1. Chemistry: The Study of Change Chapter 1

2. 1.2 The Study of Chemistry MacroscopicMicroscopic

3. Matter is anything that occupies space and has mass. Chemistry is the study of matter and the changes it undergoes. 1.4 Sugar Water Gold What about air? Yes it is matter

4. MAJOR AREAS OF CHEMISTRY Organic Chemistry - the study of matter which is carbon- based. Inorganic Chemistry - the study of matter containing all other elements (Inorganic is everything else). Analytical Chemistry - analyze matter to determine identity and composition (involves qualitative and quantitative) (eg. Determination of Alcohol in blood) Biochemistry - the study of life at the molecular level (DNA sequencing) Physical Chemistry - attempts to explain the way matter behaves (eg. Kinetics)

5. Chemistry Related Other Fields Medical Science Pharmaceutical Industry Oil Industry Paints and Coatings

6. Chemistry: A Science for the 21 st Century Health and Medicine Sanitation systems Surgery with anesthesia Vaccines and antibiotics Energy and the Environment Fossil fuels Solar energy Nuclear energy

7. Chemistry: A Science for the 21 st Century Materials and Technology Polymers, ceramics Food and Agriculture Genetically modified crops Specialized fertilizers

8. An element is a substance that cannot be separated into simpler substances by chemical means (Listed in the Periodic Table) 114 elements have been identified 82 elements occur naturally on Earth gold, aluminum, lead, oxygen, carbon 32 elements have been created by scientists technetium, americium, seaborgium Metals, nonmetals, metalloids Latin names: aurum (Au), ferrum (Fe),natrium (Na)

10. Classification of Matter 1.Based on Physical State: a.Gas: does not have a fixed shape or a fixed volume. Both the volume & the shape of the gas is that of the container. b.Liquid: Has a fixed volume but variable shape = shape of the container. c.Solid: Has fixed shape and volume

11. The Three States of Matter (water in three forms) solid liquid gas

12. Pure Substances Classification of Pure Substances 1.Elements: These only have 1 type of atoms. 2.Compounds: composed of atoms of two or more elements. eg. H 2 O chemically united in a fixed proportion – 2 H atoms and 1 O atom. It can be separated only chemically. Ammonia (NH 3 ) – 1N, 3H (combined at high pressure) Elements exist in monoatomic or polyatomic form. For eg., Hydrogen and oxygen exist as H 2 and O 2, that is diatomic molecules where as Ca exist as monoatomic. Chemical Symbols: Cobalt – Co, Sodium – Na (natium-Latin) CO is not cobalt, carbon monoxide

13. A mixture is a combination of two or more substances in which the substances retain their distinct identities. 1. Homogenous mixture – composition of the mixture is the same throughout. 2. Heterogeneous mixture – composition is not uniform throughout. soft drink, milk, sugar solution cement, iron filings in sand

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

15. Physical and Chemical Properties Physical Properties: will keep the identity of the substance. Eg. boiling point (B.P), melting point (M.P), density, mass, volume, area. (water changes the physical state by heating and cooling, not the composition, so the M.P and B.P are physical properties) Chemical Properties: involves chemical change. eg. rusting, burning of H in O H 2 O Needs a chemical change to bring the original substances, not physical like boiling or melting.

16. An extensive property of a material depends upon how much matter is being considered - additive An intensive property of a material does not depend upon how much matter is is being considered. e.g., density of water 0.99713g/mL at 25  C mass (M) length volume (V) density – M/V color Extensive and Intensive Properties

17. Hypothesis, Theory, Law A hypothesis is an educated guess, based on observation. Needs to prove by experiments to be true.hypothesis A scientific theory summarizes a hypothesis or group of hypotheses that have been supported with repeated testing.theory A law generalizes a body of observations. Eg., Newton's Law of Gravity F = Gm 1 m 2 /r 2

18. Classifications of Matter

19. Measurement in Chemistry In SI units, not in English units. Meaurement is always necessary to collect data Mass, volume etc. with proper equipments or glassware Always use unit. Mass = 5 for a salt is useless. Needs unit like g or kg. –A measurement is useless without its units. –weight – force that gravity exerts on an object

20. International System of Units (SI) Revised Metric System

21. Truly systematic -1 meter = 10 decimeters = 100 centimeters Basic Units of the Metric System Massgramg Lengthmeter m volumeliterL

23. Matter - anything that occupies space and has mass. mass – measure of the quantity of matter SI unit of mass is the kilogram (kg) 1 kg = 1000 g = 1 x 10 3 g In Chemistry, the smaller g is more convenient. weight – force that gravity exerts on an object

