Chapters 1 and 2: Fundamental Concepts and Units of Measurement

1. Chapters 1 and 2: Fundamental Concepts and Units of Measurement Brady & Senese 5th Ed

2. Natural phenomena and measured events; universally consistent ones can be stated as a natural law. Observations : Hypothesis: Tentative proposal that explains observations. revised if experiments do not support it Procedure to test hypothesis; measures one variable at a time. Experiment: Set of conceptual assumptions that explains data from accumulated experiments; predicts related phenomena. Model (Theory): altered if predictions do not support it Further Experiment: Tests predictions based on model. Scientific Method

3. The Scientific Method • Observe results. • Propose a simple and logical hypothesis (tentative explanation for observations). • Hypothesis should correctly predict experimental results. • Hypothesis can be tested by experiment and revised if necessary. • Further testing supporting the hypothesis upgrades it to a theory or model. • When a theory is tested and confirmed by many experiments, it eventually becomes a law.

4. Changes in Matter • Chemical change- a process that results in the formation of a new substance • Evidence? Formation of a new solid, new liquid, new gas, temperature change, or an unexpected color change • Physical change- a process that results in no new substance, but that may change the state of those present, or the proportions 1.3. Matter is Composed of Elements, Compounds, and Mixtures

5. Learning Check: Chemical Or Physical Change? 1.3. Matter is Composed of Elements, Compounds, and Mixtures

6. Matter Can Be Classified By Its Properties: • Matter is either a pure substance or a mixture • Mixtures may be separated using physical methods 1.3. Matter is Composed of Elements, Compounds, and Mixtures

7. What Is An Element? • Elements - substances that cannot be decomposed into simpler substances • shown on the periodic table as symbols: “K” for potassium and “Na” for sodium • made of identical atoms, either singly or in groups 1.3. Matter is Composed of Elements, Compounds, and Mixtures

8. Weird Science • Eleven symbols bear no resemblance to their English names - their names are derived from other languages • Some of these are used in naming, and these are highlighted 1.3. Matter is Composed of Elements, Compounds, and Mixtures

9. What Is A Compound? • Compounds - formed from two or more atoms of different elements combined in a fixed proportion • Have different characteristics than the elements that compose them • Can be broken down into elements by some chemical changes 1.3. Matter is Composed of Elements, Compounds, and Mixtures

10. Mixtures • mixtures consist of varying amounts of two or more elements or compounds • Homogeneous mixtures or “solutions”- have the same properties throughout the sample • Brass, tap water • Heterogeneous mixtures-consist of two or more phases • Salad dressing, Coca-Cola ™ 1.3. Matter is Composed of Elements, Compounds, and Mixtures

11. Classification Of Matter By State Classification based on shape and volume • Solids have fixed shape and volume • Liquids have fixed volume, but take the container shape • Gases have to expand to fill the shape and volume of the container 1.4. Properties of matter can be classified in different ways

12. Properties Of Matter • Chemical properties describe the behavior of the matter that leads to the formation of a new substance: the "reactivity" of the substance • Physical properties can be observed about the matter alone, without changing the composition 1.4. Properties of matter can be classified in different ways

13. Learning Check: Chemical or Physical Property? 1.4. Properties of matter can be classified in different ways

14. Intensive And Extensive Properties • Intensive properties are independent of sample size • Examples: color, texture and temperature • Extensive properties depend on sample size • Examples: volume and mass • Properties used to identify substances are always intensive • Density, color, and texture are often helpful in identification, but temperature is not 1.4. Properties of matter can be classified in different ways

15. Measurements are Observations • Qualitative observations are non-numerical-- ask “what” or “how” or “why” • Quantitative observations are numerical--ask “how much” and are also called measurements • This course is general chemistry with quantitative analysis 1.5 Measurements are essential to describe properties

16. Measurements: • Always involve a comparison • Require units • Involve numbers that are inexact (estimated). This uncertainty is due to the limitations of the observer and the instruments used • In science, all digits in a measurement up to and including the first estimated digit are recorded 1.5 Measurements are essential to describe properties

17. Measurements and units • In the U.S., we use the Imperial(USCS)System • The scientific community (and most of the world) uses the metric system • Variations in the metric system exist, thus a standard system is used: International System of Units (SI) • SI units we will use now: • Length (m) Mass (kg) Time (s) Temperature (K) 1.5 Measurements are essential to describe properties

