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MOUNT VERNON HIGH SCHOOL IB CHEMISTRY 1: 2011-2012 IB CHEMISTRY 1 WORKBOOK 1 2 Course Content IB Chemistry 1 Mrs. McManus and Mrs. Townley Measurement and Data Processing Atomic Theory (SL) Periodicity (SL) Bonding (SL) Stoichiometry (SL) Energetics (SL) Kinetics (SL) and Equilibrium (SL) Acids and Bases (SL) Oxidation and Reduction (SL) Topic 11 Topic 2 Topic 3 Topic 4 Topic 1 Topic 5 Topics 6 and 7 Topic 8 Topics 9 The topics above will be covered during the course of the year. The numbered topics are taken directly from the IB Program syllabus. This workbook contains objectives, worksheets, labs and study guides for the topics shown. 3 FCPS Student Laboratory Safety Rules You and your classmates will be participating in many hands-on laboratory activities this year. Some of these activities may require the use of materials or pieces of equipment that are potentially harmful if not handled in a safe manner. Carefully review the following rules for student conduct. After you have reviewed these rules with your instructor, sign the statement indicating you understand and agree to follow these rules. Review these rules with your parent(s)/guardian(s) and have them sign the form, indicating that they understand the risks and rules you will follow and support your adherence to these safety rules. Failure to abide by these rules will result in removal from the lab and forfeiture of the grade. 1. No unauthorized experiments are permitted. Do not modify or change the design of your experiment without instructor’s written approval. 2. Avoid loose fitting clothing, open-toed shoes or sandals, dangling jewelry, and tie back long hair. Ponytails that could fall over the shoulders must be folded over and tied. Loose clothing or hair can catch on fire or knock over equipment causing accidents. 3. Wear appropriate eye protection at all times (safety goggles). Chemicals and sharp objects can damage unprotected eyes. If you splash chemicals in your eyes, notify your instructor immediately and begin flushing eyes for 15-30 minutes in the eye wash. 4. Absolutely no horseplay of any kind is permitted in lab. This can cause accidents. 5. No visiting by friends is allowed during lab sessions. Keep your attention on the laboratory at hand – more attention equals fewer accidents. 6. Eating, drinking, applying cosmetics, chewing on gum, pencils or fingernails is never permitted in the labs. Assume everything you handle in lab is contaminated. Eating or drinking in the lab could allow you to ingest poisons. Wash your hands before leaving the lab. This prevents your taking trace chemicals to other locations. 7. Do not deliberately smell or taste chemicals unless instructed to do so. When testing odors, use a wafting motion of your hand to direct the odors to your nose. 8. Fire is to be used with caution in the laboratory. Some organic solvents are highly flammable. No flammable solvents should be around an open flame. Follow your teacher’s instructions when heating liquids so as to avoid fires. 9. Burners and hot plates require special care in usage. Never leave a heat source unattended when in use. Remember that these as well as heated apparatus remain hot long after heating, so handle with care. Be aware of cords and heating surfaces. If a fire alarm occurs, turn off Bunsen burners and hot plates. 10. Never look directly into a test tube or flask. Do not point the opening of these at anyone especially when heating. Liquids can quickly become boiling hot which can cause their contents to eject forcefully from a tube. Do not use fingers as rubber stoppers. 11. Read bottle labels carefully and thoroughly. Always verify that you have selected the correct chemical. Report unlabelled containers to the instructor. Mixing the incorrect chemicals can have harmful results. 12. Never contaminate reagents by pouring unused portions into stock bottles. 4 13. Follow waste disposal procedures specified by your teacher. Do not dump chemicals into trashcans or put into the sink unless instructed to do so. Never put any solid material into the sink!! 14. Clean up broken glass immediately. 15. If you spill anything during lab, notify your instructor immediately (acids and bases must be neutralized before cleanup). 16. Wash your skin if contact is made with any chemical immediately and notify your instructor. Chemical burns must be washed for 15-30 minutes. 17. Immediately report any unusual odors, broken equipment or unsafe situations to the instructor. 18. Immediately report any spill, cut, burn or other injury to the instructor. 19. No chemicals or supplies may be taken from storerooms or labs for home use unless approved by teacher for independent (science fair) projects. 20. Know the location and use of the following laboratory safety equipment: eye wash station, fire extinguisher, safety shower, fire blanket, fume hood. 21. Know the basic procedures to be followed in the event of a fire at your lab table, fire on a person, cut or chemical burn, or chemical in the eyes. 5 Internal Assessment Form Student:____________________________ Teacher:_______________________ Date:_________ Lab Title:___________________________ Topic/Option:_____________________ Hours:________ Students receive 2 marks for completely fulfilling an aspect, 1 mark for partially fulfilling it and 0 marks for failing to fulfill the aspect. Criterion Design Data Collection/Processing Conclusion and Evaluation Aspect 1 Aspect 2 Aspect 3 Formulates a focused problem/research question and identifies relevant variables. Designs a method for effective control of the variables. Develops a method that allows for the collection of sufficient relevant data. 2 1 0 Records appropriate quantitative and associated qualitative raw data, including units and uncertainties where relevant. 2 1 0 Processes the quantitative raw data correctly. 2 1 0 Presents processes data appropriately and, where relevant, includes errors and uncertainties. 2 1 0 States a conclusion, with justification, based on a reasonable interpretation of the data. 2 1 0 Evaluates weaknesses and limitations. 2 1 0 Suggests realistic improvements in respect of identified weaknesses and limitations. 2 1 0 2 6 1 0 2 1 0 Total Marks 7 8 Topic 11: Measurement and Data Processing 11.1 Uncertainty and Error in Measurement 11.1.1 Describe and give examples of random uncertainties and systematic errors. 11.1.2 Distinguish between precision and accuracy. 11.1.3 Describe how the effects of random uncertainties may be reduced. 11.1.4 State random uncertainty as an uncertainly range(±). 11.1.5 State the results of calculations to the appropriate number of significant figures. 11.2 Uncertainties in Calculated Results. 11.2.1 State uncertainties as absolute and percentage uncertainties. 11.2.2 Determine the uncertainties in results. 11.3 Graphical Techniques 11.3.1 Sketch graphs to represent dependences and interpret graph behavior. 11.3.2 Construct graphs from experimental data. 11.3.3 Draw best-fit lines through data points on a graph. 11.3.4 Determine the values of physical quantities from graphs. 9 Topic 11 Notes 10 Measurement & Calculations Metric System We use several types of measurements: length - meters (km, cm and mm) mass - grams (kg and mg) amount of substance – moles (mol) time – seconds (s) volume - Liters (mL) or cubic centimeters (cm3) We commonly use the prefixes: deci – 1 / 10 th centi - 1 / 100 th milli - 1 / 1000 th kilo - 1000 Occasionally you will encounter micro (µ), nano, pico, mega, and giga. Significant Digits: What do they mean? In a measurement or a calculation, it is important to know which digits of the reported number are significant. That means…if the same measurement were repeated again and again, some numbers would be consistent and some might vary. All of the digits that you are absolutely certain of plus one more that is a judgment are significant. Consider: 16.82394 cm. If all the digits are significant, everyone who measures the object will agree that it is 16.8239 cm, but everyone will not agree about the final digit; some would say …94 cm while others might say …95 cm. Rules for Recognizing Significant Digits All non-zero digits are significant. 523 g (3); 972,366 sec (6) 0’s in the MIDDLE between two non-zero digits are ALWAYS significant. 5082 m (4); 0.002008 L (4) 0’s in the FRONT of a number are NEVER significant. 0.0032 kg (2); 0.00000751 m (3) 0’s at the END of a number with a decimal point present are ALWAYS significant. 2.000 L (4); 0.00500 g (3) Significant Digits in Calculations When you perform a calculation using measurements, often the calculator gives you an incorrect number of significant digits. Here are the rules to follow to report your answers: x and ÷ : The answer has the same # of sig. digits as the measurement in the problem with the least number of sig. digits. Example: 3.71 cm x 8.1 cm = 30.051? 30. cm2 (2 sig. digits) + and – : The answer has the same # of decimal places as the measurement in the problem with the least number of decimal places. example: 3.70 cm + 8.1 cm = 11.8 cm (1 decimal place) Scientific Notation Scientific notation uses a number between 1 and 9.99… x 10 n where n is an integer. Know how to put numbers into scientific notation: 5392 = 5.392 x 103 0.000328 = 3.28 x 10-4 1.03 = 1.03 5500 = 5.5 x 103 Some 0’s in numbers are placeholders and are not a significant part of the measurement so they disappear when written in scientific notation. Ex: 0.000328 above. In scientific notation, only the three sig. digits (3.28) are written. Accuracy vs. Precision Accuracy refers to how close a measurement is to some accepted or true value (a standard). Ex: an experimental value of the density of Al is 2.69 g/mL. The accepted value is 2.70 g/mL. Your value is accurate to within 0.4%. Percent error is used to express accuracy. Precision refers to the reliability, repeatability, or consistency of a measurement. ± uncertainty and sig. digits are used to express precision. 11 Worksheet Significant Figures State the number of significant figures in each of the following measurements: 1. 8.523 cm 5. 0.00560 m 5 2. 9.50 x 10 s 6. 27 students 3. 0.0040040 kg 7. 950 g 4. 8.0000 m 8. 0.0000001 cm For each of the following measurements (a) state whether or not the zero(s) are significant and (b) state the rule that applies: 9. 9502 m 10. 52.0 cm 11. 0.0045 m 12. 370 s Calculate the answer for the following problems using the proper number of significant figures and appropriate units. 13. 15.2 cm x 0.0021 cm 14. 25. 6 m + 2365.2 m + 152.123 m 15. 28.7 cm2 / 4.00 cm 16. (3.52 x 105 m) x (1.522 x 107 m) 17. (9.0 x 1025 m) x (6.00 x 1010 m) 18. 10.000 cm – 8.99 cm 19. 0.00043256 m x 1.0 m 20. (1.0 x 103 m) x (6 x 10-8 m) Write the SI prefix that corresponds to the following: 21. 10 times smaller (101) 22. 1000 times smaller (103) 23. 1000 times larger (103) 24. 100 times smaller (102) 25. 1 000 000 times smaller (106) Read the following measurement devices to the proper number of significant digits: 26. cm 2 3 4 20 30 40 5 6 27. cm 50 60 28. cm 4 29. 5 6 5 4 mL 12 Calculations Using Significant Figures When multiplying and dividing, limit and round your answer to the least number of significant figures in any of the factors. You are only as good as your least precise measurement. Ex. 23.0 cm 432 cm 19 cm = 188.784 cm3 The answer is expressed as 1.9 102 cm3 since 19 cm has only two significant figures. When adding and subtracting, limit and round your answer to the least amount of decimal places in any of the measurements. Ex. 123.25 mL + 46.0 mL + 86.257 mL = 255.507 mL The answer is expressed as 255.5 mL since 46.0 mL has only one decimal place. Perform the following operations expressing the answer in the correct number of significant figures. 1. 1.35 m 2.467 m = ____________ _____________ (Calculator answer) (Answer w/ sig. Figs) 2. 1,305 m2 42 m = ____________ _____________ 3. 12.01 mL + 35.2 mL + 6 mL = ____________ _____________ 4. 55.46 g – 28.9 g = ____________ _____________ 5. 0.21 cm 3.2 cm 100.1 cm = ____________ _____________ 6. 0.15 cm + 1.15 cm + 1.051 cm = ____________ _____________ 3 7. 150 L 4 L = ____________ _____________ 3 2 8. 1.278 10 m 1.4267 10 m = ____________ _____________ Percentage Error Percentage error is a way to express how far off an experimentally determined value is from the accepted or true value. % error = (accepted – experimental) accepted value 100 Determine the percentage error in the following problems. Remember the rules for significant figures. 1. Experimental value = 1.24 g, Accepted value = 1.30 g 2. Experimental value = 0.124 g, Accepted value = 0.0998 g 3. Experimental value = 252 mL, Accepted value = 225 mL 4. Experimental value = 22.2 L, Accepted value = 22.4 L 5. Experimental value = 125.2 mg, Accepted value = 124.8 mg 13 Dimensional Analysis We can use equalities to change from one unit to another. For example: 1 minute = 60 seconds Therefore: 1 min 60 sec =1= 60 sec 1 min To convert one unit to another use the following steps: 1. Write the given number and unit. 2. Set up a conversion factor (like the fraction shown above) comparing the given unit to another. a. Place the given unit in the denominator of the conversion factor b. Place desired unit as numerator c. Place a “1” in front of the larger unit d. Determine the number of small units in “1” larger unit 3. Cancel units. Solve the problem Example: 25 km = _____ cm 25 km 1000 m 1 km 100 cm 1m = 2.5 x 106 cm Problems. Convert the following measurements. Show all work and conversion factors for full credit. 1. 2.5 hours = ______ sec 2. 15 cm = _____ m 3. 4.3 x 105 mm = _____ m 4. 8.00 x 10-3 km = _____ dm 5. 0.0075 m = _____ cm 6. 3.5 days = _____ sec 7. 2.5 x 103 mL = _____ L 8. 8,943 mg = _____ kg 14 Experimental Design When designing an experiment “from scratch”, there are several things to consider. The experiment should be designed so that there is a hypothesis, independent and dependent variables, constants, controls, measurable data, repeated trials and graphs. What are these components? HYPOTHESIS- an educated guess as to the relationship between the variables that can be tested. It can be put into the form of: “If I change (the independent variable), then (the dependent variable) will do something. INDEPENDENT VARIABLE- (manipulated variable) - the variable that is purposely changed by the experimenter. DEPENDENT VARIABLE- (responding variable) - the variable that responds to the manipulated changes. It is usually measured with items such as stopwatches, measuring devices, balances, etc. Try to eliminate subjectivity. Do a ranking of change if there is no other way to measure. CONSTANTS- all factors that remain the same and have a fixed value throughout the repeated trial of the experiment. Constantly check to make sure they do stay the same. CONTROL- the standard for comparing experimental effects. With living organisms, it will generally be the ones grown at the optimal conditions for that organism. REPEATED TRIALS- the number of experimental repetitions, object or organisms tested at each level of the independent variable. They help to reduce the effects of errors. How many should you have? It depends on the nature of the experiment- living things require more trials since they have more variation. GRAPH- the independent variable is always on the X axis while the dependent variable is on the Y axis. Bar graphs are used when the data is qualitative (descriptive, based on observations or categories of data). Line graphs are used when the data is quantitative (more precise, measured with tools). **VERY IMPORTANT** When designing an experiment, you should have only one independent and one dependent variable. You can have several experiments with different independent variables, but they must NOT be done at the same time! Answer the following questions about experiment design. A chemical experiment is performed by mixing substance A with substance B. A color change indicates the completion of the reaction. The time it takes for the color change to occur is recorded below. Each experiment is done in the same beaker. Experiment # 1 2 3 Volume of A (mL) 5.0 5.0 5.0 Volume of B (mL) 5.0 10.0 15.0 Temperature (oC) 25 26 24 Time of reaction (sec) 45 61 78 1. Identify the independent variable. 2. Identify the dependent variable. 3. Identify the constants. 4. Name two things that can be done to improve this experiment. 5. If a graph were to be made, what measurement (remember units) would be placed on the x-axis? 6. If a graph were to be made, what measurement (remember units) would be placed on the y-axis? 15 Experimental Design (cont.) When designing your own experiment you must choose appropriate variables and constants, create appropriate tables to collect measurable data and present the findings with charts and graphs. Let’s design an experiment! (You won’t actually have to perform the experiment, just design it!) We want to study plant growth (something you know about from biology). 