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Mec chapter 1


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MEC Chapter 1

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Mec chapter 1

  1. 1. Copyright© The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 1 Chemistry: Methods and MeasurementDenistonToppingCaret7th Edition
  2. 2. 1.1 The Discovery Process• Chemistry - The study of matter… – Matter - Anything that has mass and occupies space • A table • A piece paper – What about air? • Yes, it is matter
  3. 3. 1.1 The Discovery Process Chemistry: • the study of matter • its chemical and physical properties • the chemical and physical changes it undergoes • the energy changes that accompany those processes • Energy - the ability to do work to accomplish some change
  4. 4. 1.1 The Discovery Process THE SCIENTIFIC METHOD • The scientific method - a systematic approach to the discovery of new information Characteristics of the scientific process 1. Observation 2. Formulation of a question 3. Pattern recognition 4. Developing theories 5. Experimentation 6. Summarizing information
  5. 5. 1.1 The Discovery Process
  6. 6. 1.1 The Discovery Process Models in Chemistry • To aid in understanding of a chemical unit or system – a model is often used – good models are based on everyday experience • Ball and stick methane model – color code balls – sticks show attractive forces holding atoms together
  7. 7. 1.2 Matter and Properties• Properties - characteristics of matter – chemical vs. physical• Three states of matter 1. gas - particles widely separated, no definite shape or volume solid 2. liquid - particles closer together, definite volume but no definite shape 3. solid - particles are very close together, define shape and definite volume
  8. 8. Three States of Water(a) Solid (b) Liquid (c) Gas
  9. 9. 1.2 Matter and Properties Comparison of the Three Physical States
  10. 10. 1.2 Matter and Properties • Physical property - is observed without changing the composition or identity of a substance • Physical change - produces a recognizable difference in the appearance of a substance without causing any change in its composition or identity - conversion from one physical state to another - melting an ice cube
  11. 11. Separation by Physical PropertiesMagnetic iron is separated from other nonmagnetic substances, such as sand. This property is used as a large-scale process in the recycling industry.
  12. 12. 1.2 Matter and Properties • Chemical property - result in a change in composition and can be observed only through a chemical reaction • Chemical reaction (chemical change) - a process of rearranging, removing, replacing, or adding atoms to produce new substances hydrogen + oxygen  water reactants products
  13. 13. 1.2 Matter and Properties Classify the following as either a c h e m i c a l o r p h
  14. 14. 1.2 Matter and Properties Classify the following as either a chemical or physical change: a. Boiling water becomes steam b. Butter turns rancid c. Burning of wood d. Mountain snow pack melting in spring e. Decay of leaves in winter
  15. 15. 1.2 Matter and Properties • Intensive properties - a property of matter that is independent of the quantity of the substance - Density - Specific gravity • Extensive properties - a property of matter that depends on the quantity of the substance - Mass - Volume
  16. 16. 1.2 Matter and Properties Classification of Matter • Pure substance - a substance that has only one component • Mixture - a combination of two or more pure substances in which each substance retains its own identity, not undergoing a chemical reaction
  17. 17. 1.2 Matter and Properties Classification of Matter • Element - a pure substance that cannot be changed into a simpler form of matter by any chemical reaction • Compound - a substance resulting from the combination of two or more elements in a definite, reproducible way, in a fixed ratio
  18. 18. 1.2 Matter and Properties Classification of Matter • Mixture - a combination of two or more pure substances in which each substance retains its own identity • Homogeneous - uniform composition, particles well mixed, thoroughly intermingled • Heterogeneous – nonuniform composition, random placement
  19. 19. 1.2 Matter and Properties Classes of Matter
  20. 20. 1.3 Measurement in ChemistryData, Results, and Units• Data - each piece is an individual result of a single measurement or observation – mass of a sample – temperature of a solution• Results - the outcome of the experiment• Data and results may be identical, however usually related data are combined to generate a result• Units - the basic quantity of mass, volume or whatever quantity is being measured – A measurement is useless without its units
  21. 21. English and Metric Units • English system - a collection of1.3 Measurement in functionally unrelated units – Difficult to convert from one unit to Chemistry another – 1 foot = 12 inches = 0.33 yard = 1/5280 miles • Metric System - composed of a set of units that are related to each other decimally, systematic – Units relate by powers of tens – 1 meter = 10 decimeters = 100 centimeters = 1000 millimeters
  22. 22. Basic Units of the Metric System1.3 Measurement in Mass gram g Length meter m Chemistry Volume liter L • Basic units are the units of a quantity without any metric prefix
  23. 23. 1.3 Measurement in Chemistry
  24. 24. UNIT CONVERSION1.3 Measurement in • You must be able to convert between Chemistry units - within the metric system - between the English system and metric system • The method used for conversion is called the Factor-Label Method or Dimensional Analysis !!!!!!!!!!! VERY IMPORTANT !!!!!!!!!!!
