Honors Chemistry I 84.135

1. Honors Chemistry I84.135 Dr. Nancy De Luca Course web site: http://faculty.uml.edu/ndeluca/84.135

2. Matter Matter is anything that has mass and occupies space. It includes everything around us, including the air that we breath, our skin and bones, and the earth underneath us.

3. Properties of Matter Matter can be described by its physicalor chemicalproperties. Physical properties are a description of the substance, and include mass, color, physical state (solid, liquid or gas) at a specific temperature, density, melting or boiling point, odor, solubility, etc.

4. Properties of Matter Mercury, a metal that is a liquid at room temperature, reacts with iodine, a black shiny solid, to produce mercury (II) iodide, a red crystalline solid. mercury iodine mercury(II)iodide

5. Properties of Matter • During a physical change, the chemical identity of the substance or substances does not change. • Examples of physical changes include evaporation, filtration, and changes of state.

6. Properties of Matter When ice is melted and the liquid is then evaporated, all three forms of water are chemically the same, H2O. When salt water is boiled, the salt remains, and the water is removed as water vapor. For either process, there is no change in the identity of the substances. This is true for all physical changes.

7. Physical Changes When water boils, its chemical composition remains the same. The molecules are now farther apart.

8. Physical Change The dissolving of sugar in water is a physical change. The chemical identity of the water and the sugar remain unchanged.

9. Filtration During filtration, liquids are separated from solids by physical means. The liquid and solid maintain their chemical identity.

10. Distillation During distillation, liquids may be separated from other liquids, or from solids. The chemical identity of each component remains unchanged.

11. Properties of Matter Chemical properties are descriptions of how a substance reacts chemically. Examples include the rusting of iron in the presence of air and water, the souring of milk, or the burning of paper to form carbon dioxide and water vapor.

12. Chemical Changes As iron rusts, the iron atoms combine with oxygen in the air to form a new substance, rust, or iron (III) oxide.

13. Chemical Changes During a chemical change, atoms rearrange the way they are attached to each other, forming new substances with properties that are often quite different from the starting materials.

14. Intensive & Extensive Properties Intensive properties do not depend on the amount or quantity of matter. Melting point, chemical formula and color are intensive properties. Extensive properties depend upon the quantity of matter or sample size. Examples include length, mass and volume.

15. The Elements All matter is composed of approximately 100 elements, in various combinations, listed on the periodic table. The table groups elements with similar chemical and physical properties.

16. The Periodic Table • Periodic tables group elements with similar properties in vertical groups or families. • Metals are on the left side of the table, and non-metals are on the right. • A bold line resembling a flight of stairs usually separated metals from non-metals.

17. The Periodic Table metal/non-metalline

18. Measurements Scientists needed to establish a system of measurement and units before they could reproduce or communicate the results of their experiments. The Metric System is used, with the units of grams (for mass) and milliliters (for volume) commonly used in the chemistry laboratory.

19. Measurements Prefixes Commonly used in Chemistry: prefix namesymbolvalue exponential notation kilo  k 1,000 103 centi c 1/100 or .01 10-2 milli m 1/1,000 or .001 10-3 micro µ .000001 10-6 nano n 10-9 pico p 10-12

20. Measurements- Units • SI or International System units are used. QuantityUnitSymbol Mass kilogram kg Length meter m Time second s Temperature kelvin K

21. Measurement- Volume Volume is a derived unit. A liter is a volume that is 10cm x 10cm x 10cm, or 1000 cm3. Therefore, a milliliter (mL) is the same as a cubic centimeter (cm3 or cc).

22. In the chemistry lab, temperature is measured in degrees Celsius or Centigrade. The temperature in Kelvins is found by adding 273.15 The Fahrenheit scale has 180 oF/100 oC. This is reason for the 5/9 or 9/5 in the conversion formulas. Measurement - Temperature

23. Measurement - Units Common English-Metric Conversion Factors 2.54 cm = 1 inch 1 lb = 453.6 g 1 qt = 943 mL

24. Significant Figures When writing a number, the certainty with which the number is known should be reflected in the way it is written. Digits which are the result of measurement or are known with a degree of certainty are called significant digits or significant figures.

25. Significant Figures The goal of paying attention to significant figures is to make sure that every number accurately reflects the degree of certainty to which it is known. Likewise, when calculations are performed, the final result should reflect the same degree of certainty as the least certain quantity in the calculation.

26. Significant Figures If someone says “There are roughly a hundred students enrolled in the freshman chemistry course,” the enrollment should be written as 100 or 1 x 102. Either notation indicates that the number is approximate, with only one significant figure.

