Is the study of matter, its properties, the changes that matter undergoes, and

1. CHEMISTRY Is the study of matter, its properties, the changes that matter undergoes, and the energy associated with these changes. 1-

2. Matter anything that has mass and volume -the “stuff” of the universe: books, planets, trees, teachers, students Properties the characteristics that give each substance a unique identity Definitions Physical Properties those which the substance shows by itself without interacting with another substance such as color, melting point, boiling point, density Chemical Properties those which the substance shows as it interacts with, or transforms into, other substances such as flammability, corrosiveness 1-

3. Physical change A substance alters its physical form, not its composition Chemical change A substance is converted into a different substance Figure 1.1 The distinction between physical and chemical change. 1-

4. Table 1.1 Some Characteristic Properties of Copper Physical Properties Chemical Properties slowly forms a basic blue-green sulfate in moist air reddish brown, metallic luster easily shaped into sheets (malleable) and wires (ductile) reacts with nitric acid and sulfuric acid good conductor of heat and electricity density = 8.95 g/cm3 slowly form a deep-blue solution in aqueous ammonia melting point = 10830C boiling point = 25700C 1-

5. Figure 1.2 The physical states of matter. 1-

6. PROBLEM: Decide whether each of the following process is primarily a physical or a chemical change, and explain briefly: Sample Problem 1.1 Distinguishing Between Physical and Chemical Change (a) Frost forms as the temperature drops on a humid winter night. (b) A cornstalk grows from a seed that is watered and fertilized. (c) Dynamite explodes to form a mixture of gases. (d) Perspiration evaporates when you relax after jogging. (e) A silver fork tarnishes slowly in air. SOLUTION: (a) physical change (b) chemical change (c) chemical change (d) physical change (e) chemical change 1-

7. Energy is the capacity to do work. Potential Energy energy due to the position of the object or energy from a chemical reaction Kinetic Energy energy due to the motion of the object Potential and kinetic energy can be interconverted. 1-

8. Energyis the capacity to do work. Figure 1.3B less stable change in potential energyEQUALS kinetic energy more stable A system of two balls attached by a spring. The potential energy gained by a stretched spring is converted to kinetic energy when the moving balls are released. 1-

9. Energy is the capacity to do work. Figure 1.3C less stable change in potential energyEQUALS kinetic energy more stable A system of oppositely charged particles. The potential energy gained when the charges are separated is converted to kinetic energy as the attraction pulls these charges together. 1-

10. How do we study chemistry? Chemistry is an experimental science: base on the scientific method.

11. Hypothesis: Tentative proposal that explains observations. revised if experiments do not support it Procedure to test hypothesis; measures one variable at a time. Experiment: Set of conceptual assumptions that explains data from accumulated experiments; predicts related phenomena. Theory(Model): altered if predictions do not support it Further Experiment: Tests predictions based on model. Scientific Approach: Developing a Model Observations : Natural phenomena and measured events; universally consistent ones can be stated as a natural law. 1-

12. Alchemist at Work 1-

13. Lavoisier(1743 – 1794) • Debunked phlogiston theory • Demonstrated the true nature of combustion • Named oxygen 1-

14. A Systematic Approach to Solving Chemistry Problems • Problem statement Clarify the known and unknown. • Plan Suggest steps from known to unknown. Prepare a visual summary of steps. • Solution • Check Comment and Follow-up Problem 1-

15. PROBLEM: To wire your stereo equipment, you need 325 centimeters (cm) of speaker wire that sells for $0.15/ft. What is the price of the wire? Sample Problem 1.2 Converting Units of Length PLAN: Known - length (in cm) of wire and cost per length ($/ft) We have to convert cm to inches and inches to ft followed by finding the cost for the length in ft. length (cm) of wire Follow up Problem 1.2 A furniture factory needs 31.5 ft2 of fabric to upholster one chair. The supplier sends the fbric in bolts of exactly 200 m2. What is the maximum number of chairs that can be upholstered by 3 bolts of fabric? 1 m = 3.281 ft length (in) of wire 2.54 cm = 1 in length (ft) of wire 12 in = 1 ft Price ($) of wire 1 ft = $0.15 1-

