Chapter 1 Introduction to Chemistry

1. Chapter 1 Introduction to Chemistry

2. What is Science? “Natural abilities are like natural plants; they need pruning by study.” – Sir Francis Bacon (1561-1626)

3. The Nature of Science and Chemistry • Definitions • Science: knowledge • Sir Francis Bacon: “And thus knowledge itself is power” • Modern science, the acquisition of knowledge, is acquired by experience (experiment) • Chemistry: the study of matter and its interactions with other matter and with energy.

4. Chemistry and the Natural Sciences

5. 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.

6. Matter • 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.

7. Mass and Weight Mass on moon and earth is the same. Weight on moon and earth is the different.

8. Properties of Matter • 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

9. Physical Properties and Changes • Physical properties: can be measured without changing the composition of the sample. • mass, density, color, melting point • Physical change: a change that occurs without changing the composition of the material. • freezing, melting

10. Chemical Properties • Chemical properties: describe the reactivity of a material. • Natural gas burns in air; iron rusts. • Chemical change: at least part of the material is changed into a different kind of matter. • The digestion of sugar is a chemical change.

11. Practice State if the underlined property or changes is intensive or extensive and chemical or physical. The color of mercury is silvery. The sample of iron rusts by reaction with oxygen. The heat released by burning coal can power a city. Water boils at 100°C. A new pencil is 10 inches long.

12. Classification of Matter • Substances - a material that is chemically the same throughout. • Two types of substances • Elements cannot be broken into simpler substances. • Compounds can be broken down into elements.

13. Substances • Substance: cannot be separated into component parts by physical methods. • Compound: a substance which can be separated into simpler substances by chemical methods. • Element: a substance which cannot be separated into simpler substances by chemical methods.

14. 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. • Homogeneous mixture or solution: composition of the mixture is uniform throughout. • Alloy: a solution of a metal and another material (usually another metal).

15. Classification of Matter

16. Practice Identify the following types of matter as elements, compounds, heterogeneous mixtures, or homogeneous mixtures. • Sodium chloride • Stainless steel • Chlorine • soil

17. Measurement • Most modern science depends on measurements • Parts of a measurement • The object of the measurement • The value of the measurement • The units of the measurement • The reliability of the measurement • Example “The mass of iron was 4.0501 grams • All parts MUST be present in an answer for complete credit!!

18. 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.

19. Accuracy and Precision accurate and precise accurate but not precise precise but not accurate neither accurate nor precise

20. Accuracy and Precision

21. Significant Figures • The number of significant figures is the number of digits from the first non-zero digit through the last reported digit. • The uncertainty is at least ±1 unit in the last reported digit. • Leading zeros – zeros preceding the first non-zero digit are NEVER significant. • Trailing zeros – a decimal point is the key. • No decimal – trailing zeros are NOT significant • Decimal – trailing zeros ARE significant

22. Significant Figures • Quantities that are not limited by significant figures: • counted numbers or tallies. • defined numbers. • the power of ten in exponential notation.

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

24. Significant Figures • Since trailing zeros in numbers without a decimal points may be confusing for significant figures use scientific notation. • 100 1? Or should there have been a decimal?1 x 10211.0 x 10221.00 x 1023

25. Practice • 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

26. Uncertainty in Addition and Subtraction • The absolute uncertainty can be no smaller than the least accurate number. • 12.02 - 10.4 1.62 • 1.6 • The answer should have no more decimal places than the least accurate number.

27. Uncertainty in Multiplication and Division • Answers should have no more significant figures than the least accurate number. • 3121 x 12 = 37452 = 3.7 x 104# sig. digits 4x == NOT 37!!!!!! 37000 is questionable 2 2 2

28. 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 the mathematical operations, # of significant digits are in red. • density = 3.7 g/mL • Be mindful of what your calculator gives you!! 3 3 = = 2 3 3 2

29. Rounding • Be cautious about rounding during multiple steps. • Keep more significant figures than you need in intermediate steps. • Ex: • 2.5 x 4.50 x 11.25 = ?

30. Practice 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)

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

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

33. 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

34. Prefixes Used With SI Units 1 kilogramisequal to 1000 g.

35. 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:

36. Unit Conversions • Lets convert 5.73 g to kg. • Start with what you know!!! • Add conversion factors to cancel units • Units must be same on top & bottom to cancel 0.00573 kg grams on top grams on bottom

37. Practice Convert the following: • 17.43 km to cm • 165 μg to kg

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 Volumes can be expressed in different units depending on the size of the object. 1 m3 contains 1000 L 1 L contains 1000 mL

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

41. 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.

42. 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. What we know Answer: √ number √ units √ object √ sig figs Units cancel

43. English System

44. Practice • Express 323 milliliters in gallons. • Express 3.61 cubic feet in cubic centimeters.

45. Temperature Conversion Factors 0 Kelvin 273 373 -273o Celsius 0o 100o -460o Fahrenheit 32o 212o For water

46. Practice • Express 17.5°C in °F and in K.

47. Practice • It has been estimated that 1.0 g of seawater contains 4.0 pg of Au. The total mass of seawater in the oceans is 1.6x1012 Tg, If all of the gold in the oceans were extracted and spread evenly across the state of Georgia, which has a land area of 58,910 mile2, how tall, in feet, would the pile of Au be? Density of Au is 19.3 g/cm3. 1.0 Tg = 1012g.

48. Practice • One metal object is a cube with edges of 3.00 cm and a mass of 140.4 g. A second metal object is a sphere with radius 1.42 cm and a mass of 61.6 g. Are these objects made of the same or different metals? Assume the calculated densities are accurate to 1.00%.

49. Practice A 40-lb container of peat moss measures 14 x 20 x 30 in. A 40-lb container of topsoil has a volume fo 1.9 gal. • Calculate the density of both the peat moss and the topsoil. • How many bags of peat moss are needed to cover an area measuring 10 ft x 20 ft x 2 in?

50. Practice On a typical day, a hurricane expends the energy equivalent to the explosion of two thermonuclear weapons. A thermonuclear weapon has the explosive power of 1.0 Mton of nitroglycerin. Nitroglycerin generates 7.3 kJ of explosive power per gram of nitroglycerin. The hurricane’s energy comes from the evaporation of water that requires 2.3 kJ per gram of water evaporated. How many gallons of water does a hurricane evaporate per day?