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. The Periodic Table The modern periodic table was developed in 1872 by Dmitri Mendeleev (1834-1907). A similar table was also developed independently by Julius Meyer (1830-1895). The table groups elements with similar properties (both physical and chemical) in vertical columns. As a result, certain properties recur periodically.

19. The Periodic Table Mendeleev left empty spaces in his table for elements that hadn’t yet been discovered. Based on the principle of recurring properties, he was able to predict the density, atomic mass, melting or boiling points and formulas of compounds for several “missing” elements.

20. The Periodic Table

21. The Periodic Table metal/non-metalline

22. The Periodic Table Keep in mind: • Elements along the metal/non-metal dividing line are called semi-metals or metalloids. These elements sometimes behave like metals, and sometimes exhibit non-metallic properties and behavior. • Hydrogen, though in group IA, is not a metal. It is sometimes also placed in group 7A.

23. Properties of Metals Metals are often shiny solids at room temperature. They are malleable and can be pounded into thin sheets. They are also ductile, and can be drawn out into thin wires. Metals are good conductors of electricity.

24. Metallic Character The group IA metals react with water to produce hydrogen and the metal hydroxide. Metallic behavior increases going down a group.

25. Group IIA Metals The group II metals are less reactive, but also react with either warm water or acid to produce hydrogen gas and basic solutions. Due to their high reactivity with water and oxygen, the group IA and IIA metals are never found pure in nature.

26. Non-Metals Non-metals are usually gases or crumbly solids at room temperature. They are non conductors of electricity.

27. Group VIIA The halogens are highly reactive non-metals. They are found pure in nature in salts, but never as the pure element.

28. Group 8A – The Noble Gases

29. Group 8A – The Noble Gases Because they are so non-reactive, the noble or inert gases weren’t discovered until 1894. Although some compounds of these elements have been produced in the laboratory, these elements are known for producing characteristic colors in electric discharge lamps.

30. The Discovery of the Noble Gases It had been known since at least 1785 that air contained something in addition to oxygen and nitrogen. One of the new elements was called argon, from the Greek word for “lazy.”

31. 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.

32. 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

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

34. 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).

35. 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

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

37. 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.

38. 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.

39. 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.

40. 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.

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

42. 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.

43. 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.

44. 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.

45. 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.

46. 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.

47. 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.

48. 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.

49. 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

50. 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