1. Make a list of at least 5 variables affecting plant growth. a. b. c. d. e. 2. From the list you created, choose an independent variable, the variable you will manipulate in your experiment. ____________________________ 3. Choose a dependent variable, the variable you will measure the changes in based on your manipulations of the independent variable. ______________________ 4. List your constants in the experiment. 5. Write an appropriate title for your experiment: 6. Create a data table for your experiment (you don’t have actual numbers but you can make a table labeled with appropriate units.) 7. To make a graph of your data, label the x and y axes below with the appropriate units. 16 Uncertainty in Measurement Look at the following measurement device. 5 10 cm 1. What place value does each mark on the device represent? (tens, ones, tenths etc.) 2. Read the device to the proper place value by guessing one place smaller than the smallest place value on the device. 3. In what place value does uncertainty exist? (tens, ones, tenths, etc.) 4. Record the measure using the proper uncertainty. (± notation in the uncertain place value) Density Calculations using uncertainty Density is calculated with two measurements, mass ÷ volume. Since both measurements have uncertainty, any calculated density will have uncertainty as well. There are two rules to keep in mind When adding or subtracting measurements, add the uncertainties. When multiplying or dividing measurements, add the percent uncertainties. Example1 (add/sub): 12.34 ± .01 g - 10.11 ± .01 g ────────── 2.23 ± .02 g Example 2 (mult/div): 1.23 ± .01 g ÷ 2.0 ± .1 ml First determine the % uncertainty: .01 / 1.23 × 100 = 0.81 % .1 / 2.0 × 100 = 5 % Now divide the values and add the uncertainties: 1.23 g ÷ 2.0 ml = 0.61 g/ml 0.81 % + 5 % = 5.81% Write the final answer by converting the % back to a numerical value: 5.81% × 0.61 = 0.04 Final answer: 0.61 ± 0.04 g/ml 17 Uncertainty Worksheet Addition and Subtraction 1. 5.67 ± .01 g + 8.28 ± .01 g 2. 89.6 ± .2 cm - 25.7 ± .2 cm 3. 28 ± 1 g + 165 ± 2 g Multiplication and Division Remember, you must first find the percent uncertainties! 4. 8.75 ± .02 m × 5.67 ± .02 m 5. 21.5 ± .2 g ÷ 6. 125 ± 2 g × 3.7 ± .2 m 5.78 ± .02 m Putting it all together! Given the following measurements, calculate the density of the unknown liquid showing the uncertainty. Mass of empty graduated cylinder: 25.00 ± .01 g Mass of cylinder plus liquid: 28.54 ± .01 g Volume of liquid in cylinder: 2.0 ± .1 ml (Show each step of the calculations) Mass of liquid: Density of liquid: 18 Lab- Density Determination Purpose: To practice measurement and calculation techniques with uncertainty. Procedure Part 1 1. Determine the mass of the regularly shaped object. 2. Measure the length, width, and height of the object. Remember the precision of the measuring device you use! Note the smallest marking on the measurement device and record your measurements with the proper uncertainty. 3. Using the same object, determine the volume by water displacement. Again, note the smallest marking on the measurement device and record your measurement with the proper uncertainty. Part 2 You will determine the density of water by devising a procedure to determine the relationship that exists between volume and mass of water. Collect at least 5 data points and graphically represent your findings. Be sure to write your procedure and data in your lab notebook. Identify your independent and dependent variable. Data Record all data in a neatly labeled format Calculations Show all work with units and use uncertainty notation. 1. Calculate the volume of the object using its length, width and height. 2. Calculate the density of the object using the volume found in #1. 3. Calculate the volume of the object using the water displacement data. 4. Calculate the density of the object using the volume found in #3. This answer may be slightly different than the answer in #2. 5. Make a graph of the mass and volume data collected in part 2 of the procedure. Title the graph and label the axes with units. Determine the slope of the line (remember units!) Questions Answer in complete sentences 1. Compare the two methods you used in part 1 to determine volume. Which measurement was less uncertainty? Justify your answer. 2. Define the terms intensive and extensive property. Is density an intensive or extensive property? 3. The accepted value for the density of water is 1.00 g/mL. What is your percent error for the density of water? Concluding Statements- answer the following questions 1. What was determined about the two density values found in Part 1? Explain your findings. 2. What type of relationship is shown on the density graph? 3. Why do scientists use uncertainties and significant digits? 19 Topic 11 Notes 20 Topic 2: Atomic Structure 2.1 The Atom 2.1.1 State the position of protons, neutrons and electrons in the atom. 2.1.2 State the relative masses and relative charges of protons, neutrons and electrons. 2.1.3 Define the terms mass number (A), atomic number (Z) and isotopes of an element. 2.1.4 Deduce the symbol for an isotope given its mass number and atomic number. 2.1.5 Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number, atomic number and charge. 2.1.6 Compare the properties of the isotopes of an element. 2.1.7 Discuss the uses of radioisotopes. 2.2 The Mass Spectrometer 2.2.1 Describe and explain the operation of a mass spectrometer. 2.2.2 Describe how the mass spectrometer may be used to determine relative atomic mass using the 12C scale. 2.2.3 Calculate non-integer relative atomic masses and abundance of isotopes from given data. 2.3 Electron Movement 2.3.1 Describe the electromagnetic spectrum. 2.3.2 Distinguish between a continuous spectrum and a line spectrum. 2.3.3 Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels. 2.3.4 Deduce the electron arrangement for atoms and ions up to Z=20. 12.1.3 Electron Configuration 12.1.3 State the relative energies of s, p, d, and f orbitals in a single energy level. 12.1.4 State the maximum number of orbitals in a given energy level. 12.1.5 Draw the shape of an s orbital and the shapes of the px, py, and pz orbitals. 12.1.6 Apply the Aufbau principle, Hund’s rule and the Pauli Exclusion principle to write electron configurations for atoms up to Z = 54. 21 Topic 2 Notes 22 Subatomic Particles Atoms are composed of even smaller particles known as subatomic particles. Fill in the following table summarizing the properties of these particles. Particle Mass (amu and g) Charge Location Proton Neutron Electron Define the following terms: Atomic number (Z): Mass number (A): Isotope: Ion: Atomic mass: Isotopic Notation (Shorthand!) In the boxes below, identify what each letter represents. A X Z n+/- How many electrons in 23Na+? How many neutrons in 23Na+? Write the shorthand notation for the negative 2 charged ion of sulfur with a mass number of 33. 23 Isotopes Worksheet 1. Fill in the table for each isotope listed Symbol 20 Z A Number of protons Number of electrons Number of Neutrons Number of electrons Number of Neutrons Ne 201 66 Hg Zn 27 Al 2. Complete the following table for the following atoms or ions. Symbol Z A 36 84 Number of protons 36 35 53 127 35 54 27 27 Zn 32 36 Cd2+ 112 38 X2X3+ X3- 45 103 36 50 54 75 42 33 42 24 Isotopes and Average Atomic Mass 1. Elements contain a variety of isotopes. The atomic mass is the weighted average of all the isotopes of an element. Example: A sample of cesium is 75% 133Cs, 20% 132Cs, and 5% 134Cs. What is the average atomic mass? 0.75 133 = 99.75 0.20 132 = 26.4 0.05 134 = 6.7 total = 132.85 amu = average atomic mass Determine the average atomic mass of the following mixtures of isotopes. a. 80% 127I, 17% 126I, 3% 128I b. 50% 197Au, 50% 198Au c. 15% 55Fe, 85% 56Fe d. 99% 1H, 0.8% 2H, 0.2% 3H 2. The atomic mass of boron is given in the periodic table as 10.81, yet no single atom of boron has a mass of 10.81 amu. Explain. 3. The element Europium exists in nature as two isotopes: Eu-151 and Eu-153. Their masses are 150.9196 amu and 152.9209 respectively. The average atomic mass of europium is 151.96 amu. Calculate the relative abundance of the two europium isotopes. 25 Mass Spectrometry Mass spectrometry is used to determine the relative abundance of isotopes of elements. It works by vaporizing a substance then ionizing it (turning it into an ion by hitting it with electrons) so that it can be accelerated through a tube and deflected by a magnet. The magnetic field required to deflect it gives a measure of the mass to charge ratio. The more massive the ion, the stronger the magnetic field required to deflect it so that it hits the detector. The more of a particular kind of ion, the bigger the peak on the spectra. 2. Substance is IONIZED 1. Substance is VAPORIZED 4. Ions are DEFLECTED Spectrometer is under vacuum to keep as few molecules in the machine as possible to reduce the number of particles that collide with each other in the machine. 3. Ions are ACCELERATED 5. Ions are DETECTED. Sample spectra of the isotopes of a made-up element, E. 100 Relative number of Ions 50 0 24 25 m/z (m = mass, z = charge) Clearly the most abundant isotope is E-24. The other isotope, E-25, is only about 10% as common as E-24. You can figure out the average atomic mass by: 1. Abundance of E-24: 100 Abundance of E-25: 10 2. Total amount: 100 + 10 = 110 3. % abundance E-24 = (100/110)*100 = 90.91 % 4. % abundance E-25 = (10/110)*100 = 9.09 % 5. Average atomic mass = (24 * 0.9091) + (25* 0.0909) = 24.1 (It’s important to know the percentages to several significant figures because the whole point of this is to get the average atomic mass of an element to several significant figures.) Mass spectrometry is also useful for determining the structure of organic compounds. See organic notes for info on that. 26 Mass Spectrometer The existence of isotopes can be proven with the use of a mass spectrometer. This device has 5 stages of operation: vaporization, ionization, acceleration, deflection and detection. (VIADD) Draw a simple diagram of a mass spectrometer showing where all five stages take place. Describe what happens in each of the five stages. Vaporization— Ionization— Acceleration— Deflection— Detection— What factors determine the how much a particle is deflected? The following results were obtained for a pure element (X) placed in a mass spec. 6 Relative Abundance 2 85 86 87 88 (mass/charge) Calculate the relative (average) atomic mass of element X from the above spectrum. 27 Bohr and the Atom 1. According to Bohr, electrons orbit the nucleus in what he called ________________. 2. The lowest energy level, represented by the letter n, is n = ______ . 3. Elements currently discovered have electrons occupying up to how many energy levels? ______ 4. An electron going from n =1 to n = 2 must absorb/release energy. (circle one) 5. An electron going from n = 5 to n = 4 must absorb/release energy. 6. Which is more stable, an electron in n = 4 or in n = 3? 7. Define electromagnetic radiation. 8. Define quantum. 9. Name six types of electromagnetic radiation. 10. What properties are different between infrared and ultraviolet radiation? What property is the same? 11. Wavelength and frequency are directly/inversely related. 12. Frequency and energy are directly/inversely related. 13. Write the equations showing (1) the relationship between wavelength and frequency, (2) and the relationship between frequency and energy. 14. A wave with a frequency of 1.0 x 1020 Hz is what type of radiation? 15. A wave with a frequency of 1.0 x 109 Hz is what type of radiation? 16. What is the frequency of a light with a wavelength of 555 nm? 17. What is the energy of light with a frequency of 1.5 x 1014 Hz? 18. What is the energy of a quantum of energy with a wavelength of 175 nm? 28 Electron Configurations (Schrodinger) 1. Energy levels are further classified as having sublevels. How many sublevels are contained in the following energy levels? Energy Level #of sublevels 2 4 n 2. List four types of sublevels in order of increasing energy. 3. The actual region in space having the highest probability of containing the electron is known as an ___________. 4. How many orbitals are found in each of the following types of sublevels? sublevel # of orbitals s p d f 5. What is the maximum number of electrons that can occupy any orbital? 6. State the Pauli exclusion principle. 7. State Hund's rule. 8. How many electrons can occupy each of the following sublevels? sublevel # of electrons s p d f 9. Fill in the orbital filling diagrams below for the electrons in N, Mg, S, and Fe. 10. Write the electron configurations for N, Mg, S and Fe based on your diagrams above. 29 Electrons, Configurations and Orbital Diagrams Fill in the following chart: Element # of Valence Electrons Si 4 Electron configuration Core Outer electrons config Orbital filling diagram for outer electrons 3s 3p 3s23p2 [Ne] Fe S2- Cr K+ Fill in the corresponding outer electron configurations as shown. 1s1 1s2 2p3 3d3 * # 4f7 * # 30 . 2.8.4 N 6s1 Electron Arrangement Spectroscopy: Element Identification and Emission Spectra The energy levels in atoms and ions are the key to the production and detection of light. Energy levels or "shells" exist for electrons in atoms and molecules. The colors of dyes and other compounds results from electron jumps between these shells or levels. The colors of fireworks result from jumps of electrons from one shell to another. An observation of light emitted by the elements is also evidence for the existence of shells, sublevels and energy levels. The kinds of light that interact with atoms indicate the energy differences between shells and energy levels in the quantum theory model of the atom. Typically the valence electrons are the ones involved in these jumps. The "quantum" theory was proposed more than 90 years ago, and has been confirmed by thousands of experiments. Science and education has failed to clearly describe the energy level concept to almost four generations of citizens. This experiment is an exercise aimed at throwing a little more light on the subject. (Don't laugh too hard at the joke.) Atoms have two kinds of states; a ground state and an excited state. The ground state is the state in which the electrons in the atom are in their lowest energy levels possible (atoms naturally are in the ground state). This means the electrons have the lowest possible values for "n" the principal quantum number. Specific quantized amounts of energy are needed to excite an electron in an atom and produce an excited state. An excited hydrogen atom with an electron in the n = 3 shell can release energy. If the electron in hydrogen only drops to the n = 2 shell the energy matches a pulse of red light. Energy can be added to atoms many different ways. It can be in the form of light, an electric discharge or heat. This added or extra energy is emitted when the excited electrons in the atoms give off light and fall back to lower shells. The light emitted has wavelengths and colors that depend on the amount of energy originally absorbed by the atoms. Usually each individual excited atom will emit one type of light. Since we have billions and billions of atoms we get billions of excitations and emissions. Not all atoms in a sample will absorb or be excited exactly the same. For example in hydrogen the ground state has the electron in the n= 1 shell or level. The electron in some hydrogen atoms may be excited into the n = 2 level. Other hydrogen atoms can have the electron excited into the n = 4 level. Different elements emit different emission spectra when they are excited because each type of element has a unique energy shell or energy level system. Each element has a different set of emission colors because they have different energy level spacings. We will see the emission spectra or pattern of wavelengths (atomic spectra) emitted by several different elements in this lab. We will then identify an unknown element by comparing the color of the unknown with the flame color of our knowns. You need to know that white light is the combination of all colors of the spectrum. 31 Each color has a characteristic wavelength. The wavelength is the distance between the beginning and end of a complete cycle of the light wave. All colors of light travel at the same speed, 3.0 x 108 meters/ second. The animation shows how a prism separates photons of red light from photons of blue light. The photons of different colors fall in different positions on the color spectrum. The position is determined by the wavelength. Red light has longer wavelength and is lower in energy than blue light. The wavelength of red light corresponds to the range of 700 to 600 nanometers, (7000 Ångstrom or 0.0000007 meters). Blue light has shorter wavelength in the range of 400 nm (4000 Ångstrom or 0.00000004 meter, 1 Å = 1 x 10 -10 m = 0.0000000001 meter = 1 x 10-1 nanometer). Spectroscopy is the analysis of light spectra and the way in which light interacts with matter. When light is analyzed it is commonly separated into its component colors. The light source is directed on a slit and the "beam" of light is separated using a prism or grating. The reason that the images are lines is that the light from the lamp is focused on a narrow slit. The illustration shows the separation of a light beam into its component colors. This produces an image of the slit which has the shape of a line. The resulting beam of light can be broken into the color spectrum or into its components of the spectrum emitted by the atom. You can see the specific colors emitted by the light source. A white light source will give a spectrum like the one shown above. 