Chapter 1 Chemistry: Methods and Measurement PDF

Because learning changes everything.® GENERAL, ORGANIC, AND 11TH Edition BIOCHEMISTRY...

Because learning changes everything.® GENERAL, ORGANIC, AND 11TH Edition BIOCHEMISTRY Katherine J. Denniston Danaè R. Quirk Joseph J. Topping Robert L. Caret Chapter 1 Chemistry: Methods and Measurement © McGraw Hill LLC. All rights reserved. No reproduction or distribution without the prior written consent of McGraw Hill LLC.  1.3 The Classification of Matter Properties - characteristics of matter scientists can use to categorize different types of matter. Ways to Categorize matter: 1. By State 2. By Composition © McGraw Hill LLC  Three States of Matter 1. Gas - particles widely separated, no definite shape or volume solid. 2. Liquid - particles closer together, definite volume Loading… but no definite shape. 3. Solid - particles are very close together, define shape and definite volume. © McGraw Hill LLC  Composition 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. © McGraw Hill LLC  Pure Substances Loading… Element - a pure substance that cannot be changed into a simpler form of matter by any chemical reaction. Compound - a pure substance resulting from the combination of two or more elements in a definite, reproducible way, in a fixed ratio. © McGraw Hill LLC  Mixture 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. © McGraw Hill LLC  Classes of Matter © McGraw Hill LLC Image Source Plus/Alamy Stock Photo; Image Source/Getty Images; Danaè R. Quirk Dorr, Ph.D.  Physical Property versus Physical Change 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. © McGraw Hill LLC  Physical Properties and Physical Change © McGraw Hill LLC (a) moodboard/Glow Images; (b) S.Borisov/Shutterstock; (c) WeatherVideoHD.TV  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. © McGraw Hill LLC Ken Karp/MacGraw Hill 1  Chemical Property versus Chemical Reaction Chemical property - results in a change in composition and can be observed only through a chemical reaction. Loading… Chemical reaction (chemical change) - a chemical substance is converted in to one or more different substances by rearranging, removing, replacing, or adding atoms. © McGraw Hill LLC 1  Classification of Properties Classify the following as either a chemical or physical property: a. Color b. Flammability c. Hardness d. Odor e. Taste © McGraw Hill LLC 1  Classification of Changes 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 melting in spring. e. Decay of leaves in winter. © McGraw Hill LLC 1  Intensive and Extensive Properties Intensive properties - a property of matter that is independent of the quantity of the substance. Color. Melting Point. Extensive properties - a property of matter that depends on the quantity of the substance. Mass. Volume. © McGraw Hill LLC 1  1.4 The Units of Measurement Units - the basic quantity of mass, volume or whatever quantity is being measured. A measurement is useless without its 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. © McGraw Hill LLC 1  Metric System Units 1 Mass - the quantity of matter in an object not synonymous with weight. Weight = mass × acceleration due to gravity. Standard unit is the gram (g). The pound (lb) is the common English unit. 1 lb = 453.6 g Mass must be measured on a balance (not a scale). © McGraw Hill LLC 1  Metric System Units 2 Length - the distance between two points Standard unit is the meter (m). The yard is the common English unit. 1 yd = 0.9144 m Volume - the space occupied by an object Standard unit is the liter (L). The quart is the common English unit. 1 qt = 0.9464 L Time The metric unit is the second (s). © McGraw Hill LLC 1  Metric System Prefixes Basic units are the units of a quantity without any metric prefix. © McGraw Hill LLC 1  Relationship among various volume units Volume = Length × width × height Volume = 1 dm × 1 dm × 1 dm = © McGraw Hill LLC 1  1.5 The Numbers of Measurement Information - bearing digits or figures in a number are significant figures. The measuring device used determines the number of significant figures in a measurement. The degree of uncertainty associated with a measurement is indicated by the number of figures used to represent the information. © McGraw Hill LLC 2  Significant Figures Example Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit. © McGraw Hill LLC 2  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. © McGraw Hill LLC 2  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. 