  25. 25. • Let your units do the work for you by1.3 Measurement in simply memorizing connections between units. Chemistry – For example: How many donuts are in one dozen? – We say: “Twelve donuts are in a dozen.” – Or: 12 donuts = 1 dozen donuts • What does any number divided by itself equal? 12 donuts • ONE! =1 1 dozen
  26. 26. 12 donuts =11.3 Measurement in 1 dozen Chemistry • This fraction is called a unit factor • What does any number times one equal? • That number • Multiplication by a unit factor does not change the amount – only the unit
  27. 27. • We use these two mathematical facts to1.3 Measurement in use the factor label method – a number divided by itself = 1 Chemistry – any number times one gives that number back • Example: How many donuts are in 3.5 dozen? • You can probably do this in your head but try it using the Factor-Label Method.
  28. 28. 1.3 Measurement in Start with the given information... 12 donuts 3.5 dozen × = 42 donuts Chemistry 1 dozen Then set up your unit factor... See that the units cancel... Then multiply and divide all numbers...
  29. 29. Common English System Units1.3 Measurement in Chemistry • Convert 12 gallons to units of quarts
  30. 30. Intersystem Conversion Units1.3 Measurement in Chemistry • Convert 4.00 ounces to kilograms
  31. 31. 1.3 Measurement in Chemistry 1. Convert 5.5 inches to millimeters 2. Convert 50.0 milliliters to pints 3. Convert 1.8 in2 to cm2
  32. 32. 1.4 Significant Figures and Scientific Notation• Information-bearing digits or figures in a number are significant figures• The measuring devise used determines the number of significant figures a measurement has• The amount of uncertainty associated with a measurement is indicated by the number of digits or figures used to represent the information
  33. 33. and Scientific Notation1.4 Significant Figures Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit
  34. 34. and Scientific Notation Recognition of Significant Figures1.4 Significant Figures • All nonzero digits are significant • 7.314 has four significant digits • The number of significant digits is independent of the position of the decimal point • 73.14 also has four significant digits • Zeros located between nonzero digits are significant • 60.052 has five significant digits
  35. 35. Use of Zeros in Significantand Scientific Notation1.4 Significant Figures Figures • Zeros at the end of a number (trailing zeros) are significant if the number contains a decimal point. • 4.70 has three significant digits • Trailing zeros are insignificant if the number does not contain a decimal point. • 100 has one significant digit; 100. has three • Zeros to the left of the first nonzero integer are not significant. • 0.0032 has two significant digits
  36. 36. and Scientific Notation1.4 Significant Figures How many significant figures are in the following? 1. 3.400 2. 3004 3. 300. 4. 0.003040
  37. 37. and Scientific Notation Scientific Notation1.4 Significant Figures • Used to express very large or very small numbers easily and with the correct number of significant figures • Represents a number as a power of ten • Example: 4,300 = 4.3 x 1,000 = 4.3 x 103
  38. 38. and Scientific Notation • To convert a number greater than 1 to1.4 Significant Figures scientific notation, the original decimal point is moved x places to the left, and the resulting number is multiplied by 10x • The exponent x is a positive number equal to the number of places the decimal point moved 5340 = 5.34 x 104 • What if you want to show the above number has four significant figures? = 5.340 x 104
  39. 39. and Scientific Notation1.4 Significant Figures • To convert a number less than 1 to scientific notation, the original decimal point is moved x places to the right, and the resulting number is multiplied by 10-x • The exponent x is a negative number equal to the number of places the decimal point moved 0.0534 = 5.34 x 10-2
  40. 40. and Scientific Notation Types of Uncertainty1.4 Significant Figures • Error - the difference between the true value and our estimation – Random – Systematic • Accuracy - the degree of agreement between the true value and the measured value • Precision - a measure of the agreement of replicate measurements
  41. 41. Significant Figures in Calculation ofand Scientific Notation Results1.4 Significant Figures Rules for Addition and Subtraction • The result in a calculation cannot have greater significance than any of the quantities that produced the result • Consider: 37.68 liters 6.71862 liters 108.428 liters 152.82662 liters correct answer 152.83 liters