32 Experimental Procedure Part 1 Flame tests and identification of an unknown metal. Observe and record the color of the flame for each metal ion. Remember the metal ions are paired with a nonmetal ion in an ionic formula unit. The electrical charges have to add to zero. The metal ions are converted to atoms in the flame and then excited by the heat from the Bunsen burner flame. The nonmetal ions, anions, do not get converted to atoms and do not and emit visible light like the metals do. Repeat the procedure for each known. Record the color observed for the unknown and use the color to identify the cation in the unknown. Metal ion Observed Flame color barium _________________________ calcium _________________________ sodium _________________________ copper (II) _________________________ potassium _________________________ lithium _________________________ strontium _________________________ Part 1 Flame tests for unknown elements Unknowns Flame color Identity of metal ion based on flame test Unknown 1 ____________ __________ Unknown 2 ____________ __________ Part 2 Observing line spectra with the spectroscope In the second part of the experiment you will observe the color of light emitted by excited gases of elements in sealed glass tubes called "spectrum" tubes. Direct current, DC, high voltage electrons are used to excite the atoms in the spectrum tube. High voltage means 1000 to 2000 volts. This is more than 10 times normal household voltage which is 120 volts AC. The excited atoms release the energy they gained. Some of this energy is in the form of heat and some is in the form of light. The billions of excited atoms release energy. Each excited atom releases a single pulse of light energy as it returns to the "ground" state or low energy state. There are so many pulses emitted the light appears to be continuous. The excited atoms do not all emit the same energy light because the amount of energy that excited them may differ, but there are limitations on the colors they do emit. The kind of light depends on the size of the gaps between the "shells" or energy levels in the atom. The electrons are changing "n" values in the atom. Remember "n" can have only positive whole number values like 1, 2, 3, ... up to infinity. 33 The kind of light energy that can be emitted by excited atoms is unique for an element. The pattern of "lines” or colors emitted can be used to identify an element. A powerful extension of this is the ability to measure amounts of an element by measuring the brightness of the emitted light. A spectroscope can separate the light produced by an emission tube. The color seen by the naked eye is a combination of a number of colors of light. These are separated by a prism or a diffraction grating which acts like a prism. The emission lines can be seen when you look through the spectroscope at the light source. You will be able to observe the "line" spectrum for the elements and record the spectral lines. Element Part 2 Emission line spectra for selected elements Color with naked eye Emission spectrum Hydrogen Mercury Neon Questions and observations How do these emission spectra compare in terms of colors and numbers of emission line positions? Are the spectra identical? What if anything is similar? What is different? FILL IN THE FOLLOWING TABLE WITH YOUR ANSWERS Element with greatest number of visible emission lines ________________ _ Color of light for the longest wavelength What produces the colors seen in the flame tests and the emission spectra? 34 Spectroscopy lab (cont.) Questions. 1. Explain what causes the color that you see in the flame test portion of the lab. Be very detailed in your description of the electron’s behavior. Write at least four COMPLETE sentences in proper English to describe each aspect of the process. 2. What is the significance of line spectrum? What did it prove the existence of? Why does it verify that? Write at least four COMPLETE sentences in proper English. 3. Discuss the color of light as it relates to wavelength and energy. Is red light more energetic? What is the relationship between wavelength and energy? Feel free to bring in mathematical relationships to clarify your answer. Write at least four COMPLETE sentences in proper English. 35 Emission Spectrum of Hydrogen http://www.avogadro.co.uk/light/bohr/spectra.htm The diagram above shows several energy transitions as the electron falls from one energy level to another. Each energy transition represents a very specific quantity of energy as mathematically defined by the wave equation and Planck’s constant. Let’s look specifically at the emission spectrum of hydrogen. On the spectrum below color the lines representing the VISIBLE line spectrum of hydrogen. This spectrum occurs only when electrons transition to n=2 as their final energy state. These transitions create what is known as the Balmer Series. Label the lines with initial and final energy levels. 656 nm n=3 to n=2 700 nm 1. 2. 3. 4. 350 nm Which end of the spectrum represents more energy, 700 nm or 350 nm? What color is at the 700 nm end of the spectrum? Which transition represents the GREATEST amount of energy, n=3 to n=2 or n=4 to n=2? Notice that the lines get closer and closer to one another as you approach the violet end of the visible spectrum. Why do they converge? 5. An entire new series of lines was discovered when the electrons transitioned to n=1 (Lyman Series in the UV). Would this represent smaller or larger energy differences? Would these lines converge at a maximum energy just like the visible light series? 36 Models of the Atom (SOL Content) Match the following scientists with their model/experiment: ____ 1. Dalton a. “Father of the Periodic Table” ____ 2. Thomson b. quantum theory of energy ____ 3. Heisenberg c. first to say all matter is made of atoms ____ 4. Democritus d. oil drop experiment proved charge of electron ____ 5. Bohr e. electrons act as both waves and particles ____ 6. Rutherford f. uncertainty principle ____ 7. deBroglie g. electrons occupy an energy level ____ 8. Millikan h. atom is mostly empty space, nucleus at the center ____ 9. Mendeleev i. plum pudding model ____ 10. Planck j. atom is solid sphere Answer each of the following questions. 11. Why did Dalton rely so heavily on the results of others? 12. What device did Thomson use in his experiments? What particle was found using this device? 13. Describe the Rutherford gold foil experiment. (Draw a diagram) What did it show that had not been known previously? 14. Explain how Bohr’s model of the atom corresponds with what you observed in the flame test/emission spectra lab. 37 Introduction to Nuclear Chemistry (SOL Content) 1. What can be said about the nucleus of a radioactive isotope? 2. Name two types particles a nucleus can emit in an attempt to become more stable. 3. Fill in the following table regarding types of radioactive emissions. Type of emission Alpha () Beta () Gamma () Mass Charge Composition Shielding required to stop particle 4. A beta particle is a high speed electron emitted from the nucleus. This is not to be confused with electrons in the electron cloud. (Electrons in the cloud are ignored in nuclear reactions.) If the nucleus consists of only protons and neutrons, how is it possible for an electron to be emitted from it? 5. When balancing nuclear equations, what two properties must be conserved? 6. In isotopic notation of atoms, i.e. 146C, the number in the upper left refers to the _____________, and the number in the lower left refers to the _______________. 7. Define half-life. 8. Complete the following table regarding half-lives of a radioactive isotope. Number of half-lives Amount remaining Amount decayed 0 100% or 1 0% or 0 1 50% or ½ 50% or ½ 2 25% or 75% or 3 4 5 38 Nuclear Equations and Half-Life (SOL Content) Nuclear equations are balanced in a different manner than ordinary chemical reactions. Atoms are NOT conserved! Instead the nuclear properties of mass and charge are conserved. Balance the following nuclear equations. Remember in isotopic notation the superscript number is the mass number (protons + neutrons), and the subscript, if shown, is the atomic number (number of protons). Every element has a unique and constant atomic number that can be found on the periodic table. H + 3H → _________ 1. 1 2. 238 3. 42 K → 4. 27 Al + 4He → 5. 9 U → 4 2 He 0 -1e + _________ + ________ 30 P + _________ Be + 1H → _______ + 6. _______ + 01n → 142 Ba + 4 2He 91 Kr + 3 01n Half-life is the time it takes for one half of a radioactive sample to decay. Answer the following questions pertaining to half-life. 1. The half-life of iodine-131 is 8 days. What mass of I-131 remains from a 4.00 g sample after 32 days? 2. A sample of radioactive isotope with an original mass of 10.0 grams is observed for 24 days. After that time, 1.25 grams of the isotope remains. What is its half-life? 3. After a period of four half-lives of a radioactive isotope has passed, what fraction of its original mass will remain? What fraction has decayed? 4. The half-life of K-42 is 12.4 hours. How much of a 240.0 g sample will remain after 62.0 hours? 5. A 50.0 g sample of N-16 decays to 12.5 g in 14.4 seconds. What is its half-life? 39 TOPIC 2- NOTES 40 Study Guide: Topic 2 Atomic Theory 1. Which statement is correct about the isotopes of an element? (2.1) A. They have the same mass number B. They have the same electron arrangement C. They have more protons than neutrons D. They have the same numbers of protons and neutrons (Total 1 mark) 2. How many electrons are there in one A. 10 B. 12 C. 14 D. 22 24 +2 12 Mg ion? (2.1) (Total 1 mark) 3. What is the symbol for a species that contains 15 protons, 16 neutrons and 18 electrons? (2.1) A. 31 16 B. 31 16 S 3 C. 31 15 P D. 31 15 P 3 S (Total 1 mark) 4. What is the difference between two neutral atoms represented by the symbols (2.1) A. B. C. D. 59 27 Co and 59 28 Ni? The number of neutrons only. The number of protons and electrons only. The number of protons and neutrons only. The number of protons, neutrons and electrons. (Total 1 mark) 5. What is the correct number of each particle in a fluoride ion, 19F–? (2.1) A. B. C. D. protons 9 9 9 9 neutrons 10 10 10 19 electrons 8 9 10 10 (Total 1 mark) 6. A certain sample of element Z contains 60% of 69Z and 40% of 71Z. What is the relative atomic mass of element Z in this sample? (2.2) A. 69.2 B. 69.8 C. 70.0 D. 70.2 (Total 1 mark) 41 7. Which ion would undergo the greatest deflection in a mass spectrometer? (2.2) 16 + A. O 16 2+ B. O 18 2+ C. O 16 18 + D. ( O O) (Total 1 mark) 8. Which statement is correct about a line emission spectrum? (2.3) A. Electrons absorb energy as they move from low to high energy levels. B. Electrons absorb energy as they move from high to low energy levels. C. Electrons release energy as they move from low to high energy levels. D. Electrons release energy as they move from high to low energy levels. (Total 1 mark) 9. Which statement is correct for the emission spectrum of the hydrogen atom? (2.3) A. The lines converge at lower energies. B. The lines are produced when electrons move from lower to higher energy levels. C. The lines in the visible region involve electron transitions into the energy level closest to the nucleus. D. The line corresponding to the greatest emission of energy is in the ultraviolet region. (Total 1 mark) 10. What is the electron arrangement of an Al3+ ion? (2.3) A. 2, 8 B. 2, 3 C. 2, 8, 3 D. 2, 8, 8 (Total 1 mark) 11. The electron arrangement of sodium is 2.8.1. How many occupied main electron energy levels are there in an atom of sodium? (2.3) A. 1 B. 3 C. 10 D. 11 (Total 1 mark) 12. The element vanadium has two isotopes, (2.1) (a) 50 23 V and 51 23V, and a relative atomic mass of 50.94. Define the term isotope. ………………………………………………………………………………………. ………………………………………………………………………………………. (1) (b) State the number of protons, electrons and neutrons in 50 23V. ………………………………………………………………………………………. ………………………………………………………………………………………. 42 (2) (c) State and explain which is the more abundant isotope. ………………………………………………………………………………………. ………………………………………………………………………………………. (1) (d) State the name and the mass number of the isotope relative to which all atomic masses are measured. ………………………………………………………………………………………. (1) (Total 5 marks) 13. (a) Define the term isotope. (2.2) ..................................................................................................................................... ..................................................................................................................................... ..................................................................................................................................... (2) (b) A sample of argon exists as a mixture of three isotopes. mass number 36, relative abundance 0.337% mass number 38, relative abundance 0.0630% mass number 40, relative abundance 99.6% Calculate the relative atomic mass of argon. ..................................................................................................................................... ..................................................................................................................................... ..................................................................................................................................... ..................................................................................................................................... (2) (c) State the number of electrons, protons and neutrons in the ion 56Fe3+. electrons: ............................. protons: ............................. neutrons: ........................... (2) (Total 6 marks) 14. Identify the numbers of protons, neutrons and electrons in the species 33S2–. (2.1) .............................................................................................................................................. .............................................................................................................................................. (Total 1 mark) 43 15. (a) Describe the following stages in the operation of the mass spectrometer. (2.2) (i) ionization (2) (ii) deflection (2) (iii) acceleration (1) (b) (i) State the meaning of the term isotopes of an element. (1) (ii) Calculate the percentage abundance of the two isotopes of rubidium 85Rb and 87Rb. (2) (iii) State two physical properties that would differ for each of the rubidium isotopes. (1) (iv) Determine the full electron configuration of an atom of Si, an Fe3+ ion and a P3– ion. (3) (Total 12 marks) 16. (a) Evidence for the existence of energy levels in atoms is provided by line spectra. State how a line spectrum differs from a continuous spectrum. (2.3) ..................................................................................................................................... ..................................................................................................................................... (1) (b) On the diagram below draw four lines in the visible line spectrum of hydrogen. (1) Low energy (c) High energy Explain how the formation of lines indicates the presence of energy levels. ..................................................................................................................................... ..................................................................................................................................... (1) (Total 3 marks) 17. State the electron arrangement, electron configuration and orbital filling diagrams for the following particles: aluminium3+, nitrogen and fluorine1-. (2.3) (Total 6 marks) 44 Topic 3: Periodicity 3.1 The Periodic Table 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number. 3.1.2 Distinguish between the terms group and period. 3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z=54. 3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table. 3.2 Physical Properties 3.2.1 Define the terms first ionization energy and electronegativity. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization, energies, electronegativity and melting points for the alkali metals (Li Cs) and the halogens (F I). 3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization, energies, electronegativity for elements across period 3. 3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table. 3.3 Chemical Properties 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3. 