Chemistry - Chapter 1 Chemical Foundations PDF

Chapter 1 Chemical Foundations Copyright ©2018 Cengage Learning. All Rights Reserved. Section 1.1 Chemistry: An Overview  CHEMISTRY: Knowing that matter is composed of various types of atoms.  Is the science that study matter, its properties, t...

Chapter 1 Chemical Foundations Copyright ©2018 Cengage Learning. All Rights Reserved. Section 1.1 Chemistry: An Overview  CHEMISTRY: Knowing that matter is composed of various types of atoms.  Is the science that study matter, its properties, the changes that matter undergoes, and the energy associated with these changes. 2 Section 1.1 Chemistry: An Overview Science  Framework for gaining and organizing knowledge  Procedure for processing and understanding certain types of information  Scientific method: The proses of studding natural phenomena, involving observations, forming laws and theories, and testing of theories by experimentation.  Lies at the center of scientific inquiry  Varies based on the nature of a specific problem and the particular investigator involved 3 Section 1.1 Chemistry: An Overview Figure 1.3 - Fundamental Steps of the Scientific Method 4 Section 1.1 Chemistry: An Overview Steps in the Scientific Method 1. Make observations  Qualitative observations do not involve numbers e.g. Sky is blue, water is liquid.  Quantitative observations (measurements) involve both a number and a unit e.g. Water boils at 100◦C, normal heart rate for adults from 60-100 beats/min. 2. Formulate a hypothesis  Hypothesis: Possible explanation for an observation 5 Section 1.1 Chemistry: An Overview Steps in the Scientific Method (continued) 3. Perform experiments to test the hypothesis  Gather new information that would enable the scientist to ascertain the validity of the hypothesis  Experiments always produce new observations that bring the process back to the beginning again 6 Section 1.1 Chemistry: An Overview Scientific Models  Theory (model): Set of tested hypotheses that gives an overall explanation of a natural phenomenon  Explanation of why nature behaves in a certain way  Constantly refined or replaced as more information becomes available  Explains observed natural behaviour in terms of human experiences Section 1.1 Chemistry: An Overview Scientific Models (continued)  Observations  Events that are witnessed and can be recorded  Natural law: A summary of observed (measurable) behavior  Example - Law of conservation of mass, which states that the total mass of materials is unaffected by a chemical change in those materials Section 1.1 Chemistry: An Overview Science: Drawbacks  Focusing on a theory may limit one’s ability to see alternative explanations  Scientists are humans, and humans have prejudices  Science is affected by profit motives, budgets, fads, wars, and religious beliefs Section 1.1 Chemistry: An Overview CONCEPT CHECK!  Which of the following is an example of a quantitative observation? a. Solution A is a darker red color than solution B b. The grass is green c. Substance A has a greater mass than substance B d. The temperature of the water is 45°C Section 1.1 Chemistry: An Overview Measurement  Consists of two essential parts, a number and a scale (unit) Examples: 20 grams 6.63 × 10-34 joule/second (W)  Standard systems of units  English system - Used in the United States  Metric system  SI system (International System)  Based on the metric system and units derived from the metric system Section 1.1 Chemistry: An Overview Table 1.1 - Fundamental SI Units Section 1.1 Chemistry: An Overview Table 1.2 - Prefixes Used in the SI System Section 1.3 Units of Measurement Prefixes Used in the SI System How many grams in 1 megagram megagram (Mg)? How many meters in 1 nanometer (nm) ? 14 Section 1.3 Units of Measurement Prefixes Used in the SI System Convert 2.6 megagram (Mg) to grams Convert 18.2 nanometer (nm) to meters 15 Section 1.3 Units of Measurement Exercise 1:  Convert from 0.03 megasecond (Ms)to femtosecond(fs).  