Understanding Scientific Concepts in Chemistry: Definitions and Measurements

1. Chemistry Fall 2003 Dr Supplee

2. Chapter 1- Definitions • Science • Methodical exploration of nature followed by a logical explanation of observations • Scientific Method • A systematic investigation of nature and requires proposing an explanation for the results of an experiment in the form of a general principle (hypothesis)

3. Chapter 1 - Definitions • Hypothesis • Initial explanation of observations • Theory • Sufficient evidence in support of the hypothesis • Model that scientifically explains the behavior of nature • Law • Does not explain behavior • States a measurable relationship under different experimental conditions

4. Chapter 1 – Definition Examples • Hypothesis • Dalton proposed that all matter was composed of small individual particles (atoms) • Theory • 100 years later Atomic Theory which explains the composition of substances as well as the behavior of gases • Law • Boyle’s Law P1V1= P2V2 at constant temperature • If volume decreases than pressure increases at constant temperature

5. Chapter 1- Definitions Summary Scientific Theory Natural Law Analyze more data Hypothesis Analyze initial observations Experiment

6. Chapter 1 – Modern Chemistry • Organic Chemistry • Chemistry of carbon containing compounds • (C, H, O, and N) • Inorganic Chemistry • Chemistry of all other substances • Biochemistry • Chemistry of substances derived from plant substances

7. Chapter 1 – Modern Chemistry • All three have in common • Analytical Chemistry • Qualitative (what) and quantitative (how much) analyses • Physical Chemistry • Theoretical and mathematical explanations of chemical behavior

8. Relevance to daily life Interesting topics Fun experiments CHEMISTRY Career Opportunities Benefits to society Applications

9. Chapter 2- Scientific Measurements • Introduction to Laboratory • Work alone • Handout • Due 9/15/03 • Measurement Uncertainty • Plus/minus factor ( error) • Metric versus English Units • Conversion factors • Significant Figures • Rounding rules

10. Precision versus Accuracy Precise, not accurate True Value Precision –how close two measurements of the same quantity are to each other Accurate, not precise Accuracy – how close an experimental observations to the true value

11. Chapter 2- Scientific Measurements • Measurement • a number with units • Uncertainty • the estimated unit amount • plus/minus associated with measurement • Mass • Amount of matter an object possesses • Weight • Force exerted by gravity on an object

12. Volume Amount of space occupied by a solid, gas or liquid Chapter 2- Scientific Measurements

13. Significant Digits/ Figures • Digits are significant when the do more than hold a decimal place • A place holder zero is NEVER significant • determines measurement uncertainty (error analysis) • Does not apply for exact numbers, only measured numbers

14. Significant Digits/ Figures Rule • Rule #1 • Count the number of nonzero digits left to right • Do not count place holder zeros

15. Significant Figure Rounding Rules • After all calculations are complete determine significant figures and then round • 5 or greater round-up to the nearest whole number • less than 5 truncate

16. Scientific Notation • Exponential numbers (power of 10) Base 10exponent • The number 10 is raised to the nth power • Numbers greater than 1 the exponent is positive • Numbers less than 1 the exponent is negative • The decimal is placed after the first significant digit and sets the size of the number by using a power of 10.

17. Unit Equations, Factors and Conversions • Problem Solving Technique • Equivalent relationships • Unit equation • A simple statement of two equivalent quantities • Unit Factor • A ratio of two equivalent quantities

18. Unit Dimensional Analysis Problem Solving • Three steps 1) write down the units asked for in the answer 2) write down the value given in the problem that is related to the required answer 3) Apply a unit factor to convert the units in the given value to the units in the answer Given Value x Unit = units asked for Factor

19. Percent Concept • amount of a single quantity compared to the entire sample • one part per 100 parts one quantity x 100 = % total sample

20. Review

21. Significant Digits/ Figures • Digits are significant when the do more than hold a decimal place • If the number is less than 1, a place holder zero is NEVER significant • determines measurement uncertainty (error analysis) • Does not apply for exact numbers, only measured numbers

22. Exact Numbers • Infinite significant figures • English to English conversion factors • Metric to metric conversion factors

23. Unit Equations, Factors and Conversions • Problem Solving Technique • Equivalent relationships • Unit equation • A simple statement of two equivalent quantities • Unit Factor • A ratio of two equivalent quantities

24. Chapter 3 – The Metric System • Single basic unit for each quantity measured • Decimal system that uses a system of prefixes to enlarge or reduce a basic unit

25. Metric System Definitions • Meter equals one ten-millionth of the distance from the North Pole to the equator • Kilogram equals the mass of one a cube of water one-tenth of a meter on a side • Liter equals the volume occupied by a kilogram of water at 4 oC

