CHAPTER ONE

1. CHAPTER ONE CHEMISTRY: MATTER AND MEASUREMENT

2. Vocabulary • Chemistry • Matter • Energy • Natural Law-(scientific law)

3. Scientific Method • Observation • Hypothesis • Observation or experiment • Theory • Observation or experiment • Law

4. The Elements

5. Chemical and Physical Properties Chemical Properties– chemical changes Changes in chemical composition Relative reactivity Physical Properties– physical changes changes of state, composition remains constant density, color, solubility Extensive Properties– depend on amount present Intensive Properties– physical or chemical properties that are independent of amount present.

6. Periodic Table The periodic table is a convenient method of grouping elements according to similarity in properties and reactivity. The occurrence of these trends and reactivity by the elements is referred to as the Periodic Law. Nomenclature Group or Family vertical (up and down column) of elements on periodic table – these tend to have similar properties Period horizontal row of elements on periodic table

7. Alkaline earth metals Halogens Alkali metals Periodic Table Noble Gases I A Nonmetals VIII A II A III A IV A V A VI A VII A Transition metals Post transition metals Lanthanides Actinides Rare earths

8. Metalloid ‘line’ More metallic More metallic Metals, Nonmetals, and Metalloids Trends in metallic character

9. Representative Elements The elements in A groups on periodic chart are known as representative elements. This is because they are easiest to predict the properties of and best represent what we know about elemental structure and periodicity.

10. d - Transition Elements The elements in B groups on periodic chart are known as transition metals or d-transition elements. All are metals with a valence shell of ns1orns2. There is little variation of properties since inner d-electrons are being added to the elements. They are the ‘transition’ from pure metals to the nonmetals.

11. f - Transition Elements These are known as the inner transition metals or f-transition metals. All have an outer valence of ns2. There is very little variation in the properties of these elements since their electrons are being added 2 shells below the valence shell! Moseley’s triumph

12. Units of Measurement Mass isthe measure of the quantity of matter in a body. Mass Which weighs more, a pound of feathers or a pound of gold? the pound of feathers Which weighs more, an ounce of feathers or an ounce of gold? the ounce of gold The answers are different. Pounds and ounces can be either Troy or common.

13. Measurements in Chemistry Measurement Unit Symbol length meter m mass kilogram kg time second s current ampere A temperature Kelvin K amt. substance mole mol luminous strength candela cd

14. penta P1015 tera T1012 giga G109 mega M106 kilo k 103 deka da 10 deci d 10-1 centi c 10-2 milli m 10-3 micro  10-6 nano n 10-9 pico p 10-12 femto f 10-15 Measurements in Chemistry

15. Units of Measurement Weight measure of the gravitational attraction for a body Length 1 m = 39.37 inches 2.54 cm = 1 inch Volume 1 liter = 1.06 qt 1 qt = 0.946 liter

16. Heat & Temperature Heat and temperature, T, are not the same thing! T is a measure of the intensity of heat in a body There are 3 common temperature scales – all use water as a reference.

17. Heat & Temperature Comparison of the Temperature Scales (using water as the standard) MPBPDifference Fahrenheit 32 °F 212 °F 180 Celsius 0 °C 100 °C 100 Kelvin 273 K 373 K 100 notice that the degree sign is not used with K Celsius is also known as the Centigrade Scale The Fahrenheit equivalent of Kelvin is the Rankine scale (encountered by engineers).

18. Temperature Conversions The relationship between °C and °F is more complicated.

19. Heat and Temperature Example 1: Convert 211 °F to degrees Celsius. A quick mental check would be that 211 °F is just below the boiling point of water on the Fahrenheit scale. Therefore, the answer in °C should be just below 100 °C which would be boiling point of water on the Celsius scale.

20. Heat and Temperature Example 2: Express 548 K in Celsius degrees.

21. Derived Units Quantity Definition Derived Unit Arealength x length m2 Volume area x lengthm3 Density mass/unit volume kg/m3 Acceleration speed/unit time m/s Force mass x acceleration (kgm)/s2 Pressure force/unit area kg/(ms2) Energy force x distance (kgm2)/s2

22. Derived Units - Density Example 3: Calculate the density of a substance if 742 grams of it occupies 97.3 cubic centimeters.

23. Derived Units - Density Example 4: You need 125 g of a corrosive liquid (density 1.32 g/mL) for a reaction. What volume do you need? It is very important to be able to rearrange formulas!

24. Derived Units - Specific Gravity The density of water is nearly 1.00 g/mL at room T. As such, density is very nearly equal to specific gravity. Specific gravity is often used in the place of density because it has no units.

25. Derived Units - Specific Gravity Example 5: A 31.10 gram piece of chromium is dipped into a graduated cylinder that contains 5.00 mL of water. The water level rises to 9.32 mL. What is the specific gravity of chromium? It is common to encounter multi-step problems in chemistry.

26. The Use of Numbers Exact numbers 1 dozen = 12 things 1 inch = 2.54 cm Accuracy how closely measured values agree with the (actual) correct value Precision how closely individual measurements agree with each other (repeatability required) Significant figures digits believed to be correct by the person making or using the measurement

27. Significant Figures - Rules leading zeroes—never significant 0.000357 has three significant figures trailing zeroes—may or may not be significant (common sense may be needed) 1300 nails — counted or weighed? 2400 Cokes — in cases (maybe?) A decimal removes doubt 1300. nails — exact value 2400. soft drinks a counted value Use scientific notation to remove doubt 2.400  103 has four significant figures

28. Significant Figures - Rules imbedded zeroes, zeroes between integers are always significant 3.0604 has five significant figures You must assume that anyone using scientific notation is following the rules. Remember, the exponents “don’t count.” 2.40  105 has 3 significant figures even though the “long hand” version is 24,000

29. Significant Figures - Rules Multiply and Divide Rule — “easy” the product has the smallest number of significant figures found in the multipliers For “small numbers” use the smallest number of decimal places.

30. Significant Figures - Rules Add and Subtract Rule — subtle the answer contains the smallest number of decimal places found in the addends

31. Unit Factor Method This is a simple but important way to always get right answer. This is the easiest way to change from one unit to another. You make unit factors from statements. 1 ft = 12 in becomes 1 L = 1000 mL becomes

32. Unit Factor Method Example 6: Express 9.32 yards in millimeters. There are always multiple ways to do a problem. Alternate pathways include There are still other routes to the solution.

33. End Chapter 1 You need to be aware of the unit conversions in this chapter. You also need to begin memorizing the names of the elements and the names and charges of the common ions. A chart containing this information is found on Web CT under the icon Chapter Notes you need to print up the file entitled Common Ions and Their Valences. A list of the common elements is found within the Chapter 1 notes.