Chemistry EIGHTH EDITION by Steven S. Zumdahl Susan Zumdahl

1. ChemistryEIGHTH EDITIONby Steven S. Zumdahl Susan Zumdahl

2. Chapter 1Chemical FoundationsAn Overview • Chemistry: Fundamentally concerned with how one substance changes to another. • Science: • A process for understanding nature and its changes • Primary objective is to explain phenomena of the physical world

3. An Overview • Scientific method – the process that lies at the center of scientific inquiry. Steps in the Scientific Method are: 1. Making observations (collecting data) 2. Making a prediction (formulating a hypothesis) 3. Performing experiments to test the prediction (testing the hypothesis)

4. An Overview • Observations: • Qualitative- does not involve a number • Quantitative- involves both a number and a unit • Hypothesis: A possible explanation for an observation. An educated guess

5. An Overview • Theory (model): A set of tested hypotheses that gives an overall explanation of some natural phenomenon. • Is an interpretation • An attempt to explain why it happens • Law is a summary of observed behavior.The same observation applies to many different systems (Law of Conservation of Mass) • A law summarizes what happens

6. The Various Parts of the Scientific Method

7. Which of the following is the correct order of steps to establish a theory? • Conducting experimental work, collecting observations, making a hypothesis, establishing a theory. • Establishing a theory, conducting experimental work, collecting observations, making a hypothesis. • Making a hypothesis, collecting observations, establishing a theory, conducting experimental work • Collecting observations, making hypothesis, conducting experimental work, establishing a theory. Ans. ‘D’

8. Units of Measurement • Measurement: Compare an unknown quantity with one that is known • A Quantitative measurement consists of two parts • A number • A unit Both parts must be present for the measurement to be meaningful. Example: 20 grams

9. Units of Measurement • Measurements – mass, length, time, temperature etc. • Different systems were adopted in different parts of the world. • English system • Metric system

10. Units of Measurement In 1960, an international agreement set up a system of units called the International System or the SI system. This system is based on the metric system and units derived from the metric system. Mass/kilogram(kg), length/meter(m) Volume/not a fundamental SI unit but is derived from length. Cubic meter (m3)-more commonly used is liter/(L).

11. Prefixes p = pico (0.000000000001) 10-12 n = nano (one billionth) 10-9 µ = micro (0.000001) 10-6 m = milli (0.001) 10-3 k = kilo (1000) 103 M = mega (1000000) 106 G = giga (1000000000) 109

12. Uncertainty in Measurements • There is a limit to the accuracy of any measurement. The digit that must be estimated is called uncertain • A measurement always has some degree of uncertainty • We usually report a measurement by recording all the certain digits plus the first uncertain digit (Significant figures) • The accuracy of any result can not be any better than the data it came from • The overall accuracy is determined by the least accurate number used in the computation

13. Measurement of Volume Using a Buret.The volume is read at the bottom of the liquid curve (called the meniscus).

14. Precision Precision: Refers to the degree of agreement Among several measurements of the same quantity. Reflects the reproducibility of a given type of measurement. 1st series of measurements: 34, 35, 37, 37, 38 2nd series of measurements: 30, 35, 40, 42, 47 The precision of the 1st series is better than the 2nd

15. Accuracy Accuracy: Refers to the agreement of a particular value with the true value Average = Sum of all the values / number of values example: 34, 35, 37, 37, 38 sum of all the values = 181 number of values = 5 average = 36.2 true value = 36.0  good accuracy

16. The Difference between Precision and Accuracy

17. Errors • Random Error (indeterminate error): A measurement has an equal probability of being high or low. This type of error occurs in estimating the value of the last digit of measurement. • Systematic Error (Determinate error): This type of error occurs in the same direction each time. It is either always high or always low, often resulting from poor technique.

18. Rules for Counting Significant Figures • All non-zero digits are significant figures: 1234  4 significant figures • Zeros between non-zero digits (captive zeros) are significant figures: 205  3 significant figures • Zeros beyond decimal point at the end of the number (trailing) are significant figures: 0.24000  5 significant figures • Zeros preceding the first non-zero digit in a number are not significant figures: 0.00453  3 significant figures • Zeros at the end of whole numbers are not significant figures unless you are given information to the contrary: 3400  2 significant figures, 34.00 X 102  4 significant figures, 3400.  4 significant figures. • Exact numbers can be assumed to have an infinite number of significant figures: 2 in 2r (the circumference of a circle), 4 & 3 in 4/3r3 (the volume of a sphere), 1 in = 2.54 cm.

19. How many significant figures are in each of the following? • 12  2 significant figures (S.F.) • 1098  4 S.F. • 2001  4 S.F. • 2.001 x 103  4 S.F. • 0.0000101  3 S.F. • 1.01 x 10-5  3 S.F. • 1000.  4 S.F. (because of the decimal point) • 22.04030  7 S.F.

20. Rules for Significant Figures in Mathematical Operations Multiplication and division: The number of significant figures in the result is the same as that in the quantity with the smallest number of significant figures. Ex.: 4.56 x 1.4 = 6.38  6.4 (3 S.F.) (2 S.F.) (3 S.F.) (2 S.F.) The product should have only two significant figures since 1.4 has two significant figures.

