1. Daniel L. Reger Scott R. Goode David W. Ball www.cengage.com/chemistry/reger Chapter 1 Introduction to Chemistry

2. Definitions Science: Study of the natural universe Specifically, knowledge acquired by experience Science is both an activity and the result of the activity. Chemistry: the study of matter and its interactions with other matter and with energy. Chemistry is how matter is organized/reorganized/changed at the molecular level The Nature of Science and Chemistry

3. Chemistry: The Central Science Chemistry is often called the central science because it is an essential component of the natural and life sciences.

4. Chemistry and Astronomy Elemental composition of stars can be determined by different wavelengths of visible light emitted. When starlight passes through a planets atmosphere, certain frequencies of light disappear because they are absorbed by compounds in the atmosphere. http://cs.fit.edu/~wds/classes/cse5255/cse5255/davis/text.html

5. Chemistry and Geology Geochemistry: Study of the chemical composition of the earth Chemical transformations in solids Ex. Polymorphism. How limestone becomes marble. http://geology.com/rocks/limestone.shtmlhttp://geology.com/rocks/limestone.shtml; http://www.italartworld.com/

6. Chemistry and Biology You reach a certain level in biology where processes can only be understood in terms of chemistry. Chemistry in biology explains: Why your adrenaline levels increase when you are afraid or excited Why a body fails to produce insulin (diabetes) Why cells become cancerous Neurotransmitter (e.g., dopamine, norepinephrine) imbalances that can produce: Euphoria when you have a few beers or fall in love Depression http://cs.fit.edu/~wds/classes/cse5255/cse5255/davis/text.html

7. Scientific method: investigations that are guided by theory and earlier experiments. Hypothesis: a possible explanation for an event. Law: a statement that summarizes a large number of observations. Theory: an explanation of the laws of nature. In the realm of science, theory has a much narrower and more rigorous meaning than in general. The Scientific Method

8. Matter: anything that has mass and occupies space. Mass: the quantity of matter in an object. Weight: the force of attraction between an object and other objects. Matter, Mass and Weight Mass on moon and earth is the same. Weight on moon and earth is the different.

9. Matter can be described by different properties Property: anything observed or measured about a sample of matter. Extensive property: depends on the size of the sample. mass, volume Intensive property: independent of sample size. density, color, melting or boiling point 2 samples with same intensive properties may be the same material Properties of Matter

10. Physical properties: can be measured without changing the composition of the sample. mass, density, color, MP, BP, solubility Physical change: a change that occurs without changing the composition of the material. freezing, melting, crushing a brick into powder Physical Properties and Changes Physical Change Chemical Change

11. Chemical properties: describe the reactivity of a material. –Flammability (whether something is ignitable, flash point < 100 °F) –Combustibility (whether something will burn, flash point > 100 ° F) Iron rusts Chemical change: at least part of the material is changed into a different kind of matter. Digestion of sugar is a chemical change Acid/base neutralization Chemical Properties

12. Substances - Material that is chemically the same throughout. cannot be separated into component parts by physical methods Two types of substances Elements cannot be broken into simpler substances by chemical methods Table 1.1 (p. 09): Memorize these!!! Compounds can be separated into simpler substances (or elements) by chemical methods Always contain same elements in same proportions (H 2 O is always 11.2% H and 88.8% O) Classification of Matter - Substances

13. Mixture: matter that can be separated into simpler materials by physical methods. Heterogeneous mixture: composition of the mixture changes from one part to another. Chocolate chip cookies Italian dressing Homogeneous mixture or solution: composition of the mixture is uniform throughout. Chocolate pudding Sugar dissolves in water Alloy: a solution of a metal and another material (usually another metal). Classification of Matter - Mixtures

14. What’s the Matter?

15. Modern chemistry is largely based on experimental measurements. The confidence in measurements involves: Accuracy: agreement of a measurement with the true value. Precision: agreement among repeated measurements of the same quantity. Accuracy and Precision accurate and precise precise but not accurate accurate but not precise neither accurate nor precise

16. Significant Figures Significant figures are the numbers in a measurement that represent the certainty of the measurement, plus one number representing an estimate. Q: When is a number NOT significant? A: Look at the zeros Leading zeros are NOT significant.0.00123 Confined zeros ARE significant.0.00103 Trailing zeros ARE significant, when decimal visible0.0012300 But NOT significant if no decimal 12300 Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

17. Calculations with Significant Figures Rules for Rounding If the first nonsignificant figure to drop from your answer is ≥ 5, all nonsignificant figures dropped, last significant figure increased by 1. If the first nonsignificant figure to drop from your answer is < 5, all nonsignificant figures dropped, last significant figure stays the same. Exact Numbers Numbers with no uncertainty, or are known values. Exact numbers do not change. Ex. 1 foot is always = 12 inches. It will never be = 12.5 inches. Not used to determine sig figs N a = 1 mol = 6.02 x 10 23 π= 3.142 1m = 1000 mm

18. How many significant figures are present in each of the measured quantities? 0.0012 106 2006 900.0 1.0012 0.001060 Significant Figures