18. Derived units involve a combination of base units, including: Measurement Formula SI Units Area length × width m2 Volume length × width × height m3 Velocity distance/time m/s Acceleration velocity/time m/s2 Density mass/volume kg/m3 1.5 Measurements are essential to describe properties

19. Decimal multipliers Prefix (Symbol)= Numerical Equivalent • Giga ( G ) = 109 • Mega- ( M ) = 106 • kilo- ( k ) = 103 • centi- ( c ) = 10-2 • milli- ( m ) = 10-3 • micro- ( μ) = 10-6 • nano- ( n ) = 10-9 • pico ( p ) = 10-12 1.5 Measurements are essential to describe properties

20. Temperature • USCS: °F • Metric: °C • SI: K 1.5 Measurements are essential to describe properties

21. Temperature Conversions 1.5 Measurements are essential to describe properties

22. Measurement Error • Because each measurement involves an estimate, measurements always have error. • Record all measured numbers, including the first estimated digit • These digits are called significant digits or significant figures • Exact numbers have infinite significant digits 1.6. Measurements always contain some uncertainty

23. Significant Digits In A Measurement Are Limited By Instrument Precision • Using the first thermometer, the temperature is 21.3 ºC (3 significant digits) • Using the more precise (second) thermometer, the temperature is 21.32 ºC (4 significant digits) 1.6. Measurements always contain some uncertainty

24. Measured Numbers • Measurements are subject to error • For digital instruments record all digits displayed • For analog instruments use the markings to record as many digits as possible and estimate one digit beyond.

25. Errors Arise From A Number Of Sources Including: • Errors-inherent error due to the equipment or procedure • Changing volume due to thermal expansion or contraction (temperature changes) • Improperly calibrated equipment • procedural design allows variable measurements • Mistakes-blunders that you know that you have made. Do not use these data • Spillage • Incomplete procedures • Reading scales incorrectly • Using the measuring device incorrectly 1.6. Measurements always contain some uncertainty

26. Reducing Error: • Random Error can often be detected by making repeated measurements • The average or meanreduces data variations: it helps find a central value 1.6. Measurements always contain some uncertainty

27. Accuracy vs. Precision • An accurate measurement is close to the true or correct value, a “hole-in-one” • A precise measurement is close to the average of a series of repeated measurements • When calibrated instruments are used properly, the greater the number of significant figures, the greater is the degree of precision for a given measurement 1.6. Measurements always contain some uncertainty

28. Rules For Significant Figures (Sig Figs) • Non-zero digits are significant • Zeros between significant digits are significant • Zeros to the right of non-zero digits in a number that contains a decimal point are significant (Trailing with a decimal point) • Zeros to the left of the first nonzero digit are never counted as significant (Leading) • Zeros at the end of a number without a decimal point are assumed not to be significant (Trailing without a decimal place) 1.6. Measurements always contain some uncertainty

29. Learning Check: How Many Significant Figures Are There In The Following? 2.33 3 500.0 4 1000 1 3 .0500 1.6. Measurements always contain some uncertainty

30. Measurements Limit The Precision Of Calculated Results Rules for combining measurements depend on the type of operation performed: • Multiplication and division • The number of sig. figs in the answer should not be greater than the number of sig. figs in the factor with the fewest sig. figs 1.6. Measurements always contain some uncertainty

31. Your Turn! How many sig. figs. result from the following: 12.33 x 0.00002? • 2 • 3 • 4 • 5 • none of these Only 1! 1.6. Measurements always contain some uncertainty

32. Addition and Subtraction The answer should have the same number of decimal places as the quantity with the fewest number of decimal places (least precise). You may also use columns if decimal places aren’t present. 3.247 ← 3 decimal places 41.36 ← 2 decimal places +125.2 ← 1 decimal place 169.8 ← answer rounded to 1 decimal place 1.6. Measurements always contain some uncertainty

33. Exact Numbers • Numbers that come from definitions are exact and have no uncertainty • Typically describe relationships between numbers within the same system (ie: 1m = 100cm) • They can be assumed to contain an infinite number of significant figures 1.6. Measurements always contain some uncertainty

34. Your Turn! How many sig. figs. result from the following? • 2 • 3 • 4 • 5 • none of these 1.6. Measurements always contain some uncertainty

35. Equalities on Food Labels The contents of packaged foods • In the U.S. are listed as both metric and U.S. units. • Indicate the same amount of a substance in two different units.