26. The Metric System

27. Metric Prefixes

28. Metric Conversion Factors Practice • 1 kg =? g k = kilo = 1000 basic units 1kg = 1000g • 2s =? ns n=nano=1 1 x 10-9 2s=2 x 10-9 ns

29. Unit Conversion Factors • Ratio of two equivalent quantities • The quantity in the numerator is equal to the quantity in the denominator • If 100cm = 1 m, then the factor becomes 100 cm or 1m 1 m 100 cm

30. Metric- English Conversions

31. Unit Analysis • Recall: • Problem Solving Technique Units Given Unit Factor New unit Unit Factor Units asked for

32. Practice Problems • Work in groups of 3-4 • One student from each group puts solution in board and explains to class

33. Quiz # 4 • See Chemistry Current News Slides • Presentation to be given on Oct 6, 2003.

34. Density - Review • Lab Experiment 2 • Physical property • Defined as mass per unit volume • Liquids and solids expressed in g/ml (g/ cm3) • Gases expresses in grams per liter • Density of water is 1.00 g/ml • Floats in water density <1.00 g/ml • Sinks in water density >1.00g/ml

35. Estimating Density(page 59 and 60 ) Water, chloroform and ethyl ether are poured into a tall glass cylinder. Three known solids are added. Identify the liquids. Liquid 1 Solid 1 = ice Liquid 2 = water Solid 2 = rubber Liquid 3 Solid 3 = aluminum

36. TemperatureFahrenheit, Celsius and Kelvin • Measure of the average energy of individual particles in a system • Warmer temps = more molecules moving thus more energy • Cooler temps = slow moving molecules thus less energy • Fahrenheit oF • Celsius oC • Kelvin K

37. Temperature • oF • Freezing point of water 32 oF • Boiling point of water 212 oF • oC • Freezing point of water 0 oC • Boiling point of water 100 oC • K ( SI unit) • Absolute zero 0 K • Equal to -273.15oC

38. Temperature Conversions • oF to oC ( oF - 32 oF ) x 100 oC / 180 oF = oC • oC to oF ( oC x 180 oF / 100 oC ) +32 = oF • Kelvin oC +273

39. Heat • Heat measures the total energy • Temperature measures the average energy • Heat energy units calories or kilocalories • A calorie (cal) is defined as the amount of heat needed to raise 1 g of water 1 oC • Food Calorie equals 1 kcal = 1000 cal • SI unit = joule (J) 1 cal = 4.184 J

40. Specific Heat • Amount of heat required to bring about a given change in temperature • Observed amount • Unique for each substance • Specific heat of water is high • Change in temperature is minimal as water gains or losses heat • Surface of earth is covered in water so water helps to regulate the climates

41. Specific Heat • Amount of heat required to raise the temperature of 1 g of substance 1 oC • Units are cal/g oC Water Ice Iron Silver 1 g 1 g 1 g 1 g 1.0 oC 9.3 oC 17.7 oC 2.0 oC

42. Specific Heat • gain or loss of heat divided by mass and temperature change = specific heat How many calories are required to raise the temperature of a 3 inch iron nail weighing 7.05 g form room temperature to 100 oC? The specific heat of iron is 0.108 cal/g oC

43. Solution • Specific Heat = Heat/ (mass x D t) cal/g oC = cal / g x oC • 0.108 cal/g oC = cal / 7.05 x (100-25oC) • Solving for Heat ( energy required ) • Rearrange (0.108 cal/g oC) x 7.05 g x 75 oC = 57 cal

44. Chapter 4 Matter and Energy • Matter is any substance that has mass and occupies volume • Physical State changes • Melting solid into liquid • Sublimation solid into gas • Condensation gas into liquid • Deposition gas to solid • Freezes liquid to solid • Vaporization liquid to gas

45. Increasing temperature steam ice water melting vaporizing freezing condensing Sublimation Deposition

46. Chapter 4 Matter and Energy

47. Elements, Compounds and Mixtures • Properties of matter may be consistent throughout or they may vary • Melting point • Gold (Au) 1064 oC • Quartz 1000 – 1600 oC • Gold is homogenous – properties consistent • Quartz is heterogeneous – properties vary

48. Mixtures • Heterogeneous • Usually Solids • Separated into pure substances by physical methods which take advantage of different physical properties • Properties are not the same throughout the sample • Homogeneous • Gases or liquids • Separated into pure substances by either chemical or physical methods which take advantage of different physical properties • Properties a the same for any given sample, but can vary sample to sample

49. Mixtures • Alloy • Homogeneous mixture of two or more metals • Gold ( Au) 10 K 14 K 18 K 42 % 75% • Substance • Matter with definite composition and constant properties • Compound or an element • Compound • Broken down into elements by chemical reactions • Element • Cannot be broken down further by chemical reactions

50. Matter Mixtures Physical Separate Substances Heterogeneous Homogeneous Compounds Elements