19. How many significant figures are present in each of the measured quantities? 0.00122 1063 20064 900.04 1.00125 0.0010604 Significant Figures

20. Trailing zeros in numbers without a decimal point may not be significant. Avoid ambiguity by using scientific notation. 1001, 2 or 3 1 x 10 2 1 1.0 x 10 2 2 1.00 x 10 2 3 Significant Figures

21. Determine the number of significant figures: 100.100.0 30505437,000 125,904,0004.80 x 10 -3 4.800 x 10 -3 0.0048 Test Your Skill

22. Determine the number of significant figures 100.100.0 30505437,000 125,904,0004.80 x 10 -3 4.800 x 10 -3 0.0048 Answer: 100.3100.04 305055437,0003-6 125,904,0006-94.80 x 10 -3 3 4.800 x 10 -3 40.00482 Test Your Skill

23. Scientific Notation Scientific notation provides a convenient way to express very large or very small numbers. Numbers written in scientific notation consist of a product of two parts in the form M x 10 n, where M is a number between 1 and 10 (but not equal to 10) and n is a positive or negative whole number. The number M is written with the decimal in the standard position.

24. Scientific Notation (continued) STANDARD DECIMAL POSITION The standard position for a decimal is to the right of the first nonzero digit in the number M. SIGNIFICANCE OF THE EXPONENT n A positive n value indicates the number of places to the right of the standard position that the original decimal position is located. A negative n value indicates the number of places to the left of the standard position that the original decimal position is located.

25. Scientific ↔ Standard Notation Converting from scientific notation to standard numbers 1.1 x 10 2 = 1.1 x 10 x 10 = 1.1 x 100 = 110Decimal  1.1 x 10 -2 = 1.1/ (10 x 10) = 1.1/100 = 0.011  Decimal Converting Exponents When you move the decimal (l or r), the exponent will be equal to number of places you moved the decimal.

26. Standard to Scientific Notation 60023.5  345.233  -345.233  0.00345  0.10345  1.42  6.00235 × 10 4 3.45233 × 10 2 -3.45233 × 10 2 3.45 × 10 -3 1.0345 × 10 -1 1.42 × 10 0

27. Calculations with Significant Figures Sum (addition) or difference (subtraction) must contain the same number of places to the right of the decimal (prd) as the quantity in the calculation with the fewest number of places to the right of the decimal (i.e., the least accurate number). Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011 Ex. 01 Ex. 02

28. Calculations with Significant Figures Product (multiplication) or quotient (division) must have same number of sig figs as value with the fewest number of sig figs (least accurate number). Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011 Ex. 01 Ex. 02

29. Determine accuracy in the same order as you complete the mathematical operations, # of significant digits are in red. density = 3.7 g/mL 3 2 2.79 g 8.34 mL - 7.58 mL v m 2.79 g 0.76mL = 3 3 3 2 = Mixed Operations

30. Evaluate each expression to the correct number of significant figures. (a) 4.184 x 100.620 x (25.27 - 24.16) (b) (c) 8.925 - 8.904 x 100% 8.925 9.6 x 100.65 8.321 + 4.026 Test Your Skill Answers: (a) 467; (b) 0.24%; (c) 1.2 x 10 2

31. Calculate each to the correct number of significant figures. 0.1654 + 2.07 - 2.114 8.27 x (4.987 - 4.962) 9.5 + 4.1 + 2.8 + 3.175 4 (4 is exact) x 100% 9.025 - 9.024 9.025 Test Your Skill

32. Calculate each to the correct number of significant figures. 0.1654 + 2.07 - 2.114 = 0.12 8.27 x (4.987 - 4.962) = 0.21 9.5 + 4.1 + 2.8 + 3.175 4 (4 is exact) x 100% 9.025 - 9.024 9.025 Test Your Skill = 4.89 = 0.1%

33. Math Operations with Scientific Notation Multiplication Division Addition/Subtraction Convert numbers to the same exponents (5.00 x 10 2 ) + (6.01 x 10 3 ) = (0.500 x 10 3 ) + (6.01 x 10 3 ) = (5.00 x 10 2 ) + (60.10 x 10 2 ) = (5.00 + 60.10) x 10 2 = (65.10 x 10 2 ) = 6510 (6.01 x 10 3 ) - (5.00 x 10 2 ) = (6.01 x 10 3 ) - (0.500 x 10 3 ) = (60.10 x 10 2 ) - (5.00 x 10 2 ) = (60.10 - 5.00) x 10 2 = (55.10 x 10 2 ) = 5510

34. Examples of Math Operations Multiplication a. (8.2 X 10 -3 )(1.1 X 10 -2 ) = (8.2 X 1.1)(10 (-3+(-2)) ) = 9.0 X 10 -5 b. (2.7 X 10 2 )(5.1 X 10 4 ) = (2.7 X 5.1)(10 2+4 ) = 13.77 X 10 6 Now change to Scientific Notation 1.4 X 10 7 Division a. 3.1 X 10 -3 = (3.1/1.2)(10 -3-2 ) = 2.6 X 10 -5 1.2 X 10 2 b. 7.9 X 10 4 = (7.9/3.6)(10 4-2 ) = 2.2 X 10 2 3.6 X 10 2 Adding/Subtracting 3.05 X 10 3 + 2.95 X 10 3 = (3.05 + 2.95)(10 3 ) = 6.0 X 10 3