36. Conversion Factors A conversion factor Is a fraction obtained from an equality. Equality: 1 in. = 2.54 cm Is written as a ratio with a numerator and denominator. Can be inverted to give two conversion factors for every equality. 1 in. and 2.54 cm 2.54 cm 1 in.

37. Conversion Factors in a Problem A conversion factor • May be obtained from information in a word problem. Example 1:The price of one pound (1 lb) of red peppers is$2.39. 1 lb red peppers and $2.39 $2.39 1 lb red peppers Example 2: The cost of one gallon (1 gal) of gas is $2.94. 1 gallon of gas and $2.94 $2.94 1 gallon of gas

38. Percent as a Conversion Factor A percent factor Gives the ratio of the parts to the whole. % = Parts x 100 Whole Use the same units for the parts and whole. Uses the value 100 and a unit for the whole. Can be written as two factors. Example: A food contains 30% (by mass) fat. 30 g fat and 100 g food 100 g food 30 g fat

39. Problem Setup Write the initial and final units. Write a unit plan to convert the initial unit to the final unit. Write equalities and conversion factors. Use conversion factors to cancel the initial unit and provide the final unit. Unit 1 x Unit 2 = Unit 2 Unit 1 Initial x Conversion = Final unit factor unit

40. Setting up a Problem How many minutes are 2.5 hours? Initial unit = 2.5 hr Final unit = ? min Plan= hr min Setup problem to cancel hours (hr). InitialConversion Final unit factor unit 2.5 hr x 60 min = 150 min (2 SF) 1 hr

41. Using Two or More Factors Often, two or more conversion factors are required to obtain the unit needed for the answer. Unit 1 Unit 2 Unit 3 Additional conversion factors are placed in the setup to cancel each preceding unit Initial unit x factor 1 x factor 2 = Final unit Unit 1 x Unit 2 x Unit 3 = Unit 3 Unit 1 Unit 2

42. Example: Problem Solving How many minutes are in 1.4 days? Initial unit: 1.4 days Factor 1 Factor 2 Plan: days hr min Set up problem: 1.4 days x 24 hr x 60 min = 2.0 x 103 min 1 day 1 hr 2 SF Exact Exact =2 SF

43. Check the Unit Cancellation • Be sure to check your unit cancellation in the setup. • The units in the conversion factors must cancel to give the correct unit for the answer. What is wrong with the following setup?1.4 day x 1 day x 1 hr 24 hr 60 min Units = day2/min is not the unit needed Units don’t cancel properly.

44. Percent Factor in a Problem • If the thickness of the skin fold at the waist indicates an 11% body fat, how much fat is in a person with a mass of 86 kg? percent factor 86 kg x 11 kg fat 100 kg = 9.5 kg fat

45. USCS And Metric Units Are Related Using “Critical Links” USCS to Metric Metric to USCS Length 1 in. = 2.54 cm 1 m = 39.37 in 1 yd = 0.9144 m 1 km = 0.6215 mi 1 mi = 1.609 km Mass 1 lb = 453.6 g1 kg = 2.205 lb 1 oz = 28.35 g Volume 1 gal = 3.785 L 1 L = 1.0567 qt 1 qt = 946.4 mL 1 oz (fluid) = 29.6 mL It is also useful to know that 1 mL = 1 cm3=1 cc 1.7 Units can be converted using the factor-label method

46. intensive property defined as the ratio of an object’s mass (m) to volume (v), d = m/v characteristic of pure substances at a specified temperature Since most substances expand when heated, densities decrease when heated. units : g/L for gases and g/mL for solids and liquids. Density(d) 1.8. Density is a useful intensive property

47. Densities of Common Substances

48. Volume by Displacement • A solid completely submerged in water displaces its own volume of water. • The volume of the solid is calculated from the volume difference. • 45.0 mL - 35.5 mL = 9.5 mL = 9.5 cm3

49. Density Using Volume Displacement The density of the zinc object is then calculated from its mass and volume. mass = 68.60 g = 7.2 g/cm3 volume 9.5 cm3

50. Sink or Float • In water, • Ice floats because the density of ice is less than the density of water. • Aluminum sinks because its density is greater than the density of water.