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Chapter 1 Introduction to Chemistry

Chapter 1 Introduction to Chemistry. The Nature of Science and Chemistry. Definitions Science: Study of the natural universe Specifically, knowledge acquired by experience Science is both an activity and the result of the activity.

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Chapter 1 Introduction to Chemistry

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  1. Chapter 1 Introduction to Chemistry

  2. The Nature of Science and Chemistry • 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

  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.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. The Scientific Method • 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.

  8. Matter, Mass and Weight • 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. Mass on moon and earth is the same. Weight on moon and earth is the different.

  9. Properties of Matter • 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

  10. Physical Properties and Changes • 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 Change Chemical Change

  11. Chemical Properties • 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

  12. Classification of Matter - Substances • 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 (H2O is always 11.2% H and 88.8% O)

  13. Classification of Matter - Mixtures • 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).

  14. What’s the Matter?

  15. Accuracy and Precision • 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. accurate but not precise precise but not accurate neither accurate nor precise accurate and 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 visible 0.0012300 But NOT significant if no decimal 12300 Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th 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 • Na = 1 mol = 6.02 x 1023 • π= 3.142 • 1m = 1000 mm

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

  19. Significant Figures • How many significant figures are present in each of the measured quantities? • 0.00122106320064900.041.001250.0010604

  20. Significant Figures • Trailing zeros in numbers without a decimal point may not be significant. Avoid ambiguity by using scientific notation. • 100 1, 2 or 31 x 10211.0 x 10221.00 x 1023

  21. Test Your Skill • Determine the number of significant figures: 100. 100.030505 437,000125,904,000 4.80 x 10-3 4.800 x 10-3 0.0048

  22. Test Your Skill • Determine the number of significant figures 100. 100.030505 437,000125,904,000 4.80 x 10-3 4.800 x 10-3 0.0048 • Answer: 100. 3100.0 4305055437,000 3-6125,904,0006-94.80 x 10-3 34.800 x 10-34 0.0048 2

  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 10n, 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 102 = 1.1 x 10 x 10 = 1.1 x 100 = 110 Decimal  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 × 104 3.45233 × 102 -3.45233 × 102 3.45 × 10-3 1.0345 × 10-1 1.42 × 100

  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). Ex. 01 Ex. 02 Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  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). Ex. 01 Ex. 02 Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

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

  30. Test Your Skill 8.925 - 8.904 x 100% 8.925 9.6 x 100.65 + 4.026 8.321 • Evaluate each expression to the correct number of significant figures. (a) 4.184 x 100.620 x (25.27 - 24.16) (b) (c) Answers: (a) 467; (b) 0.24%; (c) 1.2 x 102

  31. Test Your Skill 9.025 - 9.024 x 100% 9.025 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)

  32. Test Your Skill 9.025 - 9.024 x 100% 9.025 Calculate each to the correct number of significant figures .0.1654 + 2.07 - 2.114 = 0.128.27 x (4.987 - 4.962) = 0.21 9.5 + 4.1 + 2.8 + 3.175 = 4.89 4 (4 is exact) = 0.1%

  33. Math Operations with Scientific Notation Multiplication Division Convert numbers to the same exponents Addition/Subtraction (5.00 x 102) + (6.01 x 103) = (0.500 x 103) + (6.01 x 103) = (5.00 x 102) + (60.10 x 102) = (5.00 + 60.10) x 102 = (65.10 x 102) = 6510 (6.01 x 103) - (5.00 x 102) = (6.01 x 103) - (0.500 x 103) = (60.10 x 102) - (5.00 x 102) = (60.10 - 5.00) x 102 = (55.10 x 102) = 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 102)(5.1 X 104) = (2.7 X 5.1)(102+4) = 13.77 X 106 Now change to Scientific Notation 1.4 X 107 Division a.3.1 X 10-3 = (3.1/1.2)(10-3-2) = 2.6 X 10-5 1.2 X 102 b. 7.9 X 104 = (7.9/3.6)(104-2) = 2.2 X 102 3.6 X 102 Adding/Subtracting 3.05 X 103 + 2.95 X 103 = (3.05 + 2.95)(103) = 6.0 X 103

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

  36. Common Prefixes Used With SI Units PrefixAbbreviationMeaning mega- M 106 kilo- k 103 centi- c 10-2 milli- m 10-3 micro- m 10-6 nano- n 10-9 pico- p 10-12

  37. Unit Conversion Factors • 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: What would be some other examples?

  38. Conversion Among Derived Units • Volume is the product of three lengths. • The standard unit of volume is the cubic meter (m3).100 cm = 1 m(100 cm)3 = (1 m)3106 cm3 = 1 m3 • Two important non-SI units of volume are the liter and milliliter.1 liter (L) = 1000 mL = 1000 cm31 mL = 1 cm3

  39. Volume 1 m3 contains 1000 L 1 L contains 1000 mL 1 mL = 1 cm3 or 1 cc Volumes can be expressed in different units depending on the size of the object.

  40. Using Unit Conversions • Express a volume of 1.250 L in mL, cm3, and m3

  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, 7th 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, 7th Edition, 2011

  43. Density • Density: mass per unit volume • Density, in SI base units, is kg/m3(kg m-3). • Most commonly used density units are g/cm3 (g cm-3 or g/mL) for solids and liquids, and g/L for gases.

  44. Conversions Between Equivalent Units • The density of Ti is 4.50 g/cm3 or 4.50 g = 1 cm3 • Calculate the volume of 7.20 g Ti.

  45. Conversions Between Equivalent Units • The density of Ti is 4.50 g/cm3 or 4.50 g = 1 cm3. • Calculate the volume of 7.20 g Ti.

  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, 7th 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, 7th Edition, 2011

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