3x1019 fs. Section 1.3 Units of Measurement Exercise 1: Convert from 1500 yard to meter knowing that 1meter =1.09361 yard  1 m= 1.09361 yard X m= 1500 yard xm=1500 yard/1.09361yard=1371.604137 meter Section 1.3 Units of Measurement Exercise 2: A golfer putted a golf ball 6.8 ft across a green. How many inches does this represent? (1ft=12in)  81.6 in Exercise 3: An iron sample has a mass of 4.50 lb. What is the mass of this sample in grams? (1 kg = 2.2046 lbs)  2.04x103 g Section 1.3 Units of Measurement Exercise 4:  It is estimated that 5g of Uranium are exist per 1000kg of earth crust. At this concentration, what mass of Uranium is present in 1.5mg of the earth crust? A)7.5 Tg (teragram) B) 7.5 µg(microgram) C)7.5ng(nanogram). Section 1.3 Units of Measurement Volume is one of the important physical quantity in chemistry, which is not a fundamental SI unit but it is derived from length. V=L x W x H=m3 1L=1dm3 1mL=1cm 3 20 Section 1.3 Units of Measurement Figure 1.6 - Common Types of Laboratory Equipment Used to Measure Liquid Volume Section 1.3 Units of Measurement Critical Thinking  What if you were not allowed to use units for one day?  How would this affect your life for that day? Section 1.4 Uncertainty in Measurement Certain and Uncertain Digits  Certain digits  Numbers that remain the same regardless of who measures them  Uncertain digits  Digits that must be estimated and therefore vary  While reporting a measurement, record all certain digits plus the first uncertain digit 23 Section 1.4 Uncertainty in Measurement Measurement of Volume Using a Buret  Volume is read at the bottom of the liquid curve, which is called the meniscus  Meniscus of the liquid occurs at about 20.15 mL  Certain digits - 20.1  Uncertain digit - 20.15 24 Section 1.4 Uncertainty in Measurement Uncertainty  A measurement always has some degree of uncertainty  Uncertainty of a measurement depends on the precision of the measuring device  Significant figures: Numbers in which the certain digits and the first uncertain digit are recorded  Uncertainty in the last number is always assumed to be ±1 unless otherwise indicated 25 Section 1.4 Uncertainty in Measurement Example 1.1 - Uncertainty in Measurement  In analyzing a sample of polluted water, a chemist measured out a 25.00-mL water sample with a pipet  At another point in the analysis, the chemist used a graduated cylinder to measure 25 mL of a solution  What is the difference between the measurements 25.00 mL and 25 mL? Section 1.4 Uncertainty in Measurement Precision and Accuracy  Accuracy: Agreement of a particular value with the true value  Precision: Degree of agreement among several measurements of the same quantity  Represents the reproducibility of a given type of measurement Section 1.4 Uncertainty in Measurement Figure 1.8 - The Results of Several Dart Throws Show the Difference between Precise and Accurate Section 1.4 Uncertainty in Measurement Types of Errors  Random error (intermediate error)  Measurement has an equal probability of being low or high  Occurs in estimating the value of the last digit of a measurement  Systematic error (determinate error)  Occurs in the same direction each time  Either always high or always low Section 1.4 Uncertainty in Measurement An Illustration of the Types of Errors Large random Small random Small random errors errors and a errors and no large systematic systematic error error Section 1.4 Uncertainty in Measurement Example 1.2 - Precision and Accuracy  To check the accuracy of a graduated cylinder, a student filled the cylinder to the 25-mL mark using water delivered from a buret and then read the volume delivered Section 1.4 Uncertainty in Measurement Example 1.2 - Precision and Accuracy (continued)  Following are the results of five trials:  Is the graduated cylinder accurate? Why? Section 1.4 Uncertainty in Measurement Join In (3)  The boiling point of a liquid was measured in the lab, with the following results: Trial Boiling point 1 22.0°C ± 0.1 2 22.1°C ± 0.1 3 21.9°C ± 0.1  The actual boiling point of the liquid is 28.7°C Section 1.4 Uncertainty in Measurement Join In (3) (continued)  The results of the determination of the boiling point are: a. accurate and precise b. precise but inaccurate c. accurate but imprecise d. inaccurate and imprecise Section 1.5 Significant Figures and Calculations Uncertainty in Measurements and Significant Figures: In Chemistry there are tow types of numbers: -EXACT NUMBERS: Numbers with defined values such as: Dozen=12, 1 kg=1000 g, 1 inch = 2.54 cm,……..etc. or counting (5 apples ) SUCH EXACT NUMBERS DO NOT CONTAIN SIGNIFICANT FIGURES. -INEXACT NUMBERS: Measured numbers (using tools of measurements) All measured number are inexact because of the uncertainty in the measuring devices and the individual who use them. TO INDICATE THE UNCERTAINETY IN SUCH MEASURED NUMBERS, THESE NUMBERS MUST BE REPORTED USING SIGNIFICANT FIGURES.  