16. SI Base Units Table 1. 2 time second s temperature kelvin K electric current ampere A amount of substance mole mol luminous intensity candela cd Physical Quantity (Dimension) Unit Abbreviation Unit Name mass kilogram kg length meter m 1-

17. What is the metric system? • The first standardized system of measurement, based on the decimal was proposed in France about 1670. • It was created to develop a unified, natural, universal system of measurement. • In 1790 King Louis XVI of France assigned a group to begin this task. • At that time, every country had their own system of weights and measures. England had three different systems just within its own borders!! • It was called the "metric" system, based on the French word for measure. • As of 2005, only three countries, the United States, Liberia, and Myanmar, have not changed over to the metric system. • The official modern name of the metric system is the International System of Units or abbreviated SI.

18. What are the metric measurements that we are learning about? • Meter • Length • Liter • Volume • Gram • Weight or Mass

19. How to remember different lengths in the metric system: • When you think of a millimeter (mm) think of: • The thickness of a dime. • When you think of a centimeter (cm) think of: • The width of your pinky.

20. How to remember different lengths in the metric system: • When you think of a meter (m) think of: • The height of the doorknob. • When you think of a kilometer (km) think of: • A little more than half of a mile

21. How to understand the ruler: • One side has inches (duh!) • The other side is the metric side. • The big numbers represent centimeters • The smaller lines represent millimeters. • How many millimeters are in a centimeter? • 10

22. How to remember different volumes in the metric system: • When you think of a liter (L) think of: • About the size of a bottle of water. • When you think of 5 milliliters (mL) think of: • One teaspoon • When you think of 2 kiloliters (kL) think of: • A hot tub

23. How to remember different masses in the metric system: • When you think of a gram (g) think of: • A paper clip • When you think of a kilogram (kg) think of: • A little more than 2 pounds.

24. What is the order of the metric system? • King Henry Died by Drinking Chocolate Milk • King: Kilo • Henry: Hecto • Died: Deca • By: Base (m, L, g) • Drinking: Deci • Chocolate: Centi • Milk: Milli

25. Metric Prefixes • Metric Units • The metric system has prefix modifiers that are multiples of 10.

26. Place Values of Metric Prefixes

27. 1 m = 100 cm 1 cm = 10 mm 1 m = 1,000 mm 1 km = 1,000 m 1 g = 100 cg 1 cg = 10 mg 1 g = 1,000 mg 1 kg = 1,000 g Measurement Conversions to Know: • 1 L = 100 cL • 1 cL = 10 mL • 1 L = 1,000 mL • 1 kL = 1,000 mL

28. Metric Conversion Challenge Write the correct abbreviation for each metric unit. 1) Kilogram _____ 4) Milliliter _____ 7) Kilometer _____ 2) Meter _____ 5) Millimeter _____ 8) Centimeter _____ 3) Gram _____ 6) Liter _____ 9) Milligram _____ Try these conversions. 10) 2000 mg = _______ g 15) 5 L = _______ mL 20) 16 cm = _______ mm 11) 104 km = _______ m 16) 198 g = _______ kg 21) 2500 m = _______ km 12) 480 cm = _____ m 17) 75 mL = _____ L 22) 65 g = _____ mg 13) 5.6 kg = _____ g 18) 50 cm = _____ m 23) 6.3 cm = _____ mm 14) 8 mm = _____ cm 19) 5.6 m = _____ cm 24) 120 mg = _____ g

29. Table 1.3 Common Decimal Prefixes Used with SI Units 1-

30. Densities of Some Common Substances* Table 1.5 Hydrogen Gas 0.0000899 Oxygen Gas 0.00133 Grain alcohol Liquid 0. 789 Water Liquid 0.998 Table salt Solid 2.16 Aluminum Solid 2.70 Lead Solid 11.3 Gold Solid 19.3 Substance Physical State Density (g/cm3) *At room temperature(200C) and normal atmospheric pressure(1atm). 1-

31. Derived Units Area (m2), volume (m3), density (kg/m3), speed (m/s) etc are some of the derived quantities that you are familiar with. They are derived quantities rather than fundamental quantities because they can be expressed using one or more of the seven base units. Volume, the amount of space occupied by an object, is measured in SI units by the cubic meter (m3), defined as the amount of space occupied by a cubic 1 meter long on each edge.