27. Significant Figures If the enrollment is exactly one hundred students, the number should be written with a decimal point, as 100. , or 1.00 x 102. Note that in either form, the number has three significant figures.

28. Significant Figures The rules for counting significant figures: 1. Any non-zero integer is a significant figure.

29. Significant Figures 2. Zeros may be significant, depending upon where they appear in a number. a) Leading zeros (one that precede any non-zero digits) are not significant. For example, in 0.02080, the first two zeros are not significant. They only serve to place the decimal point.

30. Significant Figures – Zeros (cont’d) b) Zeros between non-zero integers are always significant. In the number 0.02080, the zero between the 2 and the 8 is a significant digit. c) Zeros at the right end of a number are significant only if the number contains a decimal point. In the number 0.02080, the last zero is the result of a measurement, and is significant.

31. Significant Figures Thus, the number 0.02080 has four significant figures. If written in scientific notation, all significant digits must appear. So 0.02080 becomes 2.080 x 10-2.

32. Significant Figures 3. Exact numbers have an unlimited number of significant figures. Examples are 100cm = 1m, the “2” in the formula 2πr, or the number of atoms of a given element in the formula of a compound, such as the “2” in H2O. Using an exact number in a calculation will not limit the number of significant figures in the final result.

33. Significant Figures - Calculations When calculations are performed, the final result should reflect the same degree of certainty as the least certain quantity in the calculation. That is, the least certain quantity will influence the degree of certainty in the final result of the calculation.

34. Significant Figures - Calculations There are two sets of rules when performing calculations. One for addition and subtraction, and the other for multiplication and division. • For Multiplication and Division: The result of the calculation should have the same number of significant figures as the least precise measurement used in the calculation.

35. Significant Figures - Calculations Multiplication & Division: Example: Determine the density of an object with a volume of 5.70 cm3 and a mass of 8.9076 grams.

36. Significant Figures - Calculations Multiplication & Division: Example: Determine the density of an object with a volume of 5.70 cm3 and a mass of 8.9076 grams. δ = mass/volume = 8.9076 g/5.70 cm3 δ = 1.5627368 = 1.56 g/cm3

37. Significant Figures - Calculations • Addition and Subtraction: The result has the same number of places after the decimal as the least precise measurement in the calculation. For example, calculate the sum of: 10.011g + 5.30g + 9.7093g = 25.0203 = 25.02g

38. Significant Figures - Measurement • All measurements involve some degree of uncertainty. When reading a mass from a digital analytical balance, the last digit (usually one-ten thousandth of a gram) is understood to be uncertain. • When using other devices in the laboratory, such as a ruler, graduated cylinder, buret, etc., you should estimate one place beyond the smallest divisions on the device.

39. The volume should be estimated to the nearest hundredth of a milliliter, since the buret is marked in tenths of a milliliter. The correct reading is 20.15 (or 20.14 or 20.16) mL. It is understood that the last number is uncertain. Significant Figures - Measurements

40. The value of 20.15 mL indicates a volume in between 20.1 mL and 20.2 mL. If the liquid level were resting right on one of the divisions, the reading should reflect this by ending in a zero. Significant Figures - Measurements

41. Conversion of Units Many chemical calculations involve the conversion of units. An example is calculating how many grams of a product can be obtained from a given mass of a reactant. The calculation involves going from mass of reactant to moles of reactant to moles of product to grams of product. You should write in your unites for all calculations, and make sure they cancel properly.

42. Metric Conversion Factors These conversion factors are useful and worth learning. I inch = 2.54 cm 1 lb = 454.6 g 1 L = 1.0567 qt

43. Problem The density of mercury is 13.6 g/mL. What is the weight, in lbs, of a quart of mercury?

44. Accuracy & Precision Most experiments are performed several times to help ensure that the results are meaningful. A single experiment might provide an erroneous result if there is an equipment failure or if a sample is contaminated. By performing several trials, the results may be more reliable.

45. Accuracy & Precision If the experimental values are close to the actual value (if it is known), the data is said to be accurate. If the experimental values are all very similar and reproducible, the data is said to be precise. The goal in making scientific measurements is to that the data be both accurate and precise.

46. Accuracy & Precision Data can be precise, but inaccurate. If a faulty piece of equipment or a contaminated sample is used for all trials, the data may be in agreement (precise), but inaccurate. Such an error is called a systematic error. If the scientist has good technique, the results will be similar, but too high or too low due to the systematic error.

47. Random Error In many experiments, data varies a bit with each trial. The variation in the results is due to random error. Examples might be estimating the last digit for the volume in a buret. Random errors have an equal probability of being too high or too low. As a result, if enough trials are performed, the random error will average itself out.

48. Random and Systematic Errors