© Adrian Dingle’s Chemistry Pages 2004, 2005, 2006, 2007, 2008, 2009, 2010, 2011, 2012. All rights reserved. These materials may NOT be copied or redistributed.

Slides:



Advertisements
Similar presentations
Measurement and Significant Figures
Advertisements

Chapter 2 Measurement and Problem Solving
Chapter 2 Measurements and Calculations.
Chem101 Chapter 01 Chemical Foundations.
Measurements and Calculations
Measurements and Calculations Chapter 2 2.
Measurements & Calculations
Measurement & Conversions
Measurements and Calculations
Copyright©2004 by Houghton Mifflin Company. All rights reserved 1 Introductory Chemistry: A Foundation FIFTH EDITION by Steven S. Zumdahl University of.
Chapter 1: Introduction
Chapter 2 The Metric System
I. Scientific Method. The Scientific Method A logical approach to solving problems or answering questions. Starts with observation- noting and recording.
Introduction to Chemistry.  No eating or drinking!  Wear goggles at all times!  Use common sense!
Measurements and Calculations
Zumdahl • Zumdahl • DeCoste
1. To show how very large or very small numbers can be expressed in scientific notation 2. To learn the English, metric, and SI systems of measurement.
1 Measurement Quantitative Observation Comparison Based on an Accepted Scale –e.g. Meter Stick Has 2 Parts – the Number and the Unit –Number Tells Comparison.
Introductory Chemistry: A Foundation, 6 th Ed. Introductory Chemistry, 6 th Ed. Basic Chemistry, 6 th Ed. by Steven S. Zumdahl, Donald J. DeCoste University.
Why do we need it? Because in chemistry we are measuring very small things like protons and electrons and we need an easy way to express these numbers.
Unit 2. Measurement This lesson is 8 days long.
Chapter 2 Measurements and Calculations. Chapter 2 Table of Contents Return to TOC Copyright © Cengage Learning. All rights reserved 2.1 Scientific Notation.
Chemistry Chapter 2 MeasurementsandCalculations. Steps in the Scientific Method 1.Observations - quantitative - qualitative 2.Formulating hypotheses -
Measuring and Units.
. Do Now: 1. Differentiate between qualitative and quantitative observations/data. 2. True or False (explain why): A theory can be proven correct or incorrec.
Chapter 1 Matter and Measurement. What is Chemistry? The study of all substances and the changes that they can undergo The CENTRAL SCIENCE.
Measurements and Calculations 1. To show how very large or very small numbers can be expressed in scientific notation 2. To learn the English, metric,
I. Using Measurements (p )
Accuracy and Precision Accuracy refers to the how close you are to the actual value. Precision refers to the how close your measurements are to each other.
Chapter 2: Scientific Method Cartoon courtesy of NearingZero.net.
3.1 Measurements and Their Uncertainty
Measurements in Chemistry MeasurementsandCalculations.
Section 2.1 Units and Measurements
Measurements & Calculations Chapter 2. Nature of Measurement Measurement - quantitative observation consisting of two parts: Part 1 - number Part 2 -
Section 5.1 Scientific Notation and Units 1.To show how very large or very small numbers can be expressed in scientific notation 2.To learn the English,
Ch. 2.1 Scientific Method. 2.1 Goals 1. Describe the purpose of the scientific method. 2. Distinguish between qualitative and quantitative observations.
Foundations of Chemistry. Prefixes l Tera-T1,000,000,000, l giga- G 1,000,000, l mega - M 1,000, l kilo - k 1, l deci-d0.1.
Chapter One Chemical Foundations. Section 1.1 Chemistry an Overview Macroscopic World Macroscopic World Microscopic World Microscopic World Process for.
Chapter 3. Measurement Measurement-A quantity that has both a number and a unit. EX: 12.0 feet In Chemistry the use of very large or very small numbers.
Copyright©2004 by Houghton Mifflin Company. All rights reserved 1 Introductory Chemistry: A Foundation FIFTH EDITION by Steven S. Zumdahl University of.
INTRODUCTION TO CHEMISTRY CHAPTERS 1 AND 2. 1.) WHAT IS CHEMISTRY?  The study of matter and the changes that matter undergoes.
I. Using Measurements MEASUREMENT IN SCIENCE. A. Accuracy vs. Precision Accuracy - how close a measurement is to the accepted value Precision - how close.
Ch. 3, Scientific Measurement. Measurement Measurement: A quantity that has a number and a unit. Like 52 meters.
I II III I. Using Measurements MEASUREMENT. A. Accuracy vs. Precision  Accuracy - how close a measurement is to the accepted value  Precision - how.
Ch. 3, Scientific Measurement. Measurement : A quantity that has a and a. Like 52 meters.
Matter And Measurement 1 Matter and Measurement. Matter And Measurement 2 Length The measure of how much space an object occupies; The basic unit of length,
Chemistry and Calculations Chemistry Honors 2 Accuracy & Precision Precision: how closely individual measurements compare with each other Accuracy: how.
Objectives Describe the purpose of the scientific method. Distinguish between qualitative and quantitative observations. Describe the differences between.
Scientific Notation & Significant Figures in Measurement.
Conversions and Significant Figures Kelley Kuhn Center for Creative Arts.
Units of Measure & Conversions. Number vs. Quantity  Quantity - number + unit UNITS MATTER!!
I II III I. Using Measurements MEASUREMENT. A. Accuracy vs. Precision  Accuracy - how close a measurement is to the accepted value  Precision - how.
CHEMISTRY CHAPTER 2, SECTION 3. USING SCIENTIFIC MEASUREMENTS Accuracy and Precision Accuracy refers to the closeness of measurements to the correct or.
Intro Chem Grading and Class Policy handout Book : Introductory Chemistry, 5 th ed., Zumdahl Website:
1 CHEMISTRY 101 Dr. IsmailFasfous  Textbook : Raymond Chang, 10th Edition  Office Location: Chemistry Building, Room 212  Office Telephone: 4738 
Section 5.1 Scientific Notation and Units Steven S. Zumdahl Susan A. Zumdahl Donald J. DeCoste Gretchen M. Adams University of Illinois at Urbana-Champaign.
Measurements and Calculations Scientific Method Units of Measurement Using Scientific Measurements.
CH. 2 - MEASUREMENT. Observing and Collecting Data Data may be Qualitative (descriptive) Flower is red Quantitative (numerical) 100 flowers.
WHAT WE HAVE LEARNED. SCIENTIFIC NOTATION 1. Move the decimal to the right of the first non-zero number. 2. Count how many places the decimal had to.
The scientific method is a logical approach to solving problems by observing and collecting data, formulating hypotheses, testing hypotheses, and formulating.
Uncertainty in Measurement A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty.
Objectives Describe the purpose of the scientific method. Distinguish between qualitative and quantitative observations. Describe the differences between.
Data Analysis. Scientific Method Not covered in class: Review.
Measurements and Calculations Scientific Method Units of Measurement Using Scientific Measurements.
Chapter 2: Measurements and Calculations
Measurement.
Objectives To show how very large or very small numbers can be expressed in scientific notation To learn the English, metric, and SI systems of measurement.
Chapter 2 Table of Contents Section 1 Scientific Method
Measurements and Calculations
TOPIC 0B: Measurement.
Presentation transcript:

