Chapter 1.pdf Chemical Foundations

Chapter 1 Chemical Foundations Section 1.1 ModernAn Chemistry: Chemistry: Overview A Brief Glimpse Health and Medicine Sanitation systems Surgery with anesthesia Vaccines and antibiotics Gene therapy Energy and the En...

Chapter 1 Chemical Foundations Section 1.1 ModernAn Chemistry: Chemistry: Overview A Brief Glimpse Health and Medicine Sanitation systems Surgery with anesthesia Vaccines and antibiotics Gene therapy Energy and the Environment Fossil fuels Solar energy Nuclear energy 2 Section 1.1 ModernAn Chemistry: Chemistry: Overview A Brief Glimpse Materials and Technology Polymers, ceramics, liquid crystals Room-temperature superconductors? Molecular computing? Food and Agriculture Genetically modified crops “Natural” pesticides Specialized fertilizers 3 Section 1.1 Chemistry: An Overview  A main challenge of chemistry is to understand the connection between the macroscopic world that we experience and the microscopic world of atoms and molecules.  You must learn to think on the atomic level. Section 1.1 Chemistry: An Overview Atoms vs. Molecules  Matter: is composed of tiny particles called atoms.  Atom: smallest part of an element.  Molecule: Two or more atoms joined and acting as a unit. Section 1.1 Chemistry: An Overview Oxygen and Hydrogen Molecules Use subscripts when more than one atom is in the molecule. Section 1.1 Chemistry: An Overview A Chemical Reaction  One substance changes to another by reorganizing the way the atoms are attached to each other. Section 1.2 The Scientific Method Science  Science is a framework for gaining and organizing knowledge.  Science is a plan of action — a procedure for processing and understanding certain types of information.  Scientists are always challenging our current beliefs about science, asking questions, and experimenting to gain new knowledge.  Scientific method is needed. Section 1.2 The Scientific Method Fundamental Steps of the Scientific Method Process that lies at the center of scientific inquiry. Section 1.2 1.2Scientific The Experiment and Explanation Method An tnemirepxe is an observation of natural phenomena carried out in a controlled manner so that the results can be duplicated and rational conclusions obtained. A law is a concise statement or mathematical equation about a fundamental relationship or regularity of nature. A hypothesis is a tentative explanation of some regularity of nature. A theory is a tested explanation of basic natural phenomena. Example: molecular theory of gases. Note: We cannot prove a theory absolutely. It is always possible that further experiments will show the theory to be limited or that someone will develop a better theory 1 0 Section 1.2 The Scientific Method Scientific Models Law A summary of repeatable observed (measurable) behavior. Hypothesis A possible explanation for an observation. Theory (Model) Set of tested hypotheses that gives an overall explanation of some natural phenomenon. Section 1.3 Units of Measurement Nature of Measurement Measurement Quantitative observation consisting of two parts.  number  scale (unit) Examples  20 grams  6.63 × 10-34 joule·second Section 1.3 Units of Measurement The Fundamental SI Units Physical Quantity Name of Unit Abbreviation Mass kilogram kg Length meter m Time second s Temperature kelvin K Electric current ampere A Amount of substance mole mol Luminous intensity candela cd Section 1.3 Units of Measurement Prefixes Used in the SI System Prefixes are used to change the size of the unit. Section 1.3 Units of Measurement Prefixes Used in the SI System Section 1.3 Units of Measurement Mass ≠ Weight  Mass is a measure of the resistance of an object to a change in its state of motion. Mass does not vary.  Weight is the force that gravity exerts on an object. Weight varies with the strength of the gravitational field. Section 1.4 Uncertainty in Measurement  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 1.4 Uncertainty in Measurement Measurement of Volume Using a Buret  The volume is read at the bottom of the liquid curve (meniscus).  Meniscus of the liquid occurs at about 20.22 mL.  Certain digits: 20.22  Uncertain digit: 20.22 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. Section 1.4 Uncertainty in Measurement Precision and Accuracy Section 1.5 Significant Figures and Calculations Rules for Counting Significant Figures 1. Nonzero integers always count as significant figures.  