32. Density Density is the intensity properties that relates the mass of an object to its volume. Density = Mass (g)_______ Volume (mL or cm3) Because most substances change in volume when heated or cooled, densities are temperature-dependent. E.g. the density of water at 3.98C is 1.000g /mL and at 100C is 0.9584 g/mL as the volume expand. Although most substances expand when heated and contract when cooled, water behaves differently. Water contract when cooled from 3.98C to 0C.

33. Sample Problem 1.5 Calculating Density from Mass and Length PROBLEM: Lithium (Li) is a soft, gray solid that has the lowest density of any metal. If a slab of Li weighs 1.49 x 103 mg and has sides that measure 20.9 mm by 11.1 mm by 11.9 mm, what is the density of Li in g/cm3 ? 1-

34. The freezing and boiling points of water. Figure 1.12 1-

35. Temperature Scales and Interconversions Kelvin ( K ) - The “Absolute temperature scale” begins at absolute zero and only has positive values. Celsius ( oC ) - The temperature scale used by science, formally called centigrade, most commonly used scale around the world; water freezes at 0oC, and boils at 100oC. Fahrenheit ( oF ) -Commonly used scale in the U.S. for our weather reports; water freezes at 32oF and boils at 212oF. Kelvin = oC + 273.15 oC = Kelvin - 273.15 oF = (9/5) oC + 32 oC = [oF - 32 ] 5/9 1-

36. Sample Problem 1.6 Converting Units of Temperature A child has a body temperature of 38.7 oC. (a) If normal body temperature is 98.6 oF, does the child have a fever? (b) What is the child’s temperature in Kelvin? 1-

37. Rules for Determining Which Digits are Significant Leading zeros are not significant. • If the measured quantity has a decimal point start at the left of the number and move right until you reach the first nonzero digit. • Count that digit and every digit to it’s right as significant. Zeros that end a number and lie either after or before the decimal point are significant; thus 1.030 ml has four significant figures, and 5300. L has four significant figures also. Numbers such as 5300 L are assumed to only have 2 significant figures. A terminal decimal point (or a bar) is often used to clarify the situation, but scientific notation is the best! • If the measured quantity does not have a decimal point start at the right of the number and move leftt until you reach the first nonzero digit. • Count that digit and every digit to it’s left as significant. 1-

38. (a)2sf (b) 4sf (c) 5sf (d)4sf (e) 5sf (f) 4sf Sample Problem 1.7 Determining the Number of Significant Figures For each of the following quantities, determine the number of significant figures in each quantity. (a) 0.0030 L (b) 0.1044 g (c) 53,069 mL (d) 0.00004715 m (e) 57,600. s (f) 0.0000007160 cm3 SOLUTION: 1-

39. Figure 1.14 The number of significant figures in a measurement depends upon the measuring device. 32.33 oC 32.3 oC 1-

40. Rules for Rounding Off Numbers 1. If the digit removed is more than 5, the preceding number increases by 1. 5.379 rounds to 5.38 if three significant figures are retained and to 5.4 if two significant figures are retained. 2. If the digit removed is less than 5, the preceding number is unchanged. 0.2413 rounds to 0.241 if three significant figures are retained and to 0.24 if two significant figures are retained. 3.If the digit removedis 5, the preceding number increases by 1 if it is odd and remains unchanged if it is even. 17.75 rounds to 17.8, but 17.65 rounds to 17.6. If the 5 is followed only by zeros, rule 3 is followed; if the 5 is followed by nonzeros, rule 1 is followed: 17.6500 rounds to 17.6, but 17.6513 rounds to 17.7 4. Be sure to carry two or more additional significant figures through a multistep calculation and round off only the finalanswer. 1-