A digit that must be estimated in a measurement is called uncertain.  A measurement always has some degree of uncertainty. It is dependent on the precision of the measuring device.  Record the certain digits and the first uncertain digit (the estimated number). Section 1.5 Significant Figures and Calculations Rules for Counting Significant Figures 1. Nonzero integers  Always count as significant figures 2. Zeros - There are three classes of zeros  Leading zeros  Captive zeros  Trailing zeros 36 Section 1.5 Significant Figures and Calculations Rules for Counting Significant Figures (continued) 3. Exact numbers  Determined by counting and not by using a measuring device  Assumed to have an infinite number of significant figures  Can arise from definitions  Example - 2 in 2πr 37 Section 1.5 Significant Figures and Calculations Counting Significant Figures (Sig. Figs.): Classes of Zeroes  Leading zeros  Zeros that precede all the nonzero digits  Do not count as sig. figs.  Example - 0.0025 has only two sig. figs.  Captive zeros  Zeros between nonzero digits  Always count as sig. figs.  Example - 1.008 has four sig. figs. 38 Section 1.5 Significant Figures and Calculations Counting Significant Figures (Sig. Figs.): Classes of Zeroes (continued)  Trailing zeros  Zeros at the right end of the number  Significant only if the number contains a decimal point  Examples -100 has only one sig fig whereas 1.00×102 has three sig. figs. 39 Section 1.5 Significant Figures and Calculations Exponential Notation  Advantages  Number of significant figures can be easily indicated  Fewer zeros are required to write a very large or very small number  Example  0.000060 is much more conveniently represented as 6.0 × 10–5  Number has two significant figures 40 Section 1.5 Significant Figures and Calculations Determining the Number of Significant Figures Section Determining 1.5 the Number of Significant Figures Significant Figures and Calculations Section Determining 1.5 the Number of Significant Figures Significant Figures and Calculations :PROBLEMFor each of the following quantities, underline the zeros that are significant figures (sf), and determine the number of significant figures in each quantity..For (d) to (f), express each in exponential notation first 0.0030 )a( 0.1044 )b( 53,069 )c( L g mL 0.00004715 )d( 57,600. )e( 0.0000007160 )f( m s cm3 SOLUTION : 0.0030 )a( 2sf 0.1044 )b( 4sf 53.069 )c( 5sf L g mL 0.00004715 )d( 57,600. s )e( 5sf 0.0000007160 )f( 4sf m 5.7600x104 cm3 4sf 4.715x10-5 s 7.160x10-7 cm3 m 32.p34: Use exponential notation to express the number 385,500 to a. one significant figure b. two significant figures c. three significant figures d. five significant figures Section 1.5 Significant Figures and Calculations Interactive Example 1.3 - Significant Figures  Give the number of significant figures for each of the following results: a. A student’s extraction procedure on tea yields 0.0105 g of caffeine b. A chemist records a mass of 0.050080 g in an analysis c. In an experiment a span of time is determined to be 8.050×10–3 s Section 1.5 Significant Figures and Calculations Interactive Example 1.3 - Solution a. The number contains three significant figures  The zeros to the left of the 1 are leading zeros and are not significant, but the remaining zero (a captive zero) is significant b. The number contains five significant figures  The leading zeros (to the left of the 5) are not significant  The captive zeros between the 5 and the 8 are significant Section 1.5 Significant Figures and Calculations Interactive Example 1.3 - Solution (continued)  The trailing zero to the right of the 8 is significant because the number contains a decimal point c. This number has four significant figures  Both zeros are significant Section 1.5 Significant Figures and Calculations Rules for Significant Figures in Mathematical Operations  Multiplication or division  Number of significant figures in the result is the same as the number in the least precise measurement used in the calculation 4.56  1.4  6.38  corrected   