45 Topic 3 NOTES 46 Periodic Table Exercise Use the periodic table on the following page to complete the following activities. 1. Label the columns 1-18 across the top. 2. Label the rows 1-7 along the left edge. 3. Fill in the atomic numbers for the elements. 4. Fill in the symbols for the elements (be sure to write in small print). 5. Fill in the atomic masses of the elements. 6. Color H red (Color lightly enough to see the information underneath). 7. Color the remaining column one elements orange. 8. Color column two elements yellow. 9. Color columns 3-12 blue. 10. Color column 17 green. 11. Color column 18 purple. 12. Put an asterisk (*) in the top right corner of elements #80 and #35. 13. Put a dot () in the top left hand corner of elements #1, 2, 7-10, 17, 18, 36, 54, and 86. Answer the following questions about the periodic table. 1. What phase of matter are most elements at room temperature and pressure? 2. At room temperature and pressure, what phase of matter are the elements you marked with an asterisk (*)? 3. At room temperature and pressure, what phase of matter are the elements you marked with a dot? 4. If a symbol for an element has two letters in it, is the second letter upper or lower case? 5. Why do some elements have symbols that seem unrelated to their common names, for example sodium has the symbol Na? 6. Seven elements on the table occur in nature as diatomic molecules. Name those seven elements and write their symbols. 7. As new elements are created/discovered, how are they initially named and how many letters are in their symbols? 47 48 Periodicity What two physical properties determine the strength of the force of attraction between two oppositely charged particles? (1) (2) Atomic Radii Define atomic radius: In general, atomic radii (increase/decrease) as you move down the periodic table. Why? In general, atomic radii (increase/decrease) as you move across the periodic table. Why? Define ionic radius: When an atom gains electrons it forms a (cation/anion) with a (negative/positive) charge. The newly formed ion is (larger/smaller) that the atom from which it was formed. When an atom loses electrons it forms a (cation/anion) with a (negative/positive) charge. The newly formed ion is (larger/smaller) that the atom from which it was formed. For each list of particles, place them in order of increasing size. a) K, Na, Rb b) Fe2+, Fe, Fe3+ c) N-, N, N+ d) F, N, O e) Cl-, S2-, Ar, K+1 Electronegativity Define electronegativity: In general, electronegativity (increases/decreases) as you move down the periodic table. Why? In general, electronegativity (increases/decreases) as you move across the periodic table. Why? In a bond formed between Na and F, which atom more strongly attracts the electrons in the bond? Metals tend to have a (high/low) electronegativity and tend to (gain/lose) electrons when bonded. They tend to form (positive/negative) ions. Nonmetals tend to have a (high/low) electronegativity and tend to (gain/lose) electrons when bonded. They tend to form (positive/negative) ions. 49 Ionization Energy (Ionisation energy) Define first ionization energy: In general, ionization energy (increases/decreases) as you move down the periodic table. Why? In general, ionization energy (increases/decreases) as you move across the periodic table. Why? The more tightly an electron is held to the nucleus the (higher/lower) the ionization energy. From each pair listed below pick the species with the specified property: a) largest radius Rb or Cs b) largest electronegativity Rb or Cs c) largest ionization energy Rb or Cs d) largest radius Cl or I e) largest elctronegativity Cl or I f) largest ionization energy Cl or I g) largest radius Na or Mg h) largest electronegativity Na or Mg i) largest ionization energy Na or Mg j) largest radius K or K+ k) largest radius Br or Br- 50 Halogens 1. Define the term diatomic. 2. List the seven diatomic elements. 3. How many valence electrons does a halogen atom have? 4. What is the valence shell electron configuration for all halogen atoms? 5. What charge ion do halogens tend to make? 6. Write the appropriate symbols for the following species: a. chlorine b. chloride 7. Explain the difference between a halogen and a halide. 8. What is the periodic trend of the melting point of the halogens? 9. All the halogens have similar chemical properties. Why? 10. If aluminum forms a compound with chlorine with the formula of AlCl3, what would you expect to be the formula of the compound formed between aluminum and bromine? 11. The compound potassium fluoride has the formula KF. What would you expect to be the formula of potassium iodide? 51 Periodic Trends in the Alkali Metals and Halogens 1. What is the trend in melting point as you go down the group of alkali metals? 2. What is the trend in melting point as you go down the group of halogens? 3. pH is a measure of the acidity of a solution. A pH of less than 7 means the solution is acidic/basic. A pH greater than 7 means the solution is acidic/basic. Acidic solutions are associated with high concentrations of the hydrogen (H+)/hydroxide ion(OH-). Basic solutions are associated with high concentrations of the hydrogen (H+)/hydroxide ion(OH-). 4. Describe the reaction of alkali metals in water: a. List visual observations b. What gas is produced? c. Is heat produced or absorbed? d. Complete the general form of the equation of an alkali metal reacting with water. Alkali metal + water → _______ + _______ + _______ e. Is the resulting solution acidic or basic? Justify. f. What is the trend of reactivity as you go down the group of alkali metals? g. What charge ion do alkali metals make? 5. Answer the following questions about reactions with halogens: a. What charge ion would you expect the halogens to make? b. What is the general trend in reactivity as you go down the group? c. Is fluorine capable of replacing bromide? d. Is iodine capable of replacing chloride? e. Complete the following equation: chlorine + bromide → _____ + _____ f. Write the symbols for the equation in “e” above. 6. Answer the following questions about oxide compounds: a. What is the formula for sodium oxide? b. What is the formula for aluminum oxide? c. When a metallic oxide is place in water a basic/acidic solution is formed with a pH above/below 7. d. When a nonmetallic oxide is place in water a basic/acidic solution is formed with a pH above/below 7. e. Define amphoteric. f. An amphoteric substance will behave like a base/acid when in the presence of a strong acid. g. List an example of an amphoteric substance. h. What is the trend of acidic/basic nature of the period 3 oxides? 52 Study Guide: Topic 3 Periodicity 1. For which element are the group number and the period number the same? A. Li B. Be C. B D. Mg (Total 1 mark) 2. What is the total number of p orbitals containing one or more electrons in germanium (atomic number 32)? A. 2 B. 3 C. 5 D. 8 (Total 1 mark) 3. Which is correct about the element tin (Sn) (Z = 50)? Number of main energy levels Number of electrons in containing electrons highest main energy level A. 4 4 B. 4 14 C. 5 4 D. 5 14 (Total 1 mark) 4. Which equation represents the first ionization energy of fluorine? A. F(g) + e– F–(g) B. F–(g) F(g) + e– C. F+(g) F(g) + e– D. F(g) F+(g) + e– (Total 1 mark) 5. Which factors lead to an element having a low value of first ionization energy? I. large atomic radius II. high number of occupied energy levels III. high nuclear charge A. I and II only B. I and III only C. II and III only D. I, II and III 6. Which of the following properties of the halogens increase from F to I? I. Atomic radius II. Melting point III. Electronegativity A. I only B. I and II only C. I and III only D. I, II and III (Total 1 mark) (Total 1 mark) 53 7. Which statement is correct for a periodic trend? A. Ionization energy increases from Li to Cs. B. Melting point increases from Li to Cs. C. Ionization energy increases from F to I. D. Melting point increases from F to I. (Total 1 mark) 8. Which series is arranged in order of increasing radius? A. Ca2+ < Cl– < K+ B. K+ < Ca2+ < Cl– C. Ca2+ < K+ < Cl– D. Cl– < K+ < Ca2+ (Total 1 mark) 9. Which pair of elements reacts most readily? A. Li + Br2 B. Li + Cl2 C. K + Br2 D. K + Cl2 (Total 1 mark) 10. Rubidium is an element in the same group of the periodic table as lithium and sodium. It is likely to be a metal which has a A. B. C. D. high melting point and reacts slowly with water. high melting point and reacts vigorously with water. low melting point and reacts vigorously with water. low melting point and reacts slowly with water. (Total 1 mark) 11. Explain why (i) the first ionization energy of magnesium is lower than that of fluorine. (2) (ii) magnesium has a higher melting point than sodium. (3) (Total 5 marks) 54 12. (a) (i) State the meaning of the term electronegativity and explain why the noble gases are not assigned electronegativity values. (2) (ii) State and explain the trend in electronegativity across period 3 from Na to Cl. (2) (iii) Explain why Cl2 rather than Br2 would react more vigorously with a solution of I–. (2) (b) State the acid-base properties of the following period 3 oxides. MgO Al2O3 P4O6 Write equations to demonstrate the acid-base properties of each compound. (7) (Total 13 marks) 13. Describe the acid-base character of the oxides of the period 3 elements Na to Ar. For sodium oxide and sulfur trioxide, write balanced equations to illustrate their acid-base character. (Total 3 marks) 55 Topic 3 NOTES 56 Topic 4: Bonding 4.1 Ionic Bonding 4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged ions. 4.1.2 Describe how ions can be formed as a result of electron transfer. 4.1.3 Deduce which ions will be formed when elements in groups 1, 2, and 3 lose electrons. 4.1.4 Deduce which ions will be formed when elements in groups 5, 6, and 7 gain electrons. 4.1.5 State that transition elements can form more than one ion. 4.1.6 Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values. 4.1.7 State the formula of common polyatomic ions formed by nonmetals in periods 2 and 3. 4.1.8 Describe the lattice structure of ionic compounds. 4.2 Covalent Bonding 4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei. 4.2.2 Describe how the covalent bond is formed as a result of electron sharing. 4.2.3 Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom. 4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength. 4.2.5 Predict whether a compound of two elements would be covalent from the position of the elements in the periodic able or from their electronegativity values. 4.2.6 Predict the relative polarity of bonds from electronegativity values. 4.2.7 Predict the shape and bond angles for species with four, three and two negative charge centres on the central atom using the valence shell electron pair repulsion theory (VSEPR). 4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities. 4.2.9 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C60 fullerene). 4.2.10 Describe the structure of and bonding in silicon and silicon dioxide. 4.3 Intermolecular Forces 4.3.1 Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from the structural features of molecules. 4.3.2 Describe and explain how intermolecular forces affect the boiling points of substances. 4.4 Metallic Bonding 4.4.1 Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons. 4.4.2 Explain the electrical conductivity and malleability of metals. 4.5 Physical Properties 4.5.1 Compare and explain the properties of substances resulting from different types of bonding. 57 Topic 4 NOTES 58 Ionic Bonding Basics Define the following terms: Ion: Cation: Anion: Stable Octet: Polyatomic ion: 1. When an atom loses electrons, what type of ion is formed? 2. When an atom gains electrons, what type of ion is formed? 3. Why do atoms form ions? 4. The most common charges of ions can be determined by placement in the periodic table. Give the most common charge of an ion formed from the following families of atoms: a. alkali metals b. alkaline earth metals c. transition metals d. column 15 or 5A e. column 16 or 6A f. halogens g. noble gases 5. When metals form ions they tend to (lose/gain) electrons. 6. When nonmetals form ions they tend to ______ electrons. 7. When ions come together to form an ionic compound, what is the overall charge of the compound? Fill in the following chart for polyatomic ions Name Formula Charge ammonium carbonate hydrogen carbonate acetate nitrate nitrite phosphate sulfate hydrogen sulfate sulfite chlorate hydroxide 59 Naming and Writing Formulas for Ionic Compounds Write the names for the following: 1. BaCl2 11. K2O 2. Zn(NO3)2 12. Li2SO4 3. Ca(OH)2 13. Al2S3 4. Sr(ClO3)2 14. (NH4)3PO4 5. Mg3N2 15. SrF2 6. NaC2H3O2 16. NaHCO3 7. AgI 17. KNO2 8. Mg(HSO4)2 18. NH4C2H3O2 9. K2SO3 19. BaSO4 10. AlP 20. Na2O 21. calcium carbonate 31. sodium sulfate 22. sodium sulfide 32. aluminum nitrate 23. ammonium chloride 33. barium acetate 24. strontium bromide 34. magnesium phosphate 25. aluminum hydroxide 35. lithium sulfite 26. barium nitrate 36. potassium oxide 27. calcium acetate 37. ammonium sulfide 28. magnesium iodide 38. sodium nitride 29. potassium phosphide 39. calcium hydrogen sulfate 30. lithium carbonate 40. aluminum oxide Write the formulas for the following: 60 Stock System Most transition metals and the elements tin and lead have multiple oxidation states (they can have more than one ionic charge.) To differentiate between the different charges a Roman numeral is used to express the CHARGE of the ion. Name the following compounds. All require the use of the stock system. 1. CuCl 6. CuSO4 2. PbI2 7. Fe(NO3)3 3. SnCl4 8. Pb(OH)4 4. Fe2O3 9. SnSO4 5. CrF3 10. HgO Write the formulas for the following compounds. 11. chromium (III) sulfate 16. tin (IV) oxide 12. iron (III) bromide 17. copper (II) phosphate 13. lead (II) acetate 18. iron (II) sulfide 14. copper (I) oxide 19. tin (II) fluoride 15. mercury (II) hydroxide 20. lead (IV) carbonate Naming Simple Organic Compounds (Alkanes) Hydrocarbons with the formula of CnH2n+2 are known as alkanes. They have their own naming system using root words for the number of carbons in the compound. List the root words for the following number of carbon atoms in the compound: 13572468Name the following compounds: 1. C3H8 4. C8H18 2. CH4 5. C2H6 3. C5H12 6. C4H10 Write formulas for the following compounds: 7. hexane 9. ethane 8. methane 10. propane 61 Types of Chemical Bonds Classify the following compounds as containing either ionic bonds (metal cation bonded to a nonmetal anion), covalent bonds (nonmetal atom + nonmetal atom), or both (compound containing a polyatomic ion substituted for one of the ions in an ionic compound. 1. 2. 3. 4. 5. 11. CO2 Li2O H2O NaNO3 NO2 6. 7. 8. 9. 10. KOH N2O5 CaSO4 SrBr2 FeCl3 When two atoms have greatly different electronegativities the bond formed between them is said to be which type, ionic or covalent? Naming Molecular Compounds Molecular compounds contain covalently bonded nonmetal atoms. A prefix system is used to indicate the number of atoms of each element. The second element ends in the suffix “–ide.” Following is a list of prefixes: one: mono-* four: tetratwo: difive: pentathree: trisix: hexa*Note: mono- is not used on the first element, only the second. Name the following molecular compounds. 1. 2. 3. 4. 5. CO2 CO N2O N2O5 SF6 6. N2O4 7. ClF3 8. PCl5 9. SiO2 10. SO3 Write formulas for the following compounds. 1. sulfur dioxide 2. dinitrogen trioxide 3. nitrogen monoxide 4. sulfur tetrafluoride 5. carbon disulfide 6. oxygen difluoride 62 Lewis Dot Diagrams and Review of Covalent Bonding 1. What is a single covalent bond? 2. What is a double covalent bond? 3. What is a triple covalent bond? 4. What is a dative (coordinate covalent) bond? Electron dot diagrams are ways to represent the 2-dimensional structures of two types of structures formed by covalent bonding, molecules (such as water or carbon dioxide) and polyatomic ions (such as nitrate or ammonium). The symbol for each element represents the nucleus and core or kernel electrons. The dots represent the valence electrons. Steps for constructing Lewis dot diagrams There are many methods for drawing dot diagrams. Your teacher will demonstrate the steps for you. Use the space below to summarize the steps or rules. 1. 2. 3. 4. 5. Examples: HCl C2H6 CH2Cl2 NH3 63 Practice Exercises 1. CH4 2. Cl2 3. H2 4. PH3 5. CHI3 6. SO2 7. CH3OH 8. N2 9. H2Te 10. H2CO 11. OF2 12. HCN 13. BF3 14. PCl3 15. SiO2 16. CO2 17. C2H4 18. C2H2 64 Polyatomic ions: The only difference between these and the regular covalent compounds is you have to account for the charge when you add up your valence electrons. A positive charge means you have less electrons, a negative charge means you have more electrons. 19. ClO3- 20. NO3- 21. ClO2- 22. NO2- 23. ClO- 24. NH4+ 25. SO4-2 26. NH2- 27. H3O+ 28. OH- When you have completed your structures, circle the number of any structure that contains a dative (coordinate covalent) bond. 65 Bond Length and Strength Draw Lewis dot diagrams for the following compounds: 1a. C2H6 b. C2H4 c. C2H2 2. Identify the number of bonds between the two carbon atoms in each of the structures above. a. b. c. 3. Which compound has the strongest C-C bond? 4. Which compound has the longest C-C bond? 5. How are bond lengths and bond strength related? 