41. Precision and Accuracy Errors in Scientific Measurements Precision - Refers to reproducibility orhow close the measurements are to each other. Accuracy - Refers to how close a measurement is to the real value. Systematic error - Values that are either all higher or all lower than the actual value. Random Error - In the absence of systematic error, some values that are higher and some that are lower than the actual value. 1-

42. Uncertainties in Scientific Measurement • All measurements are subject to error. • Types of Errors: • 1)   Systematic error: arise because to some extent, measuring instruments have built-in, or inherent, errors. • Random errors:they arise from intrinsic limitation in the sensitivity of the instrument and inability of observer to read a scientific instrument and give results that may be either too high or too low. • In talking about the degree of uncertainty in a measurement, we use the words accuracy and precision. • There is no relationship between accuracy and precision, since an experiment can have small random errors and still give inaccurate results due to large systematic errors.

43. precise and accurate precise but not accurate Figure 1.16 Precision and accuracy in the laboratory. 1-

44. random error systematic error Precision and accuracy in the laboratory. Figure 1.16 continued 1-

45. Dimensional Analysis • Dimensional analysis is the method of calculation in which one carries along the units for quantities. Suppose we want to find the volume (V) of a cube, give l, the length of a side of the cube. Because V = l3, if l = 5.00cm, we find that V = (5cm)3 = 125cm3. There is no guesswork about the unit of volume here; it is cubic centimeter (cm3). Suppose, however, that we wish to express the volume in liter (L), a metric unit that equals 103 cubic centimeters. We can write this equality as • 1L = 103 cm3. • If we divide both sides of the equality by the right-hand quantity, we get • Observe that units are treated in the same way as algebraic quantities. Note too that the right-hand side now equals 1 and there are no units associated with it. Because it is always possible to multiply any quantity by 1 without changing that quantity, we can multiply our previous expression for volume by the factor 1 L/103 without changing the actual volume. We are changing only the way in which we express this volume:

46. The ratio 1 L/ 103 cm3 is called a conversion factor because it is a factor equal to 1 that converts a quantity expressed in one unit to one expressed in another unit. • E.g. Convert 8.45 kg to milligrams • 1 kg = 106mg • E.g. What is the density of a substance in g/mL, if a sample with a volume of 0.085 liters has a mass of 1700 mg?

47. What are we studying in chemistry? Chemistry is the science that studies the composition and properties of matter. Matteris anything that occupies space, displays a property known as mass, and possesses inertia.Matter is the name scientists have given to everything that you can touch, or see, or feel Composition refers to the parts or components of a sample of matter and their relative proportions. Any characteristic that can be used to describe or identify matter is called a property. We classify matter (1) by its physical state as a solid, liquid, or gas, and (2) by its chemical constitution as an element, compound, or mixture.

48. Properties • Properties can be classified as • Intensive or extensive, depending on whether their value changes with the size of the sample. Intensive properties, like temperature and melting point, have values that do not depend on the amount of sample. • Extensive properties, like length and volume, have values that do depend on the sample size. • 2)Physical or chemical. • ~ Physical properties and physical change • A physical property, are those characteristics like temperature, color and melting point, is one that a sample of matter displays without changing its composition. • In physical change some of the physical properties of a sample may change, but its composition remains unchanged. • E.g. Liquid water and ice (solid water) certainly different in many way but water remains 11.9% hydrogen and 88.81% oxygen by mass.

49. Physical Properties Chemical Properties Temperature Amount Color Odor Melting point Solubility Electrical conductivity Hardness Rusting (of iron) Combustion (of coal) Tarnishing (of silver) Hardening (of cement) ~ Chemical properties and chemical changeA chemical property is the ability (or inability) of a sample of matter to undergo a change in composition under stated condition. In a chemical change (or chemical reaction), one of more kinds of matter are converted to new kinds with different compositions. E.g. the rust that occurs when a bicycle is left out in the rain is due to the chemical combination of oxygen with iron to give the new substance iron oxide. Rust is therefore a chemical properties of iron. Table 1 Some Examples of Physical and Chemical Properties

50. Classification of Matter