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. © McGraw Hill LLC 2  How many significant figures are in the following? 1. 3.400 2. 3004 3. 300. 4. 0.003040 © McGraw Hill LLC 2  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: © McGraw Hill LLC 2  Scientific Notation Rules 1 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 The exponent x is a positive number equal to the number of places the decimal point moved. What if you want to express the above number with three significant figures? © McGraw Hill LLC 2  Scientific Notation Rules 2 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. The exponent x is a negative number equal to the number of places the decimal point moved. © McGraw Hill LLC 2  Scientific Notation Example When a number is exceedingly large or small, scientific notation must be used to input the number into a calculator: 0.0000000000000000000000066466 g must be entered into calculator as: © McGraw Hill LLC 2  Represent the following numbers in scientific notation: 1. 0.00018 2. 3004 3. 300. Loading… 4. 0.00304 © McGraw Hill LLC 2  Accuracy and Precision Accuracy - the degree of agreement between the true value and the measured value. Error - the difference between the true value and our estimation. Random. Systematic. Precision - a measure of the agreement of replicate measurements. Deviation – amount of variation present in a set of replicate measurements. © McGraw Hill LLC 3  Exact (Counted) and Inexact Numbers Inexact numbers have uncertainty (degree of doubt in final significant digit) 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. © McGraw Hill LLC 3  Rules for Rounding 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: © McGraw Hill LLC 3  Round off each number to three significant figures: 1. 61.40 2. 6.171 3. 0.066494 © McGraw Hill LLC 3  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 correct answer: 152.83 liters © McGraw Hill LLC 3  Report the result of each to the proper number of significant figures: 1. 4.26 + 3.831 2. 8.321 − 2.4 © McGraw Hill LLC 3  Adding and Subtracting in Scientific Notation There are two ways to solve the following: SOLUTION 1: convert both numbers to standard form and add 0.00000947 + 0.000093 0.00010247 correct answer: © McGraw Hill LLC 3  Addition Example There are two ways to solve the following: SOLUTION 2: change one of the exponents so that both have the same power of 10, then add changes to correct answer: © McGraw Hill LLC 3  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. Which number has the fewest significant figures? has only 2 The answer is therefore, 3.0 © McGraw Hill LLC 3  1.6 Unit Conversion Factor-Label Method (Dimensional Analysis) Uses Conversion Factors to: Convert from one unit to another within the same system. Convert units from one system to another. © McGraw Hill LLC 3  English Unit Conversion - Example To convert from one unit to another you must know the conversion factor, which is the relationship between the two units. The Relationship: 1 gal = 4 qt The Conversion Factor: © McGraw Hill LLC 4  Using Conversion Factors Convert 12 gallons to quarts. The Relationship (English system): 1 gal = 4 qt The Conversion Factor: Data Given: 12 gal. Use Conversion Factor with gal in denominator. © McGraw Hill LLC 4  Using Conversion Factors - Solution Convert 12 gallons to quarts. Solution: Write the Data Given. Multiply by the Conversion Factor with the unit of the Data Given (gal) in the denominator. Desired Result © McGraw Hill LLC 4  Unit Conversion - Example Convert 360 feet to miles. The Relationship (English system): 5280 ft = 1 mi The Conversion Factor: Data Given: 360 ft. Use Conversion Factor with ft