6. Given the following compounds: CO2, CO and CH3OCH3. (dot diagrams are essential!) a. Which has the strongest C-O bond? b. Which has the longest C-O bond? 7. Given N2 and N2H4. a. Which has the strongest N-N bond? b. Which has the longest N-N bond? 66 Electronegativity Related to Bonding 1. Define electronegativity. 2. What is the trend in electronegativity as you move across the periodic table from left to right? 3. What is the trend in electronegativity as you move down the periodic table from top to bottom? 4. From the following pairs, choose the atom with the highest electronegativity. a. B or C b. Se or Te c. Ca or Ba 5. If the difference in electronegativity between two atoms in a bond is greater than 1.7, the bond is said to be more than 50% ionic/covalent in character. 6. If the difference in electronegativity between two atoms in a bond is less than 1.7, the bond is said to be more than 50% ionic/covalent in character. 7. Is there any bond that is 100% ionic? (Complete transfer of electrons) If so, give an example. 8. Is there any bond that is 100% covalent? (evenly shared electrons) If so, give an example. 9. Define "polar covalent bond." 10. State whether or not the following bonds are ionic, polar covalent or nonpolar covalent: a. C -- F b. O -- O c. N -- O d. Na -- F 11. Draw electron dot diagrams for the following molecular compounds (all covalent bonds) a. H2O b. O2 c. NF3 d CO2 e. C2H2 12. In #11, how many bonds are in each compound? Are they polar, nonpolar or a mixture of both? Number of Bonds polarity a b c d e 67 Bonding and Molecular Geometry Lewis dot diagrams show how valence electrons are involved in bonding. They do not, however, show the geometric arrangements of the atoms in a molecule. We will use VSEPR theory for this purpose. To understand the geometry, we must be able to draw the Lewis diagram and look at the electrons around the central atom. Once we know the geometry, we are better able to predict properties of molecules. 1. What does VSEPR stand for? 2. What is the difference between a bonded pair of electrons and a nonbonded (or unshared) pair of electrons? 3. How many bonded and nonbonded pairs of electrons are on the central atom in the following molecules? (Note: you need to draw the electron dot diagram to be able to answer this) a. H2O b. CO2 c. NH3 d. SO2 4. Why do electron pairs repel one another? 5. If two electron pairs are surrounding a central atom, what is the maximum angle of separation they can achieve? 6. If three electron pairs are surrounding a central atom, what is the maximum angle of separation they can achieve? 7. If four electron pairs are surrounding a central atom, what is the maximum angle of separation they can achieve? 8. Which takes up more room, a bonded pair or nonbonded (unshared) pair of electrons? 9. Fill in the following table: Number of electron pairs on central atom Angle of separation between pairs 2 3 4 68 Name of geometrical shape made by the electron pairs Molecular Geometry and Polarity 1. How does a polar bond differ from a nonpolar bond? 2. How does a polar bond differ from an ionic bond? 3. How is electronegativity difference used to predict bond type? What value separates ionic from polar covalent bonds? 4. What is a dipole (polar molecule)? 5. What two criteria must be met for a molecule to be polar? 6. How can a molecule such as CO2 contain polar bonds yet still be a nonpolar molecule? 7. Fill in the following table: Number of Number of bonds Number of bond angle molecular bonds and (remember unshared (lone, shape unshared pairs multiple bonds unbonded) pairs on central atom count as one bond) 2 2 0 3 3 0 3 2 1 4 4 0 4 3 1 4 2 2 8. Of the shapes you listed in the table above, which can be symmetrical if all bonds are alike? 9. Complete the following table: molecule electron dot diagram number of bonds and unshared pairs geometric shape BF3 NBr3 SO2 CI4 69 polar bonds? polarity of molecule Molecular Geometry/Polarity Review Fill in the following table. Molecule Electron Dot Diagram Molecular Geometry H2S PF3 SiCl4 CO2 SO3 70 Bond Angle Molecular Polarity Allotropes of Carbon 1. Define allotrope: 2. List the names of the three common allotropes of carbon: Navigate to the following website: (The links are posted on Blackboard 24/7) http://www.edinformatics.com/interactive_molecules/diamond.htm The website poses the question, “Why is graphite soft and diamond hard if both are made of pure carbon?” Look at the macromolecule of graphite. Hold down the left mouse button over the structure and move it around. Study the structure of graphite. If you have a scrolling mouse, scroll in and out over the graphic. Enlarge the graphic enough so that the individual atoms are clearly visible. Double click on the center of one atom and move the mouse to the center of the next atom Click once on the second atom. Move the mouse to a third atom and double click. A bond angle should be displayed. Answer the following questions about the structure of graphite: 3. Each carbon atom is bonded to how many others? 4. What is the bond angle between all the C-C-C bonds? 5. What holds carbon atoms in one layer together? 6. What holds the layers to other layers? 7. What are the physical properties of graphite? Now look at the macromolecule of diamond which is further down the web page. Follow the instructions above to determine the bond angle. Answer the following questions about diamond: 8. Each carbon atom is bonded to how many others? 9. What is the bond angle between all the C-C-C bonds? 10. What are the physical properties of diamond? 11. Why is graphite soft and slippery while diamond is so hard when both are pure carbon? Now study the structure of fullerene at the following website: http://www.edinformatics.com/interactive_molecules/fullerene.htm 12. Describe the bonding in fullerene, C60. Silicon is in the same family as Carbon. Study the information at the website below: http://web1.caryacademy.org/chemistry/rushin/studentprojects/elementwebsites/silicon/Structure.htm 13. What is the structure of pure silicon crystals? 14. What are the properties of pure silicon? 15. Research the structure and properties of the common compound, SiO2. Are the properties and structure more like graphite or diamond? 71 Intermolecular Forces 1. Covalent bonds hold atoms/molecules to one another. Intermolecular forces (weak forces) hold atoms/molecules to one another. 2. Intermolecular forces are found in what type of solid? 3. List the three intermolecular forces in increasing order of strength. a. b. c. 4. Describe hydrogen bonding. 5. A molecule must contain a hydrogen atom covalently bonded to what other element(s) in order for a hydrogen bond to form between two molecules? 6. What are the properties of the elements listed in number 5 that give rise to hydrogen bonding? 7. Label the covalent bonds and the hydrogen bonds on the following diagram of HF molecules. H F ---- H F --- H F ---- H F 8. What properties are associated with molecular compounds that contain hydrogen bonds versus those that do not? 9. Describe dipole-dipole attractions. (permanent dipoles) 10. What types of molecules exhibit this type of attraction? 11. What is the only intermolecular force that attracts nonpolar molecules to one another in the solid phase? 12. Describe van der Waals forces (London dispersion forces, or temporary dipoles). 13. What factors determine the magnitude of van der Waals forces? 72 14. What properties are associated with molecules attracted solely by van der Waals forces? 15. For each substance listed below, state the type of intermolecular force that will hold the substance together as a solid. a. H2O b. HCl c. C2H6 d. CH3OH e. CO2 f. H2S g. HF 16. For each pair listed below, choose the one that you would expect to have the highest melting point and give an explanation for your choice. a. H2O or H2S b. C2H6 or C5H12 73 Types of Solids 1. Complete the following table about the four types of solids. Type of Solid Properties Type of Particles Force attracting particles metallic ionic covalent network molecular (including group 18) 2. Describe metallic bonding. How is it different from ionic and covalent bonding? 3. A solid has a high melting point, does not dissolve in water and does not conduct electricity. What type of solid is it most likely to be? 4. A solid has a moderate melting point, does not dissolve in water but conducts electricity. What type of solid is it? 5. A solid has a high melting point, dissolves in water and does not conduct electricity. What type of solid is it? 6. State which type of solid would be formed by each substance listed below. a. Fe g. graphite b. NaCl h. water c. carbon dioxide i. brass d. diamond j. potassium sulfate e. sodium k. sugar, C12H22O11 f. copper (II) chloride l. helium 74 Unit 4 Notes 75 Bonds, Polarity and Solubility Purpose: To relate solubilities of various combinations of substances to their molecular polarities. Pre-lab assignment: Fill in the following table for the substances used in this lab Substance bond type (ionic molecular polarity or covalent) I2 CuCl2 C2H5OH (ethanol) Mineral oil H2O Procedure Part A. 1. Place three clean, dry test tubes in a test tube rack. 2. Put a small amount of water in the first test tube. (Two-finger depth is fine.) Place an equal amount of ethanol in the second test tube, and an equal amount of mineral oil in the third. (***Caution: Avoid breathing the vapors and do not let the liquids come in contact with your skin!!) 3. To each test tube add 1-2 crystals of iodine. (Just the tiny little crystals!!) Stopper the tubes and shake. Record your observations. (Write "soluble" if the crystals dissolved, or "insoluble" if they did not.) Dispose of the liquids in the waste beaker. 4. Repeat steps 1-2. 5. This time add 1-2 crystals of copper (II) chloride to each test tube. Record your observations as before. Part B. 1. You will be using the three solvents that you used in Part A, in addition to hexane, a nonpolar solvent. If two liquids dissolve in one another in all proportions they are said to be miscible. If the two liquids do not dissolve you will see two distinct layers and they are said to be immiscible, and the least dense liquid will float on top of the more dense liquid. For each combination of liquids (see data table B for combinations) pour about 1 ml of each into a clean test tube. Record your observations. (Write either miscible or immiscible.) 76 Data Table Part A Solvent water I2 CuCl2 ethanol mineral oil Data Table Part B Liquid water ethanol mineral oil ethanol mineral oil hexane Questions 1. Why did iodine dissolve in mineral oil and not water? 2. Why did copper (II) chloride dissolve in water but not in mineral oil? 3. Did all substances dissolve in ethanol? For any that did not dissolve, describe the bonding. 4. Alcohols (such as ethanol) are said to be intermediates. Why do you think this is so? 5. Fill in the following blanks: a. Polar solvents tend to dissolve __________ solutes. b. Nonpolar solvents tend to dissolve __________ solutes. c. Polar liquids are miscible with __________ liquids. d. Nonpolar liquids are miscible with __________ liquids. e. Alcohols are miscible with ____________ liquids. f. Ionic compounds only dissolve in _____________ solvents. 77 Evaporation and Intermolecular Attractions In this experiment, Temperature Probes are placed in various liquids. Evaporation occurs when the probe is removed from the liquid’s container. This evaporation is an endothermic process that results in a temperature decrease. The magnitude of a temperature decrease is, like viscosity and boiling temperature, related to the strength of intermolecular forces of attraction. In this experiment, you will study temperature changes caused by the evaporation of several liquids and relate the temperature changes to the strength of intermolecular forces of attraction. You will use the results to predict, and then measure, the temperature change for several other liquids. You will encounter two types of organic compounds in this experiment—alkanes and alcohols. The two alkanes are n-pentane, C5H12, and n-hexane, C6H14. In addition to carbon and hydrogen atoms, alcohols also contain the -OH functional group. Methanol, CH3OH, and ethanol, C2H5OH, are two of the alcohols that we will use in this experiment. You will examine the molecular structure of alkanes and alcohols for the presence and relative strength of two intermolecular forces—hydrogen bonding and dispersion forces. Figure 1 MATERIALS Power Macintosh or Windows PC Vernier computer interface Logger Pro two Temperature Probes 6 pieces of filter paper (2.5 cm X 2.5 cm) 2 small rubber bands (orthodontic bands) masking tape methanol (methyl alcohol) ethanol (ethyl alcohol) 1-propanol 1-butanol n-pentane n-hexane 78 PROCEDURE 1. Obtain and wear goggles! 2. Prepare the computer for data collection. Make sure that the vertical axis has temperature is scaled from -10 to 30C and the horizontal axis is scaled from 0 to 250 seconds. 3. Wrap Probe 1 and Probe 2 with square pieces of filter paper secured by small rubber bands as shown in Figure 1 (rubber bands from the orthodontist work best). Roll the filter paper around the probe tip in the shape of a cylinder. The paper should be even with the probe end. 4. Stand Probe 1 in the 1-propanol container and Probe 2 in the 1-butanol container. Make sure the containers do not tip over. 5. After the probes have been in the liquids for at least 45 seconds, begin data collection by clicking Collect . Monitor the temperature for 15 seconds to establish the initial temperature of each liquid. Then simultaneously remove the probes from the liquids and hold them so the probe tips extend 5 cm over the edge of the table top as shown in Figure 1. 6. When both temperatures have reached minimums and have begun to increase, click Stop to end data collection. Click the Statistics button, , then click OK to display a box for both probes. Record the maximum (t1) and minimum (t2) values for 1-propanol and 1-butanol. 7. For each liquid, subtract the minimum temperature from the maximum temperature to determine t, the temperature change during evaporation. 8. Clear data by selecting “Data” and “Clear All Data”. 9. Repeat Steps 2-8 using methanol (Probe 1) and ethanol (Probe 2). 10. Repeat Steps 2-8 using pentane (Probe 1) and hexane (Probe 2). 10. Clean up. 79 PRE-LAB EXERCISE Prior to doing the experiment, complete the Pre-Lab table. The name and formula are given for each compound. Draw a structural formula for a molecule of each compound. Then determine the molecular mass of each of the molecules. Dispersion forces exist between any two molecules, and generally increase as the molecular mass of the molecule increases. Next, examine each molecule for the presence of hydrogen bonding. Before hydrogen bonding can occur, a hydrogen atom must be bonded directly to an N, O, or F atom within the molecule. Tell whether or not each molecule has hydrogen-bonding capability. Substance Formula ethanol C2H5OH 1-propanol C3H7OH 1-butanol C4H9OH n-pentane C5H12 methanol CH3OH n-hexane C6H14 Structural Formulas 80 Molecular Mass Intermolecular Force DATA TABLE Substance t1 (°C) t2 (°C) t (t1–t2) (°C) ethanol 1-propanol 1-butanol n-pentane methanol n-hexane Data Analysis and Questions We would like to determine the relationship between the strength of the intermolecular forces and the molar mass. From the data you collected, construct a graph that shows this relationship. 81 EXPLAIN (Analysis Questions): 1. Two of the liquids, n-pentane and 1-butanol, had nearly the same molecular weights, but significantly different t values. Explain the difference in t values of these substances, based on their intermolecular forces. 2. Which of the alcohols studied has the strongest intermolecular forces of attraction? The weakest intermolecular forces? Explain using the results of this experiment. 3. Which of the alkanes studied has the stronger intermolecular forces of attraction? The weaker intermolecular forces? Explain using the results of this experiment. 4. Compare the molar masses of butane (C4H10) and pentane. Can you make a prediction of Δt based on these masses? Which of the two will have stronger intermolecular forces? Explain. 5. How is the length of the carbon chain related to the molar mass? How does the length of the carbon chain relate to the strength of the intermolecular forces? 