© Adrian Dingle’s Chemistry Pages 2004, 2005, 2006, 2007, 2008, 2009, 2010, 2011, All rights reserved. These materials may NOT be copied or redistributed in any way, except for individual class instruction.

What is Chemistry? Chemistry can be described as the science that deals with matter, and the changes that matter undergoes. It is sometimes called the central science because so many naturally occurring phenomena involve chemistry and chemical change.

Scientific Problem Solving Scientific (logical) problem solving involves several steps: 1.State the problem and make observations. Observations can be quantitative (those involving numbers or measurement) or qualitative (those not involving numbers). 2.Formulate a possible solution (this is a hypothesis). 3.Perform experiments to test the hypothesis. The results and observations from these experiments lead to the modification of the hypothesis and therefore further experiments.

Scientific Problem Solving Eventually after several experiments the hypothesis may graduate to become a theory. A theory gives a universally accepted explanation of the problem. Of course, theories should be constantly challenged and may be refined as and when new data and new scientific evidence comes to light.

Scientific Problem Solving Theories are different to laws. Laws state what general behavior is observed to occur naturally. For example The Law of Conservation of Mass exists since it has been consistently observed that during all chemical changes mass remains unchanged (i.e. it is neither created nor destroyed).

Measurement Measurements, and subsequently calculations, allow the determination of some of the quantitative properties of a substance. For example, mass and density.

Scientific notation Measurements and calculations in chemistry often require the use of very large or very small numbers. In order to make handling them easier, such numbers can be expressed using scientific notation. All numbers expressed in this manner are represented by a number between 1 and 10 multiplied by 10 raised to a power.

Scientific notation The number of places the decimal point has moved determines the power of 10. If the decimal point has moved to the left then the power is positive, to the right, negative.

For example, the number is converted to scientific notation by using the number 4.2. In the process the decimal point has moved four places to the left, so the power of 10 used is = 4.2 x 10 4 The number on the other hand, is converted to scientific notation by using the number 1.2. In the process the decimal point has moved four places to the right, so the power of 10 used is = 1.2 x 10 -4

Scientific Notation Practice Convert the following numbers to scientific notation. (a) (b) 356 (c) (d) (e) 12200

Scientific Notation Practice Convert the following scientific notation numbers to non-scientific notation numbers. (a) 4.2 x 10 3 (b) 2.15 x (c) 3.14 x (d) 9.22 x 10 5 (e) 9.57 x 10 2

SI Units Units tell us the scale that is being used for measurement. Prefixes are used to make writing very large or small numbers easier.

