysis Uses a Conversion Factor to move from one unit to another • Start with the original number and unit • Multiply by a conversion factor with the unit you want in the numerator and the unit you started with in the denominator. • Cancel units. • Perform numerical calculations.

  42. Practice hint: dimensional analysis • How many molecules are in 3.00 moles of N2? • How many moles of Na are in 1.10 x 1023 atoms?

  43. Practice hint: dimensional analysis • How many molecules are in 3.00 moles of N2? 3.00 moles X 6.02 X 1023 molecules = 1 Mole • How many moles of Na are in 1.10 x 1023 atoms? 1.1 X 1023 atoms X 1 mole = 6.02 X 1023 atoms

  44. Molar Mass • the mass (think grams) of one mole of a substance • To calculate the molar mass of a compound, you add up the molar masses of all the elements in that compound

  45. Molar Mass Practice • What is the mass of 1.00 mole of Carbon? Of Nitrogen? • Find the molar mass for: • SO3 • Na2SO4

  46. Molar Mass Practice • When you see 1.00 mole = _?_ g, think “g means GO to the PERIODIC TABLE” to find the molar mass. http://www.webelements.com/

  47. Practice • 2.0 mol of H2O = ___g • 456 g of NaCl = ____ mol

  48. Changes in matter • ALL changes involve ENERGY. Physical Changes • Usually involve small amounts of energy • Involve the particles moving closer together or farther apart • Don’t change the identity of the substance

  49. Physical Changes

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