82 EVALUATE 1. Which of the following compounds has the strongest intermolecular forces? (Circle your answer and give an explanation as to why you chose that molecule. Ethylene Glycol isobutane OH OH H C C H H H propane H CH3 H H C C C H H H H H H H HC C C H H H H 2. Which of the following compounds would have the strongest intermolecular forces? Circle your answer and explain your choice. Isobutane (See above for structure) Benzene (C6H6 carbon ring) Propane (See above for structure) 3. Without using any resources, rank the following chemicals in order of increasing boiling point, 7 is the substances with the highest boiling point. Next to the rank, indicate what intermolecular forces are present in each substance. Compound Boiling Point Rank Intermolecular Forces Present NaCl CO2 H2 CH3OH C6H12O6 H2O CH4 4. Based on polarity, which of the following substance would you expect to dissolve in water? CH4 N2 NH3 CaCl2 83 H2S CO2 Study Guide: Topic 4 Bonding 1. Which statement is true for most ionic compounds? A. They contain elements of similar electronegativity. B. They conduct electricity in the solid state. C. They are coloured. D. They have high melting and boiling points. 2. Which fluoride is the most ionic? A. NaF B. CsF C. MgF2 D. BaF2 3. Which statement is a correct description of electron loss in this reaction? 2Al + 3S Al2S3 A. Each aluminium atom loses two electrons. B. Each aluminium atom loses three electrons. C. Each sulfur atom loses two electrons. D. Each sulfur atom loses three electrons. (Total 1 mark) (Total 1 mark) (Total 1 mark) 4. Element X is in group 2, and element Y in group 7, of the periodic table. Which ions will be present in the compound formed when X and Y react together? A. X+ and Y– B. X 2+ and Y– C. X+ and Y2– D. X2– and Y+ (Total 1 mark) 5. What is the formula for the compound formed by calcium and nitrogen? A. CaN B. Ca2N C. Ca2N3 D. Ca3N2 (Total 1 mark) 6. When the species BF2+, BF3 and BF4– are arranged in order of increasing F−B−F bond angle, what is the correct order? A. BF3, BF4–, BF2+ B. BF4–, BF3, BF2+ C. BF2+, BF4–, BF3 D. BF2+, BF3, BF4– 7. Which statement is true for compounds containing only covalent bonds? A. They are held together by electrostatic forces of attraction between oppositely charged ions. B. They are made up of metal elements only. C. They are made up of a metal from the far left of the periodic table and a non-metal from the far right of the periodic table. D. They are made up of non-metal elements only. (Total 1 mark) (Total 1 mark) 84 8. Which molecule is non-polar? A. H2CO B. SO3 C. NF3 D. CHCl3 9. How many electrons are used in the carbon-carbon bond in C2H2? A. 4 B. 6 C. 10 D. 12 (Total 1 mark) (Total 1 mark) 10. What is the valence shell electron pair repulsion (VSEPR) theory used to predict? A. The energy levels in an atom B. The shapes of molecules and ions C. The electronegativities of elements D. The type of bonding in compounds (Total 1 mark) 11. According to VSEPR theory, repulsion between electron pairs in a valence shell decreases in the order A. lone pair-lone pair > lone pair-bond pair > bond pair-bond pair. B. bond pair-bond pair > lone pair-bond pair > lone pair-lone pair. C. lone pair-lone pair > bond pair-bond pair > bond pair-lone pair. D. bond pair-bond pair > lone pair-lone pair > lone pair-bond pair. (Total 1 mark) 12. When C2H4, C2H2 and C2H6 are arranged in order of increasing C–C bond length, what is the correct order? A. C2H6, C2H2, C2H4 B. C2H4, C2H2, C2H6 C. C2H2, C2H4, C2H6 D. C2H4, C2H6, C2H2 (Total 1 mark) 13. In the molecules N2H4, N2H2, and N2, the nitrogen atoms are linked by single, double and triple bonds, respectively. When these molecules are arranged in increasing order of the lengths of their nitrogen to nitrogen bonds (shortest bond first) which order is correct? A. N2H4, N2, N2H2 B. N2H4, N2H2, N2 C. N2H2, N2, N2H4 D. N2, N2H2, N2H4 (Total 1 mark) 14. Which species has a trigonal planar shape? A. CO32– B. SO32– C. NF3 D. PCl3 (Total 1 mark) 85 15. In which substance is hydrogen bonding present? A. CH4 B. CH2F2 C. CH3CHO D. CH3OH 16. Which compound has the highest boiling point? A. CH3CH2CH3 B. CH3CH2OH C. CH3OCH3 D. CH3CHO (Total 1 mark) (Total 1 mark) 17. Which substance is most similar in shape to NH3? A. GaI3 B. BF3 C. FeCl3 D. PBr3 (Total 1 mark) 18. What are responsible for the high electrical conductivity of metals? A. Delocalized positive ions B. Delocalized valence electrons C. Delocalized atoms D. Delocalized negative ions (Total 1 mark) 19. What intermolecular forces are present in gaseous hydrogen? A. Hydrogen bonds B. Covalent bonds C. Dipole-dipole attractions D. Van der Waals’ forces (Total 1 mark) 20. Which statement best describes the attraction present in metallic bonding? A. the attraction between nuclei and electrons B. the attraction between positive ions and electrons C. the attraction between positive ions and negative ions D. the attraction between protons and electrons (Total 1 mark) 21. Which substance has the lowest electrical conductivity? A. Cu(s) B. Hg(l) C. H2(g) D. LiOH(aq) (Total 1 mark) 22. Which substance is most soluble in water (in mol dm–3) at 298 K? A. CH3CH3 B. CH3OCH3 C. CH3CH2OH D. CH3CH2CH2CH2OH (Total 1 mark) 86 23. Arrange the following in decreasing order of bond angle (largest one first), and explain your reasoning. NH2–, NH3, NH4+ (Total 5 marks) 24. (a) An important compound of nitrogen is ammonia, NH3. The chemistry of ammonia is influenced by its polarity and its ability to form hydrogen bonds. Polarity can be explained in terms of electronegativity. (i) Explain the term electronegativity. …………………………………………………………………………………… …………………………………………………………………………………… (2) (ii) Draw a diagram to show hydrogen bonding between two molecules of NH3. The diagram should include any dipoles and/or lone pairs of electrons …………………………………………………………………………………… …………………………………………………………………………………… …………………………………………………………………………………… …………………………………………………………………………………… …………………………………………………………………………………… (3) (iii) State the H–N–H bond angle in an ammonia molecule. ……………………………………………………………………………………… (1) (iv) Explain why the ammonia molecule is polar. …………………………………………………………………………………… …………………………………………………………………………………… …………………………………………………………………………………… (1) 87 (b) Ammonia reacts with hydrogen ions forming ammonium ions, NH4+. (i) State the H–N–H bond angle in an ammonium ion. …………………………………………………………………………………… (1) (ii) Explain why the H–N–H bond angle of NH3 is different from the H–N–H bond angle of NH4+; referring to both species in your answer. …………………………………………………………………………………… …………………………………………………………………………………… …………………………………………………………………………………… …………………………………………………………………………………… …………………………………………………………………………………… …………………………………………………………………………………… (3) (Total 11 marks) 25. (i) Use the VSEPR theory to predict and explain the shape and the bond angle of each of the molecules SCl2 and C2Cl2 …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… (6) 88 (ii) Deduce whether or not each molecule is polar, giving a reason for your answer. …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… (3) (Total 9 marks) 26. Three scientists shared the Chemistry Nobel Prize in 1996 for the discovery of fullerenes. Fullerenes, like diamond and graphite, are allotropes of the element carbon. (i) State the structures of and the bonding in diamond and graphite. …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… (2) (ii) Compare and explain the hardness and electrical conductivity of diamond and graphite. …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… (4) (iii) Predict and explain how the hardness and electrical conductivity of C60 fullerene would compare with that of diamond and graphite. …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… (4) (Total 10 marks) 89 27. (i) List the following substances in order of increasing boiling point (lowest first). CH3CHO C2H6 CH3COOH C2H5OH …………………………………………………………………………………………… (2) (ii) State whether each compound is polar or non-polar, and explain the order of boiling points in (i). …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… (8) (Total 10 marks) 28. State two physical properties associated with metals and explain them at the atomic level. (Total 4 marks) 90 Topic 1: Quantitative Chemistry 1.1 The Mole Concept and Avogadro’s Constant 1.1.1 Apply the mole concept to substances. 1.1.2 Determine the number of particles and the amount of substance (in moles). 1.2 Formulas 1.2.1 Define the terms relative atomic mass (Ar) and relative molecular mass (Mr). 1.2.2 Calculate the mass of one mole of a species from its formula. 1.2.3 Solve problems involving the relationship between the amount of substance in moles, mass and molar mass. 1.2.4 Distinguish between the terms empirical formula and molecular formula. 1.2.5 Determine the empirical formula from the percentage composition or from other experimental data. 1.2.6 Determine the molecular formula when given both the empirical formula and experimental data. 1.3 Chemical Equations 1.3.1 Deduce chemical equations when all reactants and products are given. 1.3.2 Identify the mole ratio of any two species in a chemical equation. 1.3.3 Apply the state symbols (s), (l), (g) and (aq). 1.4 Mass and Gaseous Volume Relationships in Chemical Reactions 1.4.1 Calculate theoretical yields from chemical equations. 1.4.2 Determine the limiting reactant and the reactant in excess when quantities of reacting substances are given. 1.4.3 Solve problems involving theoretical, experimental and percentage yield. 1.4.4 Apply Avogadro’s law to calculate reacting volumes of gases. 1.4.5 Apply the concept of molar volume at standard temperature and pressure in calculations. 1.4.6 Solve problems involving the relationship between temperature, pressure and volume for a fixed mass of an ideal gas. 1.4.7 Solve problems using the ideal gas equation, PV=nRT. 1.4.8 Analyze graphs relating to the ideal gas equation. 1.5 Solutions 1.5.1 Distinguish between the terms solute, solvent, solution and concentration (g dm-3 and mole dm-3). 1.5.2 Solve problems involving concentration, amount of solute and volume of solution. 91 Topic 1 NOTES 92 Introduction to the Mole 1 dozen indicates 12 of something. For instance 1 dozen eggs = 12 eggs. Similarly 1 dozen atoms = 12 atoms Unfortunately, atoms are so small that we need large numbers of them before we can start experimenting with them. Chemists use a unit known as the “mole”. Like a dozen, it simply indicates a specific number of items. Instead of 12 (like the dozen), 1 mole is equal to ….. 602 000 000 000 000 000 000 000 or 6.02 x 1023 This number is extremely important in chemistry and is known as Avogadro’s Number or NA. Complete the following examples: 1.75 dozen atoms = _______ atoms 1.75 moles of atoms = ________ atoms Complete the following problems, showing your calculations. 1. 0.250 moles of atoms = ______ atoms 2. 4.5 x 1025 atoms = _______ moles of atoms 3. 5.67 x 1019 atoms = _______ moles of atoms 4. 0.0056 moles of molecules = _______ molecules 5. 3.5 moles of H2O molecules = _______ molecules 6. 3.5 moles of H2O molecules = _______ atoms 93 Percentage Composition The law of definite proportions states the ratio or proportion by mass of the elements in a compound is always the same. This means that the ratio of hydrogen to oxygen in a sample of water is the same no matter what the sample size. Percentage composition by mass is an easy way to apply this principle. Example: H2O Element present Atomic mass Number of atoms Total mass H 1.01 2 = 2.02 O 16.0 1 = 16.0 18.0 amu total mass Percent Hydrogen in water (part/whole) 100 (2.02 / 18.0 ) 100 = 11.2 % Percent Oxygen in water (part/whole) 100 (16.0 / 18.0 ) 100 = 88.8 % Therefore any sample of water is composed of 11.2 % hydrogen by mass. A 100 g sample contains 11.2 g of hydrogen. A 500 g sample of water contains (500 g ) (11.2 %) = 56 g of hydrogen. Determine the percentage composition of each of the following compounds below. 1. KMnO4 2. HCl 3. Mg(NO3)2 4. (NH4)3PO4 5. Al2(SO4)3 Using your answers from the problems above, answer the following questions: 6. How many grams of oxygen are contained in a 75 g sample of KMnO4? 7. How much magnesium is contained in a 125 g sample of Mg(NO3)2? 8. What mass of aluminum is found in a 25 g sample of Al2(SO4)3? 94 Mass and the Mole Fill in the following table. Substance Symbol Molar Mass (Ar or Mr) g·mol-1 Aluminum Sn Oxygen Iron (II)chloride MgSO4 Ammonium phosphate 1 mole of atoms equals 6.02 x 1023 atoms 1 mole of molecules equals 6.02 x 1023 molecules Avogadro’s number expresses the number of particles but not the mass. How do we find the mass of a mole? The answer is molar mass. Use the relative atomic masses, Ar, of the elements in a compound to find the relative molecular mass of a molecule, Mr. Answer the following questions about atoms, moles and mass. 1. What is the mass of 1.65 moles of gold atoms? 2. How many atoms are found in 2.25 moles of copper? 3. 1.75 moles of Na2O has what mass? 4. How many moles are contained in 175 g of Cu(NO3)2? 5. How many atoms are contained in 2.45 grams of helium? 6. What is the mass of 1.25 × 1022 molecules of CO2? 7. How many molecules are contained in 25.0 g of H2O? 8. How many atoms of hydrogen are contained in 10.0 g of H2O? 95 Empirical and Molecular Formulas An empirical formula is the simplest ratio of atoms in a compound. A molecular formula is the actual number of atoms in a compound. The molecular formula is always a whole number multiple of the empirical formula. 1. What is the empirical formula of the following compounds? a. H2O2 b. C8H18 c. H2SO4 d. N2O4 2. A compound contains 64.3 g of carbon and 10.7 g of hydrogen. What is its empirical formula? 3. A compound has the empirical formula of CH. If its molar mass is 104 g mol-1, what is its molecular formula? 4. A compound is 22.1% aluminum, 25.4% phosphorus, and 52.5% oxygen by mass. Determine its empirical formula. 5. A compound is 69.9% iron and 30.1% oxygen by mass. Determine its empirical formula. 6. A compound is 54.5% carbon, 9.1% hydrogen, and 36.4% oxygen by mass. Its molar mass is 88 g mol-1. What is its empirical and molecular formula? 7. A compound is 43.7% phosphorus and 56.3% oxygen by mass. Its molar mass is 284 g mol-1. Determine its empirical and molecular formula. 8. A hydrated ionic compound is 25.3% copper, 12.9% sulfur, 25.7% oxygen, and 36.1% water. Determine the empirical formula of this hydrate. 96 Percentage Water in a Hydrated Salt Name ______________ Purpose To determine the percent water in a hydrated salt. To determine the empirical formula of the hydrate. Background When water molecules are chemically attached to a compound it is called a hydrate. Water can be separated from the compound by heating it. As water is added back to the compound, the water molecules can reattach. Equipment Evaporating dish Spatula Hot plate Crucible tongs Balance Goggles and apron Materials Copper (II) sulfate hydrate, CuSO4 xH2O Procedure 1. 2. 3. 4. 5. 6. 7. Find the mass of a clean, dry evaporating dish. Record this mass. With the dish still on the balance, press the TARE button to rezero the balance. Add the copper (II) sulfate hydrate until you have exactly 2.00 g. Place the evaporating dish on the hot plate and turn it to a low heat at first. Gradually increase the heat as the hydrate begins to change color. Avoid any popping and spattering. (Immediately remove the evaporating dish if it starts to pop with the crucible tongs and turn down the heat before returning the dish to the hot plate.) After the blue color has disappeared, leave the dish on medium-high heat for an additional three minutes to assure that all of the water has been removed. You may use a spatula to break up any lumps or “caked” portions of the hydrate. If the edges of the solid appear to be turning brown, remove the dish from the heat momentarily and resume heating at a gentler rate. Allow the dish to cool on an insulating square for about a minute. Immediately find the mass of the dish + anhydrous salt, and record the data below. If you would like, you can attempt to reattach the water by adding some drops of water back to the anhydrous salt. Dispose of the copper sulfate in the appropriate container and wash and dry the evaporating dish when you are finished. Data a. b. c. Mass of empty evaporating dish Mass of copper (II) sulfate hydrate Mass of evaporating dish + anhydrous salt anhydrous means without water 97 ________ ±______ g ________ _____ g ________ ± _____ g Calculations-You must show all of your work. (Remember uncertainty!) 1. Find the mass of the anhydrous salt in grams. 2. Find the mass of the water lost in grams. 3. Find the percentage of water in the hydrate: Questions 1. The true value for the percent of water in this hydrate is 36.0%. What is your experimental error? Show your work. 2. Why must you measure the mass of the anhydrous salt immediately upon cooling? 3. Calculate the moles of water. Show your work. 4. Calculate the number of moles of anhydrous salt (CuSO4). Show your work. 5. How many times more moles of water are there than anhydrous salt? (Look at your answers to questions 3 and 4.) Divide your moles of water by your moles of CuSO4 to get a ratio. Based on this answer, what is the true formula for copper (II) sulfate hydrate? (CuSO4 xH2O, where x is the number of moles of water for every one mole of CuSO4) 6. If you didn’t heat your hydrate long enough, all of the water may not have been removed. If this happened, would your moles of water (x, in question #5 above) be too high or too low. Explain your answer. 