35. Quantity Unit Abbreviation Length meterm Mass kilogramkg Time seconds Temperature kelvinK Amount mole mol Electric current ampereA Luminous intensity candelacd Base Units in the SI

36. PrefixAbbreviationMeaning mega-M10 6 kilo-k10 3 centi-c10 -2 milli-m10 -3 micro-  10 -6 nano-n10 -9 pico-p 10 -12 Common Prefixes Used With SI Units

37. Unit conversion factor: a fraction in which the numerator is a quantity equal or equivalent to the quantity in the denominator, but expressed in different units The relationship 1 kg = 1000 g Generates two unit conversion factors: Unit Conversion Factors What would be some other examples?

38. Volume is the product of three lengths. The standard unit of volume is the cubic meter (m 3 ). 100 cm = 1 m (100 cm) 3 = (1 m) 3 10 6 cm 3 = 1 m 3 Two important non-SI units of volume are the liter and milliliter. 1 liter (L) = 1000 mL = 1000 cm 3 1 mL = 1 cm 3 Conversion Among Derived Units

39. Volumes can be expressed in different units depending on the size of the object. 1 m 3 contains 1000 L 1 L contains 1000 mL 1 mL = 1 cm 3 or 1 cc Volume

40. Express a volume of 1.250 L in mL, cm 3, and m 3 Using Unit Conversions

41. Factor Unit Method Examples A length of rope is measured to be 1834 cm. How many meters is this? Solution: Write down known quantity (1834 cm). Set known quantity = units of the unknown quantity (meters). Use factor (100 cm = 1 m), to cancel units of known quantity (cm) and generate units of the unknown quantity (m). Do the math. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

42. Factor Unit Method Example Q: If an arrow shot from a bow travels 30 yards in 1 second, many cm does it travel in 4 seconds? Time = 4 s Rate = 30 yards/sec 1 yard = 3 feet 1 foot = 12 in. 1 in = 2.54 cm Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

43. Density: mass per unit volume Density, in SI base units, is kg/m 3 (kg m -3 ). Most commonly used density units are g/cm 3 (g cm - 3 or g/mL) for solids and liquids, and g/L for gases. Density

44. The density of Ti is 4.50 g/cm 3 or 4.50 g = 1 cm 3 Calculate the volume of 7.20 g Ti. Conversions Between Equivalent Units

45. The density of Ti is 4.50 g/cm 3 or 4.50 g = 1 cm 3. Calculate the volume of 7.20 g Ti. Conversions Between Equivalent Units

46. Percentage The word percentage means per one hundred. It is the number of items in a group of 100 such items. PERCENTAGE CALCULATIONS Percentages are calculated using the equation: In this equation, part represents the number of specific items included in the total number of items. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

47. Percentage Calculation A student counts the money she has left until pay day and finds she has $36.48. Before payday, she has to pay an outstanding bill of $15.67. What percentage of her money must be used to pay the bill? Solution: Her total amount of money is $36.48, and the part is what she has to pay or $15.67. The percentage of her total is calculated as follows:

48. Density Calculation A 20.00 mL sample of liquid is put into an empty beaker that had a mass of 31.447 g. The beaker and contained liquid were weighed and had a mass of 55.891 g. Calculate the density of the liquid in g/mL. Solution: The mass of the liquid is the difference between the mass of the beaker with contained liquid, and the mass of the empty beaker or 55.891g -31.447 g = 24.444 g. The density of the liquid is calculated as follows:

49. Energy Calculations Q: In order to lose 1 lb/week, you need to cut 500.0 Cal from your diet each day, or use an equivalent number of joules by working each day. How many equivalent joules would you have to spend on work to achieve this each day? How many joules would you have to expend to achieve this over 7 days? To answer this, you need to first know the following: 1Cal = 1 kcal = 1000 scientific calories or 1 nutritional calorie 1 scientific calorie = 1 cal = 4.184 J

50. Temperature Scales The three most commonly-used temperature scales are the Fahrenheit, Celsius and Kelvin scales. The Celsius and Kelvin scales are used in scientific work. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

51. Temperature Conversions Readings on one temperature scale can be converted to the other scales by using mathematical equations. Converting Fahrenheit to Celsius. Converting Celsius to Fahrenheit. Converting Kelvin to Celsius. Converting Celsius to Kelvin. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

52. Express 17.5°C in °F and in K. Test Your Skill

53. Express 17.5°C in °F and in K. Answer: T F = 63.5° F; T K = 290.6 K Test Your Skill

54. Useful Conversions/Constants N a = 1 mol = 6.02 x 10 23 π= (Circumference of circle/diameter of circle) = 3.142 1 scientific calorie = 4.184 J 1Cal = 1 kcal = 1000 scientific calories or 1 nutritional calorie 1 N = (kg)(m)/s 2 g = 9.81 m/s 2 (earth’s gravity)

55. Chapter 1 Visual Summary