21. Addition and subtraction: The result has the same number of decimal places as the least precise measurement used in the calculation. Ex.: 12.11 + 18.0 + 1.013  31.128  31.1 The correct result is 31.1, since 18.0 has only one decimal place.

22. Rules for Rounding In a series of calculations, carry the extra digits through to the final result, then round. If the digit to be removed • Is less than 5, the preceding digit stays the same. Ex. 1.33 rounds to 1.3. • Is equal to or greater than 5, the preceding digit is increased by 1. Ex. 1.36 rounds to 1.4.

23. Example Perform the following mathematical operation, and Express each result to the correct number of significant figures. • 97.381 + 4.2502 + 0.99195  102.62315  102.623 • 171.5 + 72.915 - 8.23  236.185  236.2 • (9.2 x 100.65) / (8.321 + 4.026)  925.98 / 12.347  74.996  75 • 8.27 (4.987 - 4.962)  8.27(0.025)  0.20675 0.21

24. Dimensional Analysis • Dimensional analysis converts a given result from one system of units to another • The method involves using conversion factors to cancel units until you have the proper unit in the proper place.

25. How do you convert 1.53 minutes to seconds? a. Find a conversion factor (or factors) : 60 sec = 1 min b. Set up start-up and ending information with units 1.53 min. = sec c. Decide which unit to put on top. We need an answer in ‘sec’ and we need to get rid of ‘min’. Therefore, 1.53 min X 60 sec/1 min = 91.8 sec.

26. To convert from one unit to another, use the equivalence statement that relates the two units • Derive the appropriate unit factor by looking at the direction of the required change (to cancel the unwanted units). • Multiply the quantity to be converted by the unit factor to give the quantity with the desired units.

27. Convert 10.0 km to miles Equivalence statements: 1 km = 1000 m, 1 m = 1.09 yd, 1760 yd = 1 mi • Convert 50 miles/hours to feet/minute 103 ft/min

28. Temperature Measurement Three systems for measuring temperature • The Celsius scale (oC) • The Kelvin scale (K) • The Fahrenheit scale (oF)

29. The size of the temperature unit (the degree) is the same for Celsius and Kelvin scales, the difference is in their zero points. Temperature (Kelvin) = temperature (Celsius) + 273.15 TK = TC + 273.15 Temperature (Celsius) = temperature (Kelvin) – 273.15 TC = TK – 273.15

30. The Three Major Temperature Scales

31. Both the degree sizes and zero points are different for Fahrenheit and Celsius scales – need two adjustments • One for degree size • One for the zero point 180 oF = 100 oC Unit factor is 180 oF /100 oC or 9 oF / 5 oC or reciprocal Zero points: 32 oF = 0 oC (subtract 32 from the Fahrenheit temperature) (TF – 32 oF ) 5 oC / 9 oF = TC TF = TC X 9 oF / 5 oC + 32 oF

32. A person has a temperature of 102.5 oF. What is this temperature on the Celsius scale? On the Kelvin scale? TC = (TF -32 oF) 5 oC / 9 oF = (102.5 oF - 32 oF) 5 oC / 9 oF = 39.2 oC TK = TC + 273.2 = (39.2 + 273.2 ) K = 312.4 K

33. Density Density: The mass of substance per unit volume of the substance Density = mass(g)/volume(cm3)  g/cm3 Density is often used as an “identification tag” for a substance.

34. Matter: Anything occupying space and having mass. Matter exits in three states: 1. Solid is rigid, it has a fixed volume and shape. 2. Liquid has a definite volume but no specific shape, it assumes the shape of its container. 3. Gas has no fixed volume or shape, it takes on the shape and volume of its container.

35. Types of Mixtures • Homogeneous mixture: Having visibly indistinguishable parts. Physical properties are the same throughout the material. A homogeneous mixture a solution (example: vinegar). • Heterogeneous mixtures: Having visibly distinguishable parts. Physical properties are different at different points in a material (example: bottle of ranch dressing).

36. Definitions • Pure substance is one with constant composition. Pure substances can be isolated by separation techniques – distillation, filtration, chromatography. • Compound is a substance with constant composition that can be broken down into elements by chemical processes. Example: electrolysis of water produces hydrogen and oxygen. • Physical change is a change in the form of a substance but not in its chemical composition. • Chemical change: a given substance becomes a new substance or substances with different properties and different composition

37. Definitions • Elements are substances that cannot be decomposed into simpler substances by chemical or physical means. • Intensive properties: the value is independent on the amount of material present (color, temperature, melting point, freezing point, density). • Extensive properties: the value is dependent on the amount of material present (weight, length, volume).

38. The Organization of Matter

39. Summary • Scientific method: Consists of looking at the world, developing hypothesis about how it operated, and then making prediction based on these hypotheses. • Measurement: Compare an unknown quantity with a known quantity. • SI unit: International System unit (kg, m, sec). • Uncertainty: Limit to the accuracy of measurement. • Significant figures: Count the number of significant Figures and calculate after mathematical operation. • Dimensional analysis: Converts from one system of units to another. • Temperature conversion: TC = TK – 273.15 TC = (TF – 32 oF) 5 oC/9 oF • Density = m (g)/V (cm3) • Definitions: matter, homogeneous, heterogeneous, physical change, chemical change elements, intensive and extensive properties.