98 Balancing Chemical Equations 1. ___ Cl2 + ___ NaBr ___ NaCl + ___ Br2 2. ___ Ca(OH)2 + ___ HNO3 ___ Ca(NO3)2 + ___ HOH 3. ___ C2H4 + ___ O2 ___ CO2 + ___ H2O 4. ___ Fe(OH)3 ___ Fe2O3 + ___ HOH 5. ___ P2O5 + ___ H2O ___ H3PO4 6. ___ Al(NO3)3 + ___ NaOH ___ Al(OH)3 + ___ NaNO3 7. ___ KClO3 ___ KCl + ___ O2 8. ___ C3H8 + ___ O2 ___ CO2 + ___ H2O 9. ___ H3PO4 + ___ Mg(OH)2 ___ Mg3(PO4)2 + ___ HOH 10. ___ NH3 + ___ O2 ___ NO + ___ H2O 11. ___ Na2SO4 + ___ Ba(NO3)2 ___ BaSO4 + ___ NaNO3 12. ___ CaO + ___ P2O5 ___ Ca3(PO4)2 13. ___ Al + ___ CuCl2 ___ AlCl3 + ___Cu 14. ___ Ca(OH)2 + ___ H3PO4 ___ Ca3(PO4)2 + ___ HOH 15. ___ NaHCO3 ___ Na2CO3 + ___ H2O + ___ CO2 16. ___ C2H5OH + ___ O2 ___ CO2 + ___ H2O 17. ___ Fe + ___ HCl ___ FeCl3 + ___ H2 18. ___ Co(OH)3 + ___ HNO3 ___ Co(NO3)3 + ___ HOH 19. ___ Mg + ___ O2 ___ MgO 20. ___ Na + ___ Sn(NO3)2 ___ Sn + ___ NaNO3 99 Writing and Balancing Equations Write a balanced equation for each reaction. Classify as: Synthesis (S) – the combination of two or more reactants into one product Decomposition (D) – one reactant forming two or more simpler products Single Displacement (SD) – One element takes the place of another in a compound Double Displacement (DD) – Two compounds exchange their ions Combustion (C) – A compound reacts with oxygen to form oxygen-containing compounds ____ 1. sodium bromide + fluorine sodium fluoride + bromine ____ 2. calcium nitrate + copper (II) sulfate calcium sulfate + copper (II) nitrate ____ 3. potassium chlorate potassium chloride + oxygen gas ____ 4. propane + oxygen carbon dioxide + water ____ 5. calcium bromide + chromium (III) nitrate calcium nitrate + chromium (III) bromide ____ 6. iron + copper (I) nitrate iron (II) nitrate + copper ____ 7. nitrogen + hydrogen ammonia (NH3) ____ 8. hydrogen sulfate water + oxygen + sulfur dioxide ____ 9. zinc sulfide + oxygen zinc oxide + sulfur dioxide ____ 10. ammonium phosphate + lithium hydroxide ammonium hydroxide + lithium phosphate 100 Lab: Types of Chemical Reactions Pre-Lab Discussion There are many kinds of chemical reactions and several ways classify them. One useful method classifies reactions into four major types. These are: 1). Direct combination, or synthesis 2). Decomposition, or analysis 3). Single replacement, or single displacement 4). Double replacement (displacement) or exchange of ions. Not all reactions can be put into one of these categories. Many, however, can. In a synthesis reaction, two or more substances (elements, or compounds) combine to form a more complex substance. Equations for synthesis reactions have the general form A + B AB. For example, the formation of water from hydrogen and oxygenic written, 2H2 + O2 2H2O. A decomposition reaction is the opposite of a synthesis reaction. In decomposition, a compound breaks down into two or more simpler substances (elements or compounds). Equations for decomposition reactions have the form of AB A + B. The breakdown of water into its elements is an example of such a reaction: 2H20 2H2 + O2. In a single replacement reaction, one element in a compound is replaced by another, more active, element. Equations for single replacement reactions have tow general forms. In reactions in which one metal replaces another metal, the general equation is M1 + M2X ---> M1X + M2. In those cases where a nonmetal in a compound is replaced by another nonmetal the general equation is MX + Y MY + X. The following equations illustrate these types of reactions: Zinc metal replaces copper(II) ion: Zn(cr) + CuSO(aq) ZnSO4(aq) + Cu(cr) Chlorine (a nonmetal) replaces bromide ions: KBr(aq) + Cl2(g) KCl(aq) + Br2(l) In a double replacement reaction, the metal ions of two different ionic compounds can be thought of a “replacing one another.” Equations for the type of reaction have the general form AB + CD AD + CB. Most replacement reactions, both single and double, take place in aqueous solutions which contain free ions. In a double replacement reaction, one of the products must be either a precipitate, and insoluble gas, or additional water molecules. An example is the reactions between silver nitrate and sodium chloride in which the precipitate silver chloride is formed: AgNO3(aq) + NaCl(aq) AgCl(cr) + NaNO(aq). The production of a gas is an indication of chemical reaction. A burning splint placed at the mouth of test tube in which hydrogen is being generated will cause a high pitched “pop”. If a burning splint is extinguished when it is placed in a test tube, it indicates the production of carbon dioxide. 101 Purpose In this lab you will observe examples of chemical reactions and will classify each of them into one of the four types of reactions described above. In reactions in which you use a burning splint to test for a gas, the gas produced will be identified. Equipment goggles and apron burner and striker crucible tongs spatula wood splints test tubes test tube holder fine sandpaper evaporating dish Materials zinc, mossy (Zn) copper wire, 10 cm (Cu) magnesium ribbon (Mg) copper (II) carbonate (CuCO3) 6 M hydrochloric acid (HCl(aq)) 1 M copper(II) sulfate (CuSO4) 0.1 M zinc acetate [Zn(C2H3O2)] 0.1 M sodium phosphate (Na3PO4) 1 M sodium carbonate (Na2CO3) Safety Wear safety goggles and apron at all times. You will be working with open flames, heating chemicals, handling acids, and generating gaseous products. Do not handle equipment that may be hot, test it first with the back of your hand. Burning magnesium produces a very bright, hot flame. Make sure you hold the burning metal away from yourself and other students. Do not look directly at it as it burns! Neutralize any acid spill with the sodium bicarbonate solution before you wipe it up. Procedure Reaction 1. Use fine sandpaper to clean a piece of copper wire until the wire is shiny. Note the appearance of the wire in your data table. Using crucible tongs, hold the wire in the hottest part of the burner flame for 1-2 minutes. Examine the wire and note any changes caused by the heating. Reaction 2. Place an evaporating dish near the base of the burner. Examine a piece of magnesium ribbon. Using crucible tongs, hold the magnesium in the flame until it starts to burn. Do not look directly at the flame. Hold the burning magnesium away from you and directly over the evaporating dish. When the ribbon stops burning, put the remains in the evaporating dish. Examine this product carefully. Reaction 3. Place a heaping spatulaful of copper (II) carbonate into a large, clean dry test tube. Note its appearance. Using a test tube holder, heat the CuCO3 strongly for about 3 minutes. Turn off the burner and insert a burning splint into the top part of the test tube. Write your observations of what happens, and the appearances of the test tube and the product formed. 102 Reaction 4. Stand a clean dry small test tube in the test tube rack. Add about 5 ml of 6 M hydrochloric acid to the test tube. Note its appearance. CAUTION. Handle acids with care, they can cause painful burns. Do not inhale any HCl fumes. Obtain and observe a small piece of zinc metal. Carefully place it into the acid in the test tube. Observe and record what happens. Using a test tube holder, invert a second test tube over the mouth of the test tube in which the reaction is taking place. After about 30 seconds, remove the top tube, keeping it inverted, and immediately insert a burning wood splint into its mouth. Record your observations of what happens, and the appearances of the test tube’s interior. Reaction 5. Obtain another small clean dry test tube and add about 5 ml of 1M copper (II) sulfate solution to it. Place a small piece of zinc metal into the solution. Note the appearance of the solution and the zinc before and after the reaction. Reaction 6. To a clean dry test tube, add about 2 ml of 0.1M zinc acetate. Next add about 2 ml of 0.1M sodium phosphate solution to the test tube. Note observations of the solutions before they were combined, and any changes once they were combined, in your data table. Reaction 7. Into another small, clean, dry test tube add about 1 ml of 1M sodium carbonate solution. To this solution (cautiously) add about 10 drops of 6 M HCl. After the reaction has stopped, place a burning wood splint into the test tube, but not so far in that it touches the solution Note any changes. Carefully wash all equipment, being certain no solids go down the sink. Be sure your sink is clean! Return all equipment it to its proper place. Carefully wash your hands before you leave the lab. Post Lab For all seven reactions, classify the reaction and then write a complete balanced equation with phase notation. 103 Ions and Precipitates 1. Write dissociation reactions for the following aqueous solutions. a. LiNO3 b. K2SO4 c. AlBr3 d. FeCl3 2. Predict products and write balanced molecular equations for the following double replacement reactions. Include phase notation to show which product is a precipitate. a. silver nitrate + magnesium chloride b. barium nitrate + aluminum sulfate c. sodium phosphate + calcium bromide d. potassium hydroxide + iron (III) chloride 3. Write ionic and net ionic equations for each of the reactions in number 2. 4. Silver nitrate reacts with sodium chloride to form a white precipitate. Complete the visual representation of the reaction below. (Remember the law of conservation of matter!) = Ag+ = NO3- = Na+ = Cl- 104 Lab- Precipitates and Solubility Rules Purpose: To observe the formation of precipitates, and to write balanced double displacement and net ionic equations. Procedure. You are given a set of 5 solutions. React every combination of two solutions together in the multiwell plate. React only 1 drop of one solution with 1 drop of another solution. This gives ten combinations. Fill in the following data table with “PPT” if a precipitate formed, or “NR” if no reaction was observed. Solution AgNO3 KOH CoCl2 Al2(SO4)3 FeCl3 1) 2) 4) 7) AgNO3 3) 5) 8) KOH 6) 9) CoCl2 10) Post Lab 1. For each reaction that produced a precipitate, PPT, identify the precipitate. Use page 920 in your text to help. (Use the numbers in the data table to identify the reaction.) Ex: 1) AgCl(s) 2. For each reaction that produced a precipitate, write a balanced double displacement reaction using phase notation, (s), for the precipitate. Ex: 1) 3AgNO3 + FeCl3 → 3AgCl(s) + Fe(NO3)3 3. For each reation that produced a precipitate, write a net ionic equation. Ex: 1) Ag+ + Cl- → AgCl(s) 105 Mole-Map Worksheet #1 Mass (g) A molar mass A Moles A coefficient ratio 6.02 x1023 Particles A Moles B molar mass B Mass (g) B 6.02 x1023 coefficient ratio Particles B 2H2(g) + O2(g) 2H2O(g) 1. If 5.0 moles of hydrogen gas are consumed in this reaction, how many moles of water are formed? 2. If 4.0 moles of hydrogen gas are consumed in this reaction, how many moles of oxygen are required? 3. How many molecules of water can be produced from the consumption of 2.0 x 1022 molecules of hydrogen? 4. How many molecules of hydrogen are required to react completely with 4.5 x 1024 molecules of oxygen 5. If 2.5 moles of hydrogen gas are consumed how many molecules of water are produced? 6. How many moles of oxygen gas are required to use up 5.5 x 1023 molecules of hydrogen gas? 7. How many moles of water can be produced from the consumption of 25.0 grams of oxygen? 8. How many grams of water can be produced from the consumption of 10.0 grams of oxygen? 9. How many grams of hydrogen are required to completely consume 5.0 grams of oxygen? 10. How many molecules of water can be produced by the consumption of 2.5 grams of oxygen? 106 Mole-Map Worksheet #2 1. KClO3 KCl + O2 a. Balance the equation. b. If 2.5 grams of potassium chlorate are decomposed, what theoretical mass of oxygen can be formed? c. If 2.5 x 1023 molecules of oxygen are formed, what theoretical mass of potassium chloride can be produced? d. When 45.0 grams of potassium chlorate is decomposed, how many molecules of oxygen are formed? 2. Potassium bromide reacts with chlorine to form potassium chloride and bromine. a. Write a balanced equation for the reaction. b. If 25.0 grams of potassium bromide are reacted, what is the theoretical mass of bromine formed? c. If 15.2 g of bromine are actually produced, what is the percent yield? 3. Sodium phosphate reacts with barium nitrate to form two products. a. Write a balanced equation for the reaction. b. If 10.0 grams of sodium phosphate are reacted, what is the theoretical mass of mass of barium phosphate that could be produced? c. If the actual yield of barium phosphate is 17.8 g, what is the percent yield? 107 Limiting Reactants Let’s say that you want to make a cheese sandwich. It takes 2 slices of bread and 1 slice of cheese to accomplish this task. It can be written as an equation: 2 Br + Ch → Br2Ch (2 bread + 1 cheese make 1 sandwich) How many sandwiches can be made with the following starting ingredients? a. 10 slices of bread and 8 slices of cheese? b. 18 slices of bread and 7 slices of cheese? In “a” and “b” above, which reactant was limiting? a. b. In “a” and “b” above, which reactant was in excess and by how much? a. b. Now let’s try it with a real chemical system: 2 H2 + O2 → 2 H2O For each starting amount given below, state the limiting reactant and how many moles of water can be formed. Starting moles Starting moles Moles of of hydrogen of oxygen water formed 1.0 1.0 3.0 2.0 10.0 4.0 Now let’s add the relationship of mass to the equation. Remember coefficient ratios apply to moles, not mass. Therefore all masses must be converted to moles to clearly see the relationships of amounts of reactants present. Use the same equation above to answer the following questions: 2 H2 + O2 → 2 H2O 1. If 10.0 g of hydrogen and 10.0 g of oxygen are reacted, which reactant is limiting? What mass of water can be produced? 2. If 1.50 g of hydrogen and 50.0 g of oxygen are reacted, which reactant is limiting? What mass of water can be produced? 108 More Limiting Reactant Problems 1. 50.0 g of methane and 50.0 g of oxygen are combusted in a reaction vessel. a. Write a balanced equation to represent this process. b. Determine the moles of each reactant present. c. Determine which reactant is the limiting reactant. d. What mass of carbon dioxide can be formed? 2. The synthesis of ammonia is represented by the following equation: 3 H2(g) + N2(g) → 2 NH3(g) If 15.0 g of hydrogen and 75.0 g of nitrogen are reacted, what is theoretical mass of ammonia can be produced? 3. 125 g of iron and 125 g of oxygen are reacted according to the reaction shown below. 4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) What mass of iron (III) oxide will be formed? For the reactant in excess, what mass will be left unreacted? (You must calculate the amount consumed to determine this mass!!!) 109 An Introduction to the Gas Laws (internet reqd.) Boyle’s Law We are going to take a peek into properties of gases and the relationships between these properties. For an introduction, navigate to the following site and scroll through the animation. http://preparatorychemistry.com/Bishop_KMT_frames.htm Once you have read through the introductory material in the animation, look at the menu list on the left side of the page. (There is a scroll bar as well.) Scroll down to Chapter 13, and click on the “Boyle’s Law animation”. As you read through the pages in the animation answer the following questions: 1. What are the four inter-related properties of a gas? a. b. c. d. 2. What creates pressure in a gas? 3. In Boyle’s Law, what two properties are we changing? a. b. 4. In Boyle’s Law, what two properties are we keeping constant? a. b. 5. When the volume is cut in half, what happens to the pressure? 6. As the volume decreases, the pressure ___________. Is this a direct or inverse relationship? Now that we have the concepts, let’s collect some data. Navigate to the following website: http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/gaslaw/boyles_law_graph.html You will see a device for measuring the pressure and volume of a gas. We will use the mouse to drag the plunger to change the volume of the gas. Hold down the mouse and drag it to the desired volume. When you release the mouse it will automatically record a data point. Collect a total of 6 data points from a variety of different volumes. Start with a small volume and gradually increase the volume in your experiment. Record your data points in the table below: Volume (mL) Pressure (psi) 110 Once you have a variety of points recorded, press the “Graph” button on the animation. Sketch this graph on the axes below: This graph shows the relationship P×V = constant Pressure Volume (mL) Boyle’s Law Summary: If temperature and number of particles of gas are kept constant, the pressure and volume of a gas are (directly/inversely) related. This means if you increase the volume, the pressure will _____________. Let’s look at our next gas law: Charles’ Law Navigate back to http://preparatorychemistry.com/Bishop_KMT_frames.htm Scroll down to Chapter 13, Charles’ Law and click through the animation. Answer the following questions as you go along. 