6.4 Limiting term has Two significant two significant figures figures  Product should have only two significant figures 48 Section 1.5 Significant Figures and Calculations Rules for Significant Figures in Mathematical Operations (continued)  Addition or subtraction  Result has the same number of decimal places as the least precise measurement used in the calculation 12.11  Example 18.0 Limiting term has one decimal place 1.013 31.123  corrected   31.1 One decimal place 49 Section 1.5 Significant Figures and Calculations Rules for Rounding  In a series of calculations, round off only after carrying the extra digits through to the final result  If the digit to be removed is:  Less than 5 - Preceding digit stays the same  Example - 1.33 rounds to 1.3  Greater than or equal to 5 - Preceding digit is increased by 1  Example - 1.36 to 1.4  Do not round sequentially Section 1.5 Significant Figures and Calculations Interactive Example 1.4 - Significant Figures in Mathematical Operations  Carry out the following mathematical operations, and give each result with the correct number of significant figures a. 1.05 × 10–3 ÷ 6.135 b. 21 – 13.8 Section 1.5 Significant Figures and Calculations Interactive Example 1.4 - Significant Figures in Mathematical Operations (continued) c. As part of a lab assignment to determine the value of the gas constant (R), a student measured the pressure (P), volume (V), and temperature (T) for a sample of gas, where PV R T  The following values were obtained:  P = 2.560  T = 275.15  V = 8.8  Calculate R to the correct number of significant figures Section 1.5 Significant Figures and Calculations Interactive Example 1.4 - Solution (continued 1) c. PV 2.560 8.8  R  T 275.15  The correct procedure for obtaining the final result can be represented as follows: 2.560 8.8 = 22.528 = 0.0818753 275.15 275.15 = 0.082 = 8.2×10 2 R Section 1.5 Significant Figures and Calculations Interactive Example 1.4 - Solution (continued 2)  The final result must be rounded to two significant figures because 8.8 (the least precise measurement) has two significant figures  To show the effects of rounding at intermediate steps, carry out the Rounded to twocalculation significant figures as follows: 2.560 8.8  = 22.528 = 23 275.15 275.15 275.15 Section 1.5 Significant Figures and Calculations Interactive Example 1.4 - Solution (continued 3)  Now we proceed with the next calculation 23  0.0835908 275.15  Rounded to two significant figures, this result is 0.084 = 8.4 ×10–2  Note that intermediate rounding gives a significantly different result than that obtained by rounding only at the end Section 1.5 Significant Figures and Calculations Interactive Example 1.4 - Solution (continued 4)  Again, we must reemphasize that in your calculations you should round only at the end  Rounding is carried out at intermediate steps in this text (to always show the correct number of significant figures)  The final answer given in the text may differ slightly from the one you obtain (rounding only at the end) Section 1.5 Significant Figures and Calculations Join In (6)  Express 3140 in scientific notation a. 3.14 × 103 b. 3.14 × 10–3 c. 3.140 × 103 d. 3.140 × 10–3 Section 1.5 Significant Figures and Calculations Join In (7)  After performing a calculation in the lab, the display on your calculator reads “0.023060070”  If the number in the answer is to have five significant figures, what result should you report? a. 0.0230 b. 0.00231 c. 0.023060 d. 0.2367 e. 0.02306 Section 1.5 Significant Figures and Calculations Join In (8)  How many significant figures are in the number 0.03040? a. 1 b. 2 c. 3 d. 4 e. 5 Section 1.9 Density  Mass of substance per unit volume of the substance.  Common units are g/cm3 or g/mL. mass Density = volume 60 Section 1.9 Density Example #1 A certain mineral has a mass of 17.8 g and a volume of 2.35 cm3. What is the density of this mineral? mass Density = volume 17.8 g Density = 2.35 cm3 3 Density = 7.57 g/cm 61 Section 1.9 Density Example #2 What is the mass of a 49.6-mL sample of a liquid, which has a density of 0.85 g/mL? mass Density = volume 62 Section 1.9 Density Example #2 What is the mass of a 49.6-mL sample of a liquid, which has a density of 0.85 g/mL? mass Density = volume x 0.85 g/mL = 49.6 mL mass = x = 42 g 63 Section 1.9 Density Join In (10)  A 25 g cylinder of iron (d = 7.87g/mL) and a 1.0 gram pellet of copper (d = 8.96 g/mL) are placed in 500 mL of water (d = 0.9982 g/mL)  Predict whether each will float or sink in water a. Iron will float, and copper will sink b. Iron will sink, and copper will float c. Iron and copper will sink d. Iron and copper will float e. More information is needed Section 1.9 Density 70 p37: Ethanol and benzene dissolve in each other. When 100.mL of ethanol is dissolved in 1.00L of benzene, what is the mass of the mixture? Ethanol density= 0.789,Benzene density= 0.880 g/cm3 Section 1.10 Classification of Matter Matter  Anything that occupies space and has mass  Has many levels of organization and is complex  Exists in three states  Solid  Liquid  Gas 66 Section 1.10 Classification of Matter Properties of a Solid  Rigid  Fixed volume and shape  Slightly compressible Solid: The water molecules are locked into rigid positions and are close together 67 Section 1.10 Classification of Matter Properties of a Liquid  Definite volume  No specific shape  Assumes the shape of its container  Slightly compressible Liquid: The water molecules are still close together but can move around to some extent 68 Section 1.10 Classification of Matter Properties of a Gas  No fixed volume or shape  Takes on the shape and volume of its container  Highly compressible  Relatively easy to decrease the volume of Gas: The water molecules are far a gas apart and move randomly 69 Section 1.10 Classification of Matter The Organization of Matter 70 Section 1.10 Classification of Matter Definitions for Components of Matter  Element - the simplest type of substance with unique physical and chemical properties. An element consists of only one type of atom (Very small and Indivisible Particles). It cannot be broken down into any simpler substances by physical or chemical means. Molecule - a structure that consists of two or more atoms that are chemically bound together and thus behaves as an independent unit. Helium O2 Section 1.10 Classification of Matter Section 1.10 Classification of Matter Compound - a substance composed of two or more elements which are chemically combined. NaCl,NH3 Mixture - a group of two or more elements and/o compounds that are physically intermingled. Section 1.10 Classification of Matter Chemical Change  A given substance becomes a new substance or substances with different properties and different composition.  Example: Bunsen burner (methane reacts with oxygen to form carbon dioxide and water) 74 Section 1.10 Classification of Matter CONCEPT CHECK! Which of the following are examples of a chemical change?  Pulverizing (crushing) rock salt  Burning of wood  Dissolving of sugar in water  Melting a ice cream on a warm summer day 75 Section 1.10 Classification of Matter The Organization of Matter 76 Section 1.10 Classification of Matter Mixtures  Have variable composition  Classification  Homogeneous mixture: Has visibly indistinguishable parts and is often called a solution  Heterogeneous mixture: Has visibly distinguishable parts  Can be separated into pure substances, which have constant compositions, by physical methods Section 1.10 Classification of Matter Physical Change  Change in the form of a substance, not in its chemical composition.  Example: boiling or freezing water  Can be used to separate a mixture into pure compounds, but it will not break compounds into elements. 78 Section 1.10 Classification of Matter Methods for Separating Components in a Mixture 79 Section 1.10 Classification of Matter CONCEPT CHECK! Which of the following is a homogeneous mixture?  Pure water  Gasoline  Jar of jelly beans  Soil  Copper metal 80 Section 1.10 Classification of Matter Join In (11)  A solution is also a: a. heterogeneous mixture b. homogeneous mixture c. compound d. distilled mixture e. pure mixture Section 1.10 Classification of Matter 87 p38: Classify following as physical or chemical changes. a.Moth balls gradually vaporize in a closet. b. Hydrofluoric acid attacks glass, and is used to etch calibration marks on glass laboratory utensils. c. A French chef making a sauce with brandy is able to burn off the alcohol from the brandy, leaving just the brandy flavouring. d. Chemistry majors sometimes get holes in the cotton jeans they wear to lab because of acid spills.