24. Volume – SI derived unit for volume is cubic meter (m 3 ) 1 cm 3 = (1 x 10 -2 m) 3 = 1 x 10 -6 m 3 1 dm 3 = (1 x 10 -1 m) 3 = 1 x 10 -3 m 3 1 L = 1000 mL = 1000 cm 3 = 1 dm 3 1 mL = 1 cm 3

25. Density – SI derived unit for density is kg/m 3 1 g/cm 3 = 1 g/mL = 1000 kg/m 3 Chemical Application: g/cm 3 or g/mL density = mass volume d = m V A piece of platinum metal with a density of 21.5 g/cm 3 has a volume of 4.49 cm 3. What is its mass? d = m V m = d x V = 21.5 g/cm 3 x 4.49 cm 3 = 96.5 g

26. Derived Units in SI 1.Area: (length x width); unit = m x m = m 2 2. Volume: length x width x height; unit = m x m x m = m 3 3.Density: mass/volume; unit = kg/m 3 4.Speed: distance/time; unit = m/s 5.Acceleration: speed/time; unit = m/s/s = m/s 2 6.Force: mass x acceleration; unit = kg m/s 2 = N (newton) 7.Energy: force x distance; unit = kg m/s 2 x m = kg m 2 /s 2 = J (joule) 8.Pressure: force/Area; unit = kg m/s 2 ÷ m 2 = kg/ms 2 = Pa (Pascal)

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

28. Conversions between Fahrenheit and Celsius 1. Convert 75 o C to o F. 2. Convert -10 o F to o C. 1. Ans. 167 o F 2. Ans. -23 o C

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

30. Chemistry In Action On 9/23/99, $125,000,000 Mars Climate Orbiter entered Mar’s atmosphere 100 km (62 miles) lower than planned and was destroyed by heat (read p.17, ch.1) 1 lb = 1 N 1 lb = 4.45 N F=ma=4.45Newton “This is going to be the cautionary tale that will be embedded into introduction to the metric system in elementary school, high school, and college science courses till the end of time.”

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

32. Scientific Notation 568.762 n > 0 568.762 = 5.68762 x 10 2 move decimal left 0.00000772 n < 0 0.00000772 = 7.72 x 10 -6 move decimal right Addition or Subtraction 1.Write each quantity with the same exponent n 2.Combine N 1 and N 2 3.The exponent, n, remains the same 4.31 x 10 4 + 3.9 x 10 3 = 4.31 x 10 4 + 0.39 x 10 4 = 4.70 x 10 4

33. Scientific Notation Multiplication 1.Multiply N 1 and N 2 2.Add exponents n 1 and n 2 (4.0 x 10 -5 ) x (7.0 x 10 3 ) = (4.0 x 7.0) x (10 -5+3 ) = 28 x 10 -2 = 2.8 x 10 -1 Division 1.Divide N 1 and N 2 2.Subtract exponents n 1 and n 2 8.5 x 10 4 ÷ 5.0 x 10 9 = (8.5 ÷ 5.0) x 10 4-9 = 1.7 x 10 -5

34. Scientific Notation The measuring device determines the number of significant figures a measurement has. In this section you will learn –to determine the correct number of significant figures (sig figs) to record in a measurement –to count the number of sig figs in a recorded value –to determine the number of sig figs that should be retained in a calculation.

35. Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit.

36. Significant Figures Any digit that is not zero is significant 1.234 kg 4 significant figures Zeros between nonzero digits are significant 606 m 3 significant figures Zeros to the left of the first nonzero digit are not significant 0.08 L 1 significant figure If a number is greater than 1, then all zeros to the right of the decimal point are significant 2.0 mg 2 significant figures If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant 0.00420 g 3 significant figures

37. How many significant figures are in each of the following measurements? 24 mL2 significant figures 3001 g 4 significant figures 0.0320 m 3 3 significant figures 6.4 x 10 4 molecules 2 significant figures 560 kg (5.6 x 10 2 ) 2 significant figures

38. Significant Figures Addition or Subtraction The answer cannot have more digits to the right of the decimal point than any of the original numbers. 89.332 1.1+ 90.432 round off to 90.4 one digit after decimal point 3.70 -2.9133 0.7867 two digit after decimal point round off to 0.79