1. What two properties are you keeping constant in Charles’ Law? a. b. 2. What two properties are you changing? a. b. 3. When you increase the temperature of a gas, what is happening to the velocity of the particles? 4. To keep the pressure (number of collisions) constant, when the temperature increases what must happen to the volume? 5. As temperature increases, volume _______________. 6. Is this a direct or inverse relationship? 7. Would a graph of temperature vs. volume be a curve like Boyle’s law or a straight line? Sketch the graph below: This graph represents the relationship T V ÷ T = constant V 8. If the temperature drops to zero Kelvin, what happens to the volume of the gas? 111 9. What temperature is absolute zero in Kelvins? In ºC? 10. What happens to molecular motion at absolute zero? Avogadro’s Principle Navigate back to http://preparatorychemistry.com/Bishop_KMT_frames.htm Scroll down to Chapter 13, “Volume/moles animation” and click through the animation. Answer the following questions as you go along. You recognize the name from the number of particles in a mole, but Avogadro also did a lot of experimenting with gases. 1. What two properties of a gas are changing ? a. b. 2. What two properties must be kept constant? a. b. 3. As the number of moles of gas increases what happens to the volume? 4. Is this an inverse or direct relationship? 5. Would a graph of temperature vs. volume be a curve like Boyle’s law or a straight line? Sketch the graph below: This graph represents the relationship n V÷ n = constant V 6. Equal volumes of gas at the same temperature and pressure contain ________ moles of gas. 7. Which statement is true about 1 L of hydrogen gas at the same temperature and pressure as 1 L of oxygen gas? Circle the best answer. a. 1 L of hydrogen contains the most moles of gas. b. 1 L of oxygen contains the most moles of gas. c. Both gases contain the same number of moles of gas. 8. Which gas in number 7 above has the highest mass? Summarizing Graphical Relationships Navigate to the following website and play with the variables. Freeze two of the variables (click on them in the upper left hand corner) and determine the relationship between the remaining two variables. http://www.grc.nasa.gov/WWW/K-12/airplane/Animation/frglab2.html A straight line on a graph represents what type of relationship? A descending curved line on a graph represents what type of relationship? 112 Dalton’s Law Dalton’s law is a little different than the ones we learned above. Unlike Boyle’s, Charles’ and Avogadro’s Laws, this one deals with a mixture of gases. Navigate to the website below and click through the 8 pages and answer the practice questions. http://www.wwnorton.com/college/chemistry/gilbert2/tutorials/interface.asp?chapter=chapter_ 06&folder=daltons_law 1. What is the equation for Dalton’s Law? 2. In the space below show your work and answers for the practice questions on section 4 of the web site. Practice question 1: Practice question 2: Practice question 3: Practice question 4: 3. If a mixture of oxygen and nitrogen has a total pressure of 2.50 atm and the partial pressure of the oxygen is 1.73 atm, what is the partial pressure of the nitrogen? 4. A mixture of gases is 80% helium and 20% neon (by moles). If the total pressure is 760 mm Hg, what is the partial pressure of the helium? 5. 2.00 moles of gas A, 3.00 moles of gas B and 5.00 moles of gas C have a total pressure of 800 mm Hg. What is the partial pressure of each individual gas? Extra information about pressure and its units! Pressure is measured in many units. The official SI unit is the kilopascal (kPa). We also use two other units as well, the atmosphere (atm) and millimeter of mercury (mm Hg). Standard pressure (close to the air pressure in the room on any given day) is defined as the following values: 101.3 kPa = 1.00 atm = 760 mm Hg Using dimensional analysis, convert the following pressure values into the desired unit. 1. 75 kPa = _____ atm 3. 655 mm Hg = _____ atm 2. 1.55 atm = _____ kPa 4. 128 kPa = _____ mm Hg 113 Ideal Gas Law The ideal gas equation relates all four measurable properties of a gas. If you know three of the properties, this equation allows you to calculate the fourth. It has the form: PV = nRT List what each of the letters represents and the units for each: P= V= n= T= R is defined as the gas law constant. Its value does not change. State the value and units for R: Solve the following problems using the ideal gas equation. 1. What volume does 1.25 moles of gas occupy at 35ºC and 145 kPa? 2. What is the pressure of 2.38 moles of gas that has a volume of 38 dm3 at a temperature of -55ºC? 3. What volume does 10.0 g of carbon dioxide occupy at 75ºC and 85 kPa? 4. How many grams of sulfur dioxide are contained in a 2.50 dm3 container at a pressure of 135 kPa and a temperature of 28ºC? 5. What is the molar mass of a 2.50 g sample of gas that occupies 375 cm3 at a temperature of 15ºC and a pressure of 175 kPa? 114 Gas Laws -- Calculations 1. A gas occupies 1.00 cm3 at STP. What volume does it occupy at 710 mm Hg and 55C? 2. What volume does a gas occupy at 1.50 atm, if at standard pressure it has a volume of L? 3. A gas occupies 48.0 mL at 75C. What volume does it occupy at 15C? 4. A gas occupies 0.75 dm3 at 125C and 95.5 kPa. What volume does it occupy at STP? 5. Two gases are mixed in a 1.00 L container. The total pressure of the combined gases is 2.25 atm. If the pressure of one gas is 0.75 atm, what is the pressure of the second gas? 6. A mixture of three gases (A,B and C) has a pressure of 800 mm Hg. There are 1.00 mole of gas A, 0.75 mole of gas B, and 0.25 mole of gas C. What are the partial pressures of each gas? 7 What is the volume of 16.0 g of oxygen gas at 25 C and 135 kPa? 8 A fixed amount of gas has its temperature doubled and its pressure doubled. What happens to its volume? 9 A fixed amount of gas has its temperature tripled and its pressure cut in half. What happens to its volume? 115 25.0 Molar Volume of a Gas 1. What does STP stand for? 2. What is the volume of 1.0 mole of ANY gas at STP? (Remember units!!) 3. State Avogadro's hypothesis. 4. Solve the following problems using your roadmap and the following balanced equation C2H4(g) + 3 O2(g) 2 CO2(g) + 2 H2O(g) a. What volume does 8.0 g of carbon dioxide gas occupy at STP? b. What is the volume of 1.50 x 1022 molecules of oxygen gas at STP? c. What volume of carbon dioxide is produced when 10.0 dm3 of oxygen is consumed? d. What mass of water is produced when 10.0 dm3 of ethene (C2H4) at STP undergoes combustion? e. What volume of carbon dioxide gas at STP is produced when 24.0 g of oxygen gas is reacted? f. How many molecules of water are produced by the consumption of 10.0 L of ethene gas at STP? 116 Molarity Define the following terms: SoluteSolventSolution ConcentrationMolarityDilution- Solve the following problems. 1. What is the molarity of 1.75 dm3 of solution that contains 2.45 moles of NaCl? 2. What is the molarity of 0.755 dm3 of solution that contains 25 g of NaCl? 3. 1.5 dm3 of 0.35 mol dm-3 solution contains what mass of KNO3? 4. What volume of 1.45 mol dm-3 solution contains 84.0 g of MgBr2? 5. What is the molarity of 282 cm3 of solution that contains 18.7 g of potassium oxide? 6. What mass of lithium nitrate is contained in 583 mL of 2.50 mol dm-3 solution? 117 7. What is the concentration of Cl- in a 2.50 mol dm-3 solution of AlCl3? 8. What is the concentration of K+ ions in a solution prepared with 10.0 g of potassium oxide in 255 cm3 of solution? 9. 15 mL of 12 mol dm-3 HCl solution is diluted to a total volume of 225 mL. What is the new concentration? 10. What volume of 6.0 M HCl solution must be added to water to make 365 mL of 1.75 M solution? 11. 2 NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2 H2O(l) Using the reaction above, what mass of H2O can be formed from the reaction of 100.0 cm3 of 1.25 mol dm-3 H2SO4 solution with excess NaOH solution? 12. 2 NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2 H2O(l) Using the reaction above, what mass of H2O can be formed from the reaction of 150.0 cm3 of 1.00 mol dm-3 H2SO4 solution with 200.0 cm3 of 2.00 mol dm-3 NaOH solution? (This is a limiting reactant problem!) 118 Stoichiometry Review Use the following equation to solve problems 1-3. N2(g) + 3 H2(g) → 2 NH3(g) 1. How many molecules of NH3 can be produced from the reaction of 10.0 dm3 of N2 at STP with excess hydrogen? 2. If 2.0 L of N2 gas and 5.0 L of H2 gas are reacted together, what is the theoretical volume of NH3 gas that can be produced? Assume constant temperature and volume. 3. In problem #2, which reactant is in excess and what volume will be left unreacted? Use the following equation to solve problems 4-5 Ba(OH)2(aq) + 2 HCl(aq) → BaCl2(aq) + 2 H2O(l) 4. What mass of water will be produced from the reaction of 100.0 cm3 of 0.100 mol dm-3 Ba(OH)2 solution and 100.0 cm3 of 0.100 mol dm-3 HCl? 5. What volume of 0.250 mol dm-3 HCl would be required to completely neutralize 1.25 dm3 of 0.500 mol dm-3 Ba(OH)2? 119 Topic 1 NOTES 120 Study Guide: Topic 1 Quantitative Chemistry 1. Which of the following quantities has units? A. Relative atomic mass B. Relative molecular mass C. Molar mass D. Mass number (Total 1 mark) 2. What is the total number of atoms in 0.20 mol of propanone, CH3COCH3? A. 1.2×1022 B. 6.0×1023 C. 1.2×1024 D. 6.0×1024 (Total 1 mark) 3. Which contains the same number of ions as the value of Avogadro’s constant? A. 0.5 mol NaCl B. 0.5 mol MgCl2 C. 1.0 mol Na2O D. 1.0 mol MgO (Total 1 mark) 4. Which is a correct definition of the term empirical formula? A. formula showing the numbers of atoms present in a compound B. formula showing the numbers of elements present in a compound C. formula showing the actual numbers of atoms of each element in a compound D. formula showing the simplest ratio of numbers of atoms of each element in a compound (Total 1 mark) 5. The complete oxidation of propane produces carbon dioxide and water as shown below. C3H8 + __O2 __CO2 + __H2O What is the total of the coefficients for the products in the balanced equation for 1 mole of propane? A. 6 B. 7 C. 12 D. 13 (Total 1 mark) 6. When the equation below is balanced for 1 mol of C3H4, what is the coefficient for O2? C3H4 + O2 CO2 + H2O A. 2 B. 3 C. 4 D. 5 (Total 1 mark) 121 7. The equation for a reaction occurring in the synthesis of methanol is CO2 + 3H2 CH3OH + H2O What is the maximum amount of methanol that can be formed from 2 mol of carbon dioxide and 3 mol of hydrogen? A. 1 mol B. 2 mol C. 3 mol D. 5 mol (Total 1 mark) 8. Lithium hydroxide reacts with carbon dioxide as follows. 2LiOH + CO2 → Li2 CO3 + H2O What mass (in grams) of lithium hydroxide is needed to react with 11 g of carbon dioxide? A. 6 B. 12 C. 24 D. 48 (Total 1 mark) 9. 3.0 dm3 of sulfur dioxide is reacted with 2.0 dm3 of oxygen according to the equation below. 2SO2(g) + O2(g) → 2SO3(g) What volume of sulfur trioxide (in dm3) is formed? (Assume the reaction goes to completion and all gases are measured at the same temperature and pressure.) A. 5.0 B. 4.0 C. 3.0 D. 2.0 (Total 1 mark) 10. Which change in conditions would increase the volume of a fixed mass of gas? A. B. C. D. Pressure /kPa Doubled Halved Doubled Halved Temperature /K Doubled Halved Halved Doubled (Total 1 mark) 11. The temperature in Kelvin of 2.0 dm3 of an ideal gas is doubled and its pressure is increased by a factor of four. What is the final volume of the gas? A. 1.0 dm3 B. 2.0 dm3 C. 3.0 dm3 D. 4.0 dm3 (Total 1 mark) 12. What volume (in dm3) of 0.30 mol dm–3 NaCl solution can be prepared from 0.060 mol of solute? A. 0.018 B. 0.20 C. 0.50 D. 5.0 (Total 1 mark) 122 13. Which solution contains 0.1 mol of sodium hydroxide? A. 1 cm3 of 0.1 mol dm–3 NaOH B. 10 cm3 of 0.1 mol dm–3 NaOH C. 100 cm3 of 1.0 mol dm–3 NaOH D. 1000 cm3 of 1.0 mol dm–3 NaOH (Total 1 mark) 14. Assuming complete reaction, what volume of 0.200 mol dm–3 HCl(aq) is required to neutralize 25.0 cm3 of 0.200 mol dm–3 Ba(OH)2(aq)? A. 12.5 cm3 B. 25.0 cm3 C. 50.0 cm3 D. 75.0 cm3 (Total 1 mark) 15. An organic compound A contains 62.0 by mass of carbon, 24.1 by mass of nitrogen, the remainder being hydrogen. (i) Determine the percentage by mass of hydrogen and the empirical formula of A. ................................................................................................................................... ................................................................................................................................... ................................................................................................................................... ................................................................................................................................... ................................................................................................................................... ................................................................................................................................... (3) (ii) Define the term relative molecular mass. ................................................................................................................................... ................................................................................................................................... ................................................................................................................................... ................................................................................................................................... (2) (iii) The relative molecular mass of A is 116. Determine the molecular formula of A. ................................................................................................................................... ................................................................................................................................... (1) (Total 6 marks) 123 16. Copper metal may be produced by the reaction of copper(I) oxide and copper(I) sulfide according to the below equation. 2Cu2O + Cu2S 6Cu + SO2 A mixture of 10.0 kg of copper(I) oxide and 5.00 kg of copper(I) sulfide was heated until no further reaction occurred. (a) Determine the limiting reagent in this reaction, showing your working. .................................................................................................................................... .................................................................................................................................... .................................................................................................................................... .................................................................................................................................... .................................................................................................................................... .................................................................................................................................... (3) (b) Calculate the maximum mass of copper that could be obtained from these masses of reactants. .................................................................................................................................... .................................................................................................................................... .................................................................................................................................... .................................................................................................................................... (2) (Total 5 marks) 124 Topic 5: Energetics 5.1 Exothermic and Endothermic Reactions 5.1.1 Define the terms exothermic reaction, endothermic reaction and standard enthalpy change of reaction (∆H°). 5.1.2 State that combustion and neutralization are exothermic processes. 5.1.3 Apply the relationship between temperature change, enthalpy change and the classification of a reaction as endothermic or exothermic. 5.1.4 Deduce, from an enthalpy level diagram, the relative stabilities of reactants and products, and the sign of the enthalpy change for the reaction. 5.2 Calculation of Enthalpy Changes 5.2.1 Calculate the heat energy change when the temperature of a pure substance is changed. 5.2.2 Design suitable experimental procedures for measuring the heat energy changes of reactions. 5.2.3 Calculate the enthalpy change for a reaction using experimental data on temperature changes, quantities of reactants and mass of water. 5.2.4 Evaluate the results of experiments to determine enthalpy changes. 5.3 Hess’s Law 5.3.1 Determine the enthalpy change of a reaction that is the sum of two or three reactions with known enthalpy changes. 5.4 Bond Enthalpies 5.4.1 Define the term average bond enthalpy. 5.4.2 Explain, in terms of average bond enthalpies, why some reactions are exothermic and others are endothermic. 