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MEC Chapter 1

MEC Chapter 1

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Mec chapter 1 Mec chapter 1 Presentation Transcript

  • Copyright© The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 1 Chemistry: Methods and MeasurementDenistonToppingCaret7th Edition
  • 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
  • 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
  • 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
  • 1.1 The Discovery Process
  • 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
  • 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
  • Three States of Water(a) Solid (b) Liquid (c) Gas
  • 1.2 Matter and Properties Comparison of the Three Physical States
  • 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
  • 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.
  • 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
  • 1.2 Matter and Properties Classify the following as either a c h e m i c a l o r p h
  • 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
  • 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
  • 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
  • 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
  • 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
  • 1.2 Matter and Properties Classes of Matter
  • 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
  • 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
  • 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
  • 1.3 Measurement in Chemistry
  • 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 !!!!!!!!!!!
  • • 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
  • 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
  • • 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.
  • 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...
  • Common English System Units1.3 Measurement in Chemistry • Convert 12 gallons to units of quarts
  • Intersystem Conversion Units1.3 Measurement in Chemistry • Convert 4.00 ounces to kilograms
  • 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
  • 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
  • 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
  • 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
  • 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
  • and Scientific Notation1.4 Significant Figures How many significant figures are in the following? 1. 3.400 2. 3004 3. 300. 4. 0.003040
  • 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
  • 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
  • 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
  • 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
  • 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
  • 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
  • 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
  • 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
  • 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
  • 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)
  • 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
  • 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
  • 1.5 Experimental Quantities The milliliter (mL) and the cubic centimeter (cm3) are equivalent
  • 1.5 Experimental Quantities • Time - metric unit is the second • Temperature - the degree of “hotness” of an object
  • 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
  • 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
  • 1.5 Experimental Quantities • Energy
  • 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
  • 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.
  • 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
  • 1.5 Experimental Quantities cork water brass nut liquid mercury
  • 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
  • 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.
  • 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)