in denominator. © McGraw Hill LLC 4  Unit Conversion - Solution Convert 360 feet to miles. Solution: Write the Data Given. Multiply by the Conversion Factor with the unit of the Data Given (ft) in the denominator. Desired Result © McGraw Hill LLC 4  Multistep Conversion - Example Convert 0.0047 kilograms to milligrams The Relationships (metric system): and The Conversion Factors: Data Given: 0.0047 kg 1. Use Conversion Factor with kg in denominator to convert to Initial Data Result in g. 2. Use Conversion Factor with g in denominator. © McGraw Hill LLC 4  Multistep Conversion - Solution Convert 0.0047 kilograms to milligrams. Data Given × Conversion Factor = Initial Data Result Initial Data Result × Conversion Factor = Desired Result © McGraw Hill LLC 4  Multistep Conversions - Alternate Solution Convert 0.0047 kilograms to milligrams. Alternatively, solve in a single step: Data Given × Conversion Factor × Conversion Factor = Desired Result © McGraw Hill LLC 4  Practice Unit Conversions 1. Convert 5.5 inches to millimeters. 2. Convert 50.0 milliliters to pints. 3. Convert. © McGraw Hill LLC 4  1.7 Additional Experimental Quantities Temperature - the degree of “hotness” of an object. © McGraw Hill LLC 4  Conversions Between Fahrenheit and Celsius 1. Convert 75 degrees C to degrees F. 2. Convert −10 degrees F to degrees C. 1. Answer: 167 degrees F 2. © McGraw Hill LLC Answer: −23 degrees C 5  Kelvin Temperature Scale The Kelvin (K) 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. © McGraw Hill LLC 5  Energy Energy - the ability to do work. kinetic energy - the energy of motion (energy of action). potential energy - the energy of position (stored energy). Energy is also categorized by form: light. heat. electrical. mechanical. chemical. © McGraw Hill LLC 5  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. © McGraw Hill LLC 5  Units of Energy Basic Units: calorie or joule. 1 calorie (cal) = 4.184 joules (J). kilocalorie (kcal) = food Calorie. 1 kcal = 1 Calorie = 1000 calories 1 calorie = amount of heat energy required to increase the temperature of 1 gram of water 1 degree C. © McGraw Hill LLC 5  Concentration Concentration: the number or mass of particles of a substance contained in a specified volume. Often used to represent the mixtures of different substances. Concentration of oxygen in the air. Pollen counts. Proper dose of an antibiotic. © McGraw Hill LLC 5  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/cc. © McGraw Hill LLC 5  Density Examples © McGraw Hill LLC Stephen Frisch/McGraw Hill 5  Densities of Some Common Materials Substance Density (g/mL) Substance Density (g/mL) Air 0.00129 (at 0 degrees C) Mercury 13.6 Ammonia 0.000771 (at 0 degrees C) Methanol 0.792 Benzene 0.879 Milk 1.028 to 1.035 Blood 1.060 Oxygen 0.00143 (at 0 degrees C) Bone 1.7 to 2.0 Column 3, Rubber substance, and 0.9 to 1.1 Carbon column 0.001963 (at 0 degrees C) 4, density (grams 0.87 Turpentine dioxide per milliliter), is empty for last two rows. Ethanol 0.789 Urine 1.010 to 1.030 Gasoline 0.66 to 0.69 Water 1.000 (at 4 degrees C) Gold 19.3 Water 0.998 (at 20 degrees C) Hydrogen 0.000090 (at 0 degrees C) Wood balsa, least 0.3 to 0.98 dense; ebony and teak, most dense) Kerosene 0.82 Lead 11.3 © McGraw Hill LLC 5  Calculating Density A 2.00 sample of aluminum is found to weigh 5.40 g. Calculate the density in and g/mL. Use the expression: Density (d) = m/V. Substitute information given into the expression: Since = 2.70 g/mL © McGraw Hill LLC 5  Use Density in Calculation Calculate the volume, in mL, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 g. Density can be written as a Conversion Factor. Multiply the Data Given (g) by the Conversion Factor with the unit g in the denominator. © McGraw Hill LLC 6  Density Calculations 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. © McGraw Hill LLC 6  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 4 degrees C. Specific gravity is a unitless term because the 2 units cancel. Often the health industry uses specific gravity to test urine and blood samples. © McGraw Hill LLC 6