  42. 42. Rules for Multiplication and Divisionand Scientific Notation1.4 Significant Figures • The answer can be no more precise than the least precise number from which the answer is derived • The least precise number is the one with the fewest significant figures 4.2 × 103 (15.94) = 2.9688692 ×10 −8 (on calculator) 2.255 ×10 − 4 Which number has the fewest significant figures? 4.2 x 103 has only 2 The answer is therefore, 3.0 x 10-8
  43. 43. and Scientific Notation Exact and Inexact Numbers1.4 Significant Figures • Inexact numbers have uncertainty by definition • Exact numbers are a consequence of counting • A set of counted items (beakers on a shelf) has no uncertainty • Exact numbers by definition have an infinite number of significant figures
  44. 44. Rules for Rounding Off Numbersand Scientific Notation1.4 Significant Figures • When the number to be dropped is less than 5 the preceding number is not changed • When the number to be dropped is 5 or larger, the preceding number is increased by one unit • Round the following number to 3 significant figures: 3.34966 x 104 =3.35 x 104
  45. 45. and Scientific Notation1.4 Significant Figures How Many Significant Figures? Round off each number to 3 significant figures: 1. 61.40 2. 6.171 3. 0.066494
  46. 46. 1.5 Experimental Quantities• Mass - the quantity of matter in an object – not synonymous with weight – standard unit is the gram• Weight = mass x acceleration due to gravity• Mass must be measured on a balance (not a scale)
  47. 47. 1.5 Experimental Quantities • Units should be chosen to suit the quantity described – A dump truck is measured in tons – A person is measured in kg or pounds – A paperclip is measured in g or ounces – An atom? • For atoms, we use the atomic mass unit (amu) – 1 amu = 1.661 x 10-24 g
  48. 48. 1.5 Experimental Quantities • Length - the distance between two points – standard unit is the meter – long distances are measured in km – distances between atoms are measured in nm, 1 nm = 10-9 m • Volume - the space occupied by an object – standard unit is the liter – the liter is the volume occupied by 1000 grams of water at 4 oC – 1 mL = 1/1000 L = 1 cm3
  49. 49. 1.5 Experimental Quantities The milliliter (mL) and the cubic centimeter (cm3) are equivalent
  50. 50. 1.5 Experimental Quantities • Time - metric unit is the second • Temperature - the degree of “hotness” of an object
  51. 51. 1.5 Experimental Quantities Conversions Between Fahrenheit and Celsius o F - 32 o C= 1.8 o F = 1.8 ×( C) + 32 o 1. Convert 75oC to oF 2. Convert -10oF to oC 1. Ans. 167 oF
  52. 52. 1.5 Experimental Quantities Kelvin Temperature Scale • The Kelvin scale is another temperature scale. • It is of particular importance because it is directly related to molecular motion. • As molecular speed increases, the Kelvin temperature proportionately increases. K = oC + 273
  53. 53. 1.5 Experimental Quantities • Energy
  54. 54. 1.5 Experimental Quantities Characteristics of Energy • Energy cannot be created or destroyed • Energy may be converted from one form to another • Energy conversion always occurs with less than 100% efficiency • All chemical reactions involve either a “gain” or “loss” of energy
  55. 55. 1.5 Experimental Quantities Units of Energy • Basic Units: • calorie or joule • 1 calorie (cal) = 4.184 joules (J) • A kilocalorie (kcal) also known as the large Calorie. This is the same Calorie as food Calories. • 1 kcal = 1 Calorie = 1000 calories • 1 calorie = the amount of heat energy required to increase the temperature of 1 gram of water 1oC.
  56. 56. 1.5 Experimental Quantities Density and Specific Gravity • Density – the ratio of mass to volume – an extensive property – use to characterize a substance as each substance has a unique density – Units for density include: • g/mL • g/cm3 mass m d= = • g/cc volume V
  57. 57. 1.5 Experimental Quantities cork water brass nut liquid mercury
  58. 58. 1.5 Experimental Quantities Calculating the Density of a Solid • 2.00 cm3 of aluminum are found to weigh 5.40g. Calculate the density of aluminum in units of g/cm3. mass m – Use the formula d= = volume V – Substitute our values 5.40 g 2.00 cm3 = 2.70 g / cm3
  59. 59. 1.5 Experimental Quantities Air has a density of 0.0013 g/mL. What is the mass of 6.0-L sample of air? Calculate the mass in grams of 10.0 mL if mercury (Hg) if the density of Hg is 13.6 g/mL. Calculate the volume in milliliters, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 grams.
  60. 60. 1.5 Experimental Quantities Specific Gravity • Values of density are often related to a standard • Specific gravity - the ratio of the density of the object in question to the density of pure water at 4oC • Specific gravity is a unitless term because the 2 units cancel • Often the health industry uses specific gravity to test urine and blood samples density of object (g/mL) specific gravity = density of water (g/mL)