Mec chapter 1
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This document provides an overview of key concepts in chemistry, including: 1. Chemistry is the study of matter, its properties, and the changes it undergoes. The scientific method is used to systematically study matter through observation, questioning, experimentation and summarization. 2. Matter can exist in three physical states - solid, liquid, and gas. It also has physical and chemical properties and can undergo physical or chemical changes. 3. The metric system is the standard system of measurement in chemistry. It relates units decimally and is easier for conversion than the English system. Careful measurement and use of appropriate units is important in chemistry.
Mec chapter 1
- 1. Copyright© The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 1 Chemistry: Methods and Measurement Deniston Topping Caret 7th 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. 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. 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. 1.1 The Discovery Process
- 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. 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. Three States of Water (a) Solid (b) Liquid (c) Gas
- 9. 1.2 Matter and Properties Comparison of the Three Physical States
- 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. 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.
- 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. 1.2 Matter and Properties Classify the following as either a c h e m i c a l o r p h
- 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. 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. 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. 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. 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. 1.2 Matter and Properties Classes of Matter
- 20. 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
- 21. English and Metric Units • English system - a collection of 1.3 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. Basic Units of the Metric System 1.3 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. 1.3 Measurement in Chemistry
- 24. UNIT CONVERSION 1.3 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. • Let your units do the work for you by 1.3 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. 12 donuts =1 1.3 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. • We use these two mathematical facts to 1.3 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. 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. Common English System Units 1.3 Measurement in Chemistry • Convert 12 gallons to units of quarts
- 30. Intersystem Conversion Units 1.3 Measurement in Chemistry • Convert 4.00 ounces to kilograms
- 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. 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. and Scientific Notation 1.4 Significant Figures Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit
- 34. and Scientific Notation Recognition of Significant Figures 1.4 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. Use of Zeros in Significant and Scientific Notation 1.4 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. and Scientific Notation 1.4 Significant Figures How many significant figures are in the following? 1. 3.400 2. 3004 3. 300. 4. 0.003040
- 37. and Scientific Notation Scientific Notation 1.4 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. and Scientific Notation • To convert a number greater than 1 to 1.4 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. and Scientific Notation 1.4 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. and Scientific Notation Types of Uncertainty 1.4 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. Significant Figures in Calculation of and Scientific Notation Results 1.4 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. Rules for Multiplication and Division and Scientific Notation 1.4 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. and Scientific Notation Exact and Inexact Numbers 1.4 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. Rules for Rounding Off Numbers and Scientific Notation 1.4 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. and Scientific Notation 1.4 Significant Figures How Many Significant Figures? Round off each number to 3 significant figures: 1. 61.40 2. 6.171 3. 0.066494
- 46. 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. 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. 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. 1.5 Experimental Quantities The milliliter (mL) and the cubic centimeter (cm3) are equivalent
- 50. 1.5 Experimental Quantities • Time - metric unit is the second • Temperature - the degree of “hotness” of an object
- 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. 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. 1.5 Experimental Quantities • Energy
- 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. 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. 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. 1.5 Experimental Quantities cork water brass nut liquid mercury
- 59. 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
- 60. 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.
- 61. 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)