39. Scientific Notation 1.8 568.762 n > 0 568.762 = 5.68762 x 10 2 move decimal left 0.00000772 n < 0 0.00000772 = 7.72 x 10 -6 move decimal right Addition or Subtraction 1.Write each quantity with the same exponent n 2.Combine N 1 and N 2 3.The exponent, n, remains the same 4.31 x 10 4 + 3.9 x 10 3 = 4.31 x 10 4 + 0.39 x 10 4 = 4.70 x 10 4

40. Significant Figures Multiplication or 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= 16.5 3 sig figsround to 3 sig figs 6.8 ÷ 112.04 = 0.0606926 2 sig figsround to 2 sig figs = 0.061

41. Significant Figures Numbers from definitions or numbers of objects are considered to have an infinite number of significant figures The average of three measured lengths; 6.64, 6.68 and 6.70? 6.64 + 6.68 + 6.70 3 = 6.67333 = 6.67 = 7

42. Accuracy – how close a measurement is to the true value Precision – how close a set of measurements are to each other accurate & precise but not accurate & not precise

43. UNIT CONVERSION You need to be able to convert between units -within the metric system -between the English and metric system The method used for conversion is called the Factor- Label Method or Dimensional Analysis ALWAYS NEEDED

44. Dimensional Analysis Method of Solving Problems 1.Determine which unit conversion factor(s) are needed 2.Carry units through calculation 3.If all units cancel except for the desired unit(s), then the problem was solved correctly. given quantity x conversion factor = desired quantity given unit x = desired unit desired unit given unit

45. We use these two mathematical facts to do the factor label method –a number divided by itself = 1 –Example: 2/2 = 1 –any number times one gives that number back –Example: 2 x 1 = 2 Example: How many donuts are in 3.5 dozen? You can probably do this in your head but let’s see how to do it using the Factor-Label Method.

46. Start with the given information... 3.5 dozen Then set up your unit factor... See that the units cancel... Then multiply and divide all numbers... = 42 donuts

47. Dimensional Analysis Method of Solving Problems Conversion Unit 1 L = 1000 mL 1L 1000 mL 1.63 L x = 1630 mL 1L 1000 mL 1.63 L x = 0.001630 L2L2 mL How many mL are in 1.63 L?

48. The speed of sound in air is about 343 m/s. What is this speed in miles per hour? 1 mi = 1609 m1 min = 60 s1 hour = 60 min 343 m s x 1 mi 1609 m 60 s 1 min x 60 min 1 hour x = 767 mi hour meters to miles seconds to hours conversion units

49. Common Relationships Used in the English System A. Weight1 pound = 16 ounces 1 ton = 2000 pounds B. Length 1 foot = 12 inches 1 yard = 3 feet 1 mile = 5280 feet C. Volume1 gallon = 4 quarts 1 quart = 2 pints 1 quart = 32 fluid ounces Commonly Used “Bridging” Units for Intersystem conversion QuantityEnglishMetric Mass1 pound=454 grams 2.2 pounds=1 kilogram Length1 inch=2.54 centimeters 1 yard= 0.91 meter Volume1 quart=0.946 liter 1 gallon=3.78 liters

50. 1.3 Measurement in Chemistry 1.Convert 5.5 inches to millimeters (139.7 mm) 2.Convert 50.0 milliliters to pints (0.1057 pint) 3.Convert 1.8 in 2 to cm 2 (11.613 cm 2)

51. 1.7

52. where d is the density, m is the mass, and V is the volume. Generally the unit of mass is the gram. The unit of volume is the mL for liquids; cm 3 for solids; and L for gases. Derived Units The density of an object is its mass per unit volume,

53. A Density Example A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm 3. What is the density of galena?

54. A Density Example A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm 3. What is the density of galena? Density = mass volume

55. A Density Example A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm 3. What is the density of galena? Density = mass volume = 12.4 g 1.64 cm 3

56. A Density Example A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm 3. What is the density of galena? Density = mass volume = 12.4 g 1.64 cm 3 = 7.5609 = 7.56 g/cm 3

57. liquid mercury brass nut water cork

58. What is the density of a sample of bone with mass of 12.0 grams and volume of 5.9 cm 3 ? A sinker of lead has a volume of 0.25 cm 3. Calculate the mass in grams. The density of lead is 11.3 g/cm 3. What is the volume of air in liters (density = 0.00129 g/mL) occupied by 1.0 grams.

59. WORKED EXAMPLES

60. Worked Example 1.1

61. Worked Example 1.2

63. Worked Example 1.4

68. Homework Problems 1.9,1.11,1.12,1.14,1.15,1.16,1.18,1.21,1.22,1.23, 1.29,1.33,1.38,1.39,1.40,1.51,1.52,1.64,1.74, 1.80