SI Units Common SI units and prefixes are given below. Base Quantity Name of Unit Symbol Mass Kilogram kg Length Meter m Time Second s Amount of Substance Mole mol Temperature Kelvin K

Metric System PrefixSymbolMeaning GigaG10 9 MegaM10 6 Kilok10 3 Base-10 0 Decid10 -1 Centic10 -2 Millim10 -3 Microu10 -6 Nanon10 -9 Picop10 -12

Converting Units One unit can be converted to another by using a conversion factor. Application of the simple formula below will allow the conversion of one unit to another. This method of converting between units is called the factor-labeling method or dimensional analysis.

Converting Units (Unit a) (conversion factor) = Unit b The conversion factor is derived from the equivalence statement of the two units. For example, in the equivalence of 1.00 inches = 2.54 cm, the conversion factor will either be,

Converting Units The correct choice is the one that allows the cancellation of the unwanted units. For example, to convert 9.00 inches into cm, do the following calculation: 9 inch 2.54 cm 1.00 inch = 22.9 cm

Converting Units To convert 5.00 cm into inches, do the following calculation: 5.00 cm 1.00 inch 2.54 cm = 1.97 inches

Temperature There are three scales of temperature that you may come across in your study of chemistry. They are Celsius ( o C), Fahrenheit ( o F) and Kelvin (K). The conversion factors on the next slide will be useful.

Temperature

Examples Convert the following quantities from one unit to another using the following equivalence statements. Clearly show working m = yd mile = 1760 yd kg = lbs (a) 30 m to miles (b) 1500 yd to miles (c) 206 miles to m (d) 34 kg to lbs (e) 34 lb to kg

Examples Convert the following temperatures from one unit to the other. (a) 263 K to o F (b) 38 K to o F (c) 13 o F to o C (d) 1390 o C to K (e) 3000 o C to o F

Derived Units All other units can be derived from base quantities. One such unit that is very important in chemistry is volume. Volume has the unit length 3. Common units for volume are Liters (L) or milliliters (mL) mL = cm L = mL = cm 3 = dm 3

Density Density is the ratio of the mass of an object to its volume.

Density Examples A small gold nugget has volume of 0.87 cm 3. What is its mass if the density of gold is 19.3 g/cm 3 ? What volume is occupied by 35.2 g of carbon tetrachloride if its density is 1.60 g/mL?

Uncertainty, Significant figures and Rounding off When reading the scale on a piece of laboratory equipment such as a graduated cylinder or a buret, there is always a degree of uncertainty in the recorded measurement. The reading will often fall between two divisions on the scale and an estimate must be made in order to record the final digit. This estimated final digit is said to be uncertain and is reflected in the recording of the numbers by using +/-. All those digits that can be recorded with certainty are said to be certain. The certain and the uncertain numbers taken together are called significant figures.

Determining the number of significant figures present in a number 1.Any non-zero integers are always counted as significant figures. 2.Leading zeros are those that precede all of the non-zero digits and are never counted as significant figures. 3.Captive zeros are those that fall between non-zero digits and are always counted as significant figures.

Determining the number of significant figures present in a number 4. Trailing zeros are those at the end of a number and are only significant if the number is written with a decimal point. 5. Exact numbers have an unlimited number of significant figures. (Exact numbers are those which are as a result of counting e.g. 3 apples or by definition e.g kg = lb). 6. In scientific notation the 10 x part of the number is never counted as significant.

Determining the correct number of significant figures to be shown as the result of a calculation 1.When multiplying or dividing. Limit the answer to the same number of significant figures that appear in the original data with the fewest number of significant figures. 2. When adding or subtracting. Limit the answer to the same number of decimal places that appear in the original data with the fewest number of decimal places.

Rounding off In a multi-step calculation it is possible to leave the rounding until the end i.e. leave all numbers on the calculator in the intermediate steps, or round to the correct number of figures in each step, or round to an extra figure in each intermediate step and then round to the correct number of significant figures at the end of the calculation.

Rounding off Whichever method is being employed, use the simple rule that if the digit directly to the right of the final significant figure is: less than 5 then the preceding digit stays the same. perfectly equal to 5 (without any more nonzero digits) then the preceding digit should be rounded to the nearest even digit 5 with nonzero digits trailing then the preceding digit should be increased by one.

Sig Fig Examples Determine the number of significant figures in the following numbers. (a) (b) (c) 1024 (d) 4.7 x (e)

Sig Fig Examples Using a calculator carry out the following calculations and record the answer to the correct number of significant figures. (a) 34.5 x (b) 123 / 3 (c) 2.61 x x 356 (d) (e) (f)

Accuracy and Precision A distinction should be made between accuracy and precision. Accuracy relates to how close the measured value is to the actual value of the quantity. Precision refers to how close two or more measurements of the same quantity are to one another.

Accuracy and Precision Consider three sets of data that have been recorded after measuring a piece of wood that is exactly m long. (a) Which set of data is the most accurate? (b) Which set of data is the most precise?

Percentage Error The data that is derived from experiments will often differ from the accepted, published, actual value. When this occurs, a common way of expressing accuracy is percentage error.

Percentage Error Determine the percentage error in the following problem. Show all your work! Experimental Value: 1.14 g Actual Value: 1.30 g