3456 has 4 sig figs (significant figures). Section 1.5 Significant Figures and Calculations Rules for Counting Significant Figures 2. There are three classes of zeros. a. Leading zeros are zeros that precede all the nonzero digits. These do not count as significant figures.  0.048 has 2 sig figs. Section 1.5 Significant Figures and Calculations Rules for Counting Significant Figures b. Captive zeros are zeros between nonzero digits. These always count as significant figures.  16.07 has 4 sig figs. Section 1.5 Significant Figures and Calculations Rules for Counting Significant Figures c. Trailing zeros are zeros at the right end of the number. They are significant only if the number contains a decimal point.  9.300 has 4 sig figs.  150 has 2 sig figs. Section 1.5 Significant Figures and Calculations Section 1.5 Significant Figures and Calculations Exponential Notation  Section 1.5 Significant Figures and Calculations How many significant figures does each of the following numbers have? # of Sig. Figs. Scientific Notation 1. 413.97 5 102 × 4.1397 2. 0.0006 1 4–10 ×6 3. 5.120063 7 5.120063 4. 161,000 3 105 × 1.61 5. 03600. 4 103 ×3.600 Section 1.5 Significant Figures and Calculations Q)Round each of the following to three significant figures.dedeen erehw noitaton cifitneics esU. 1. 37.459 37.5 or 3.75 × 101 2. 5431978 5.43 × 106 3. 132.7789003 133 or 1.33 × 102 4. 0.00087564 8.76 × 10–4 Section 1.5 Significant Figures and Calculations Q) Round 0.00564458 to four significant figures and express the answer using scientific notation. A. 5.64  10-2 B. 5.000  10-3 C. 5.645  10-4 D. 0.56446 E. 5.645  10-3 Section 1.5 Significant Figures and Calculations Significant Figures in Mathematical Operations 1. For multiplication or division, the number of significant figures in the result is the same as the number in the least precise measurement used in the calculation. 1.342 × 5.5 = 7.381  7.4 Section 1.5 Significant Figures and Calculations e.g., 10.54 × 31.4 × 16.987 =5621.9 =5.62×103 4 S.F. ×3 S.F. × 5 S.F. = 5 S.F. = 3 S.F. e.g. 5.896 ÷ 0.008 = 737 = 7 x 102 4 S. F. ÷ 1 S.F. = 3 S.F.= 1 S.F. 31 Section 1.5 Significant Figures and Calculations Give the value of the following calculation to the correct number of significant figures.  635.4  0.0045   2.3589  A. 1.21213 B. 1.212 C. 1.212132774 D. 1.2 E. 1 32 Section 1.5 Significant Figures and Calculations Significant Figures in Mathematical Operations 2. For addition or subtraction, the result has the same number of decimal places as the least precise measurement used in the calculation. 23.445  7.83 31.275  Corrected 31.28 Section 1.5 Significant Figures and Calculations Addition and Subtraction Answer has same number of decimal places as quantity with fewest number of decimal places. 12.9753 4 decimal places e.g., 319.5+ 1 decimal place 4.398+ 3 decimal places 336.9 1 decimal place 397 0decimal places e.g., 273.15– 2 decimal places 124 0 decimal place Section 1.5 Significant Figures and Calculations Q) For each calculation, give the answer to the correct number of significant figures. 1. 10.0 g + 1.03 g + 0.243 g = 11.3 g or 2. °19.556 C– 19.552 °C = 0.00400C ° or 3. 327.5 m × 4.52 m = 1480.3 = 1.48 × 103 m2 4. 15.985 g ÷ 24.12 mL = 0.6627 g/mL or Section 1.5 Significant Figures and Calculations Q) When the expression, 412.272 + 0.00031 – 1.00797 + 0.000024 +12.8 is evaluated, the result should be expressed as: A. 424.06 B. 424.064364 C. 424.1 D. 424.064 E. 424 Section 1.5 Significant Figures and Calculations Q) For the following calculations, give the answer to the correct number of significant figures. ( 71.359m  71.357m) 0.00200 m 1.  (3.2 s × 3.67 s) 11.744 = (0.00200/12)=(1.666 × 10–4 ) =1.7× 10–4 m/s2 2. 13.674 cm × 4.35 cm × 0.35 cm  856 s + 1531.1 s 20.818665 cm3 )21/2387( = 0.0088 cms/3 30 2387.1 s Or 8.8 x 3-10 Section 1.5 Significant Figures and Calculations Rules for Counting Significant Figures 3. Exact numbers have an infinite number of significant figures.  1 inch = 2.54 cm, exactly.  9 pencils (obtained by counting). Section 1.5 Significant Figures and Calculations )1)Numbers that come from definitions. 12 in. = 1 ft 60 s = 1 min )2(Numbers that come from direct count –Number of people in small room Have no uncertainty Assume they have infinite number of significant figures The number of significant figures in a calculation result depends only on the numbers of significant figures in quantities having uncertainties Section 1.5 Significant Figures and Calculations Q) If you have 9 coins in a jar. What is the total mass of the 9 coins when each coin has a mass of 3.0 grams? 3.0g x 9=27 g The number 9 is exact and does not determine the number of significant figures Q) How many feet are there in 36.00 inches? Express the answer with the correct number A. 3 ft. of significant figures: 1 ft.=12 in B. 3.0 ft C. 3.00 ft. D. 3.000 ft. E. 3.00000 ft. Section 1.5 Significant Figures and Calculations a = 3.6 b = 5.01 c = 0.37 d=-18 Section 1.5 Significant Figures and Calculations CONCEPT CHECK! You have water in each graduated cylinder shown. You then add both samples to a beaker (assume that all of the liquid is transferred). How would you write the number describing the total volume? 3.1 mL What limits the precision of the total volume? Section 1.6 Learning to Solve Problems Systematically Questions to ask when approaching a problem What is my goal? What do I know? How do I get there? Section 1.7 Dimensional Analysis  Use when converting a given result from one system of units to another.  To convert from one unit to another, use the equivalence statement that relates the two units.  Derive the appropriate unit factor by looking at the direction of the required change (to cancel the unwanted units).  Multiply the quantity to be converted by the unit factor to give the quantity with the desired units. Section 1.7 Dimensional Analysis Example #1 A golfer putted a golf ball 6.8 ft across a green. How many inches does this represent? To convert from one unit to another, use the equivalence statement that relates the two units. 1 ft = 12 in The two unit factors are: 1 ft 12 in and 12 in 1 ft Section 1.7 Dimensional Analysis Example #1 A golfer putted a golf ball 6.8 ft across a green. How many inches does this represent?  Derive the appropriate unit factor by looking at the direction of the required change (to cancel the unwanted units). 12 in 6.8 ft   in 1 ft Section 1.7 Dimensional Analysis Example #1 A golfer putted a golf ball 6.8 ft across a green. How many inches does this represent?  Multiply the quantity to be converted by the unit factor to give the quantity with the desired units. 12 in 6.8 ft   82 in 1 ft Section 1.7 Dimensional Analysis Example #2 An iron sample has a mass of 4.50 lb. What is the mass of this sample in grams? (1 kg = 2.2046 lbs; 1 kg = 1000 g) 1 kg 1000 g 4.50 lbs   = 2.04  103 g 2.2046 lbs 1 kg Section 1.7 Dimensional Analysis Dimensional Analysis Method of Solving Problems 1. Determine which unit conversion factor(s) are needed 2. Carry units through calculation 3. If all units cancel except for the desired unit(s), then the problem was solved correctly. given quantity x conversion factor = desired quantity given desired unit given unit x = desired unit given unit 22 A person’s average daily intake of glucose (a form of sugar) is 0.0833 pound (lb). What is this mass in milligrams (mg)? (1 lb = 453.6 g.) Volume – SI derived unit for volume is cubic meter (m3) (1 m)3 = (100 cm)3 (1 m)3 =(10 dm)3 1 L = 1000 mL = 1000 cm3 = 1 dm3 1 mL = 1 cm3 1 L = 1 dm3 103 L = 1 m3 21 Q) A liquid helium storage tank has a volume of 275 L. What is the volume in m3? 275 L × 1 m3 = 0.275 m3 103 L Q) The density of liquid nitrogen at its boiling point (−196°C or 77 K) is 0.808 g/cm3. Convert the density to units of kg/m3. The world’s oceans contain approximately 1.35×109 km3 of water. What is this volume in litters? Example: Convert 3.5 m3 to cm3 Example: Convert 0.097 m to mm. Q) How many centimetres are there in 6.51 miles? 1 mi = 5280 ft 1 ft = 12 in 1 in = 2.54 cm Q) Convert speed of light from 3.00 × 108 m/s to mi/hr (1 mi = 1.609 km) = 671 = 6.71×102 1 mi = 5280 ft 1 ft = 12 in 1 in = 2.54 cm =670953813 = 671 = 6.71×102  The Toyota Camry hybrid electric car has a gas mileage rating of 56 miles per gallon. What is this rating expressed in units of kilometers per liter? 1 gal = 3.784 L 1 mile = 1.609 km mi 1 gal  1.609km 56  A. 2.38 × 101 km L–1 gal 3.784L 1 mi B. 24 km L–1 C. 23.8 km L–1 D. 2.4 × 101 km L–1 E. 9.2 km L–1 The volume of a basketball is 433.5 in3. Convert this to mm3. (1 in. = 2.54 cm) 3 3 3 A. 1.101  10-2 mm3 B. 7.104  106 mm3 C. 7.104  104 mm3 D. 1.101  104 mm3 E. 1.101  106 mm3 Section 1.7 Dimensional Analysis CONCEPT CHECK! What data would you need to estimate the money you would spend on gasoline to drive your car from New York to Los Angeles? Provide estimates of values and a sample calculation. Section 1.8 Temperature Three Systems for Measuring Temperature  Fahrenheit  Celsius  Kelvin Section 1.8 Temperature The Three Major Temperature Scales Section 1.8 Temperature Converting Between Scales TK  TC + 273.15 TC  TK  273.15  TC  TF  32 F  5C 9F TF  TC  9F 5C + 32 F Section 1.8 Temperature EXERCISE! At what temperature does °C = °F?  TF = TC  Assume: TF = TC = X Section 1.8 Temperature EXERCISE!  Since °C equals °F, they both should be the same value (designated as variable x).  Use one of the conversion equations such as:  TC  TF  32 F  5C 9F  Substitute in the value of x for both TC and TF. Solve for x. Section 1.8 Temperature EXERCISE!  TC  TF  32 F  5C 9F x  x  32 F  5C 9F x   40 So  40C =  40F Section 1.8 Temperature 1. Convert 121 ˚F to the Celsius scale.  9 TF = 5 x TC + 32 TC  (TF  32 (ºF)) 5 (℃) 9 (ºF)   Tc  121ºF 32 ºF 5 ºC 9ºF = 49.4 ºC 2. Convert 121 ˚F to the Kelvin scale. – We already have in °C so… 1K 1K TK  ( Tc  273.15ºC)  (49.4  273.15C) 1 ºC 1 ºC TK = 332.55 K  3. Convert 77 K to the Celsius scale. 1K 1 ºC TK  (TC  273.15 ºC) TC  (TK  273.15 K) 1 ºC 1K 1C TC  (77K  273.15 K) = –196ºC 1K 4. Convert 77.0 K to the Fahrenheit scale. – We already have in °C so 9 ºF ( 196 ºC)  32 ºF TF  = –321 ºF 5 ºC  The melting point of UF6 is 64.53 °C. What is the melting point of uranium UF6 on the Fahrenheit scale? 9ºF TF = 5ºC xTC + 32 A. 67.85 °F B. 96.53 °F C. 116.2 °F D. 337.5 °F E. 148.15 °F Section 1.9 Density  Mass of substance per unit volume of the substance.  Common units are g/cm3 or g/mL. mass Density = volume 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 Density = 7.57 g/cm3 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  A student weighs a piece of gold that has a volume of 11.02 cm3 of gold. She finds the mass to be 212 g. What is the density of gold? d m V 212 g d  19.2 g/cm3 11.02 cm3 Another student has a piece of gold with a volume of 1.00 cm3. What does it weigh? 19.2 g What if it were 2.00 cm3 in volume? 38.4 g 30 (Q) If the density of an object is 2.87 x 10−4 lbs/cubic inch, what is its density in g/mL? (1 lb = 454 g, 1 inch = 2.54 cm) Section 1.10 Classification of Matter Matter  Anything occupying space and having mass.  Matter exists in three states.  Solid  Liquid  Gas Section 1.10 Classification of Matter The Three States of Water Section 1.10 Classification of Matter Solid  Rigid  Has fixed volume and shape.  relatively incompressible Section 1.10 Classification of Matter Liquid  Has definite volume but no specific shape.  Assumes shape of container.  relatively incompressible Section 1.10 Classification of Matter Gas  Has no fixed volume or shape.  Takes on the shape and volume of its container.  easily compressible Section 1.10 Classification of Matter  A substance: is a kind of matter that cannot be separated into other kinds of matter by any physical process.  A mixture: is a kind of matter that can be separated by physical means into two or more substances.  An Element: A substance that cannot be decomposed by any chemical reaction into simpler substances (Fe, Au, Na etc…)  A Compound: is a substance composed of two or more elements chemically combined. H2O, NaCl, CO2 Section 1.10 Classification of Matter Mixtures  Have variable composition. Homogeneous Mixture  Having visibly indistinguishable parts; solution.  is a mixture that is uniform in its properties throughout given samples. Examples: NaCl solution, Soft drink, Air, Solder Heterogeneous Mixture  Having visibly distinguishable parts.  a mixture that consists of physically distinct parts, each with different properties Example: Sand and iron filings 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 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.  Distillation  Filtration  Chromatography 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) Section 1.10 Classification of Matter  A physical change is a change in the form of matter but not in its chemical identity. Examples: Ice melting, salt or sugar dissolved in water. (Physical property: Melting, boiling, electrical conductivity)  A chemical change, or chemical reaction, is a change in which one or more kinds of matter are transformed into a new kind of matter or several new kinds of matter. Examples: rust formation, burning butane gas in oxygen (Chemical property: Describes how a substance undergoes a chemical reaction ) 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 popsicle on a warm summer day Section 1.10 Classification of Matter Chemical Physical Magnesium burns when heated  Magnesium metal tarnishes in air  Magnesium metal melts at 922 K  Orange juice lightens when water  is added Section 1.10 Classification of Matter The Organization of Matter