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Chapter 1

Copyright  The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 1. Chemistry: Methods and Measurement. Denniston Topping Caret 5 th Edition. 1.1 The Discovery Process. Chemistry - The study of matter…

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Chapter 1

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  1. Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 1 Chemistry: Methods and Measurement Denniston Topping Caret 5th Edition

  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. 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 1.1 The Discovery Process

  4. MAJOR AREAS OF CHEMISTRY • Biochemistry - the study of life at the molecular level • Organic chemistry - the study of matter containing carbon and hydrogen • Inorganic chemistry - the study of matter containing elements, not organic • Analytic chemistry - analyze matter to determine identity and composition 1.1 The Discovery Process

  5. Physical chemistry - attempts to explain the way matter behaves pharmaceutical industry public health 1.1 The Discovery Process CHEMISTRY food science medical practitioners forensic sciences

  6. THE SCIENTIFIC METHOD • The scientific method - a systematic approach to the discovery of new information • Characteristics of the scientific process • Observation • Formulation of a question • Pattern recognition • Developing theories • Experimentation • Summarizing information 1.1 The Discovery Process

  7. 1.1 The Discovery Process

  8. 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 1.1 The Discovery Process

  9. 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

  10. Three States of Water (a) Solid (b) Liquid (c) Gas

  11. Comparison of the Three Physical States 1.2 Matter and Properties

  12. 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 1.2 Matter and Properties

  13. Separation by Physical Properties Magnetic iron is separated from other nonmagnetic substances, such as sand. This property is used as a large-scale process in the recycling industry.

  14. 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 1.2 Matter and Properties hydrogen + oxygen  water products reactants

  15. Classify the following as either a chemical or physical property: • Color • Flammability • Hardness • Odor • Taste 1.2 Matter and Properties

  16. Classify the following as either a chemical or physical change: • Boiling water becomes steam • Butter turns rancid • Burning of wood • Mountain snow pack melting in spring • Decay of leaves in winter 1.2 Matter and Properties

  17. 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 1.2 Matter and Properties

  18. Classification of Matter 1.2 Matter and Properties • 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

  19. Classification of Matter 1.2 Matter and Properties • 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

  20. Classification of Matter 1.2 Matter and Properties • 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

  21. Classes of Matter 1.2 Matter and Properties

  22. 1.3 Measurement in Chemistry Data, 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

  23. English and Metric Units • English system - a collection of functionally unrelated units • Difficult to convert from one unit to 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 1.3 Measurement in Chemistry

  24. Basic Units of the Metric System Mass gram g Length meter m Volume liter L 1.3 Measurement in Chemistry • Basic units are the units of a quantity without any metric prefix

  25. 1.3 Measurement in Chemistry

  26. UNIT CONVERSION • You must be able to convert between 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 1.3 Measurement in Chemistry !!!!!!!!!!! VERY IMPORTANT !!!!!!!!!!!

  27. Let your units do the work for you by simply memorizing connections between units. • For example: How many donuts are in one dozen? 1.3 Measurement in Chemistry • We say: “Twelve donuts are in a dozen.” • Or: 12 donuts = 1 dozen donuts • What does any number divided by itself equal? • ONE!

  28. This fraction is called a unit factor 1.3 Measurement in Chemistry • What does any number times one equal? • That number • Multiplication by a unit factor does not change the amount – only the unit

  29. We use these two mathematical facts to use the factor label method • a number divided by itself = 1 • 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. 1.3 Measurement in Chemistry

  30. Start with the given information... 3.5 dozen = 42 donuts 1.3 Measurement in Chemistry Then set up your unit factor... See that the units cancel... Then multiply and divide all numbers...

  31. Common English System Units 1.3 Measurement in Chemistry • Convert 12 gallons to units of quarts

  32. Intersystem Conversion Units 1.3 Measurement in Chemistry • Convert 4.00 ounces to kilograms

  33. Practice Unit Conversion 1. Convert 5.5 inches to millimeters 2. Convert 50.0 milliliters to pints 3. Convert 1.8 in2 to cm2 1.3 Measurement in Chemistry

  34. 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

  35. 1.4 Significant Figures and Scientific Notation Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit

  36. Recognition of 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 1.4 Significant Figures and Scientific Notation

  37. Use of Zeros in Significant 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 1.4 Significant Figures and Scientific Notation

  38. Determining Significant Figures • How many significant figures are in the following? • 3.400 • 3004 • 300. • 0.003040 1.4 Significant Figures and Scientific Notation

  39. Scientific Notation • 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 1.4 Significant Figures and Scientific Notation

  40. To convert a number greater than 1 to 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 1.4 Significant Figures and Scientific Notation

  41. 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 1.4 Significant Figures and Scientific Notation

  42. Types of Uncertainty • 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 1.4 Significant Figures and Scientific Notation

  43. Significant Figures in Calculation of Results 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 1.4 Significant Figures and Scientific Notation correct answer 152.83 liters

  44. Rules for Multiplication and Division • 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 1.4 Significant Figures and Scientific Notation Which number has the fewest significant figures? 4.2 x 103 has only 2 The answer is therefore, 3.0 x 10-8

  45. Exact and Inexact Numbers • 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 1.4 Significant Figures and Scientific Notation

  46. Rules for Rounding Off Numbers • 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 1.4 Significant Figures and Scientific Notation • =3.35 x 104

  47. How Many Significant Figures? Round off each number to 3 significant figures: • 61.40 • 6.171 • 0.066494 1.4 Significant Figures and Scientific Notation

  48. 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)

  49. 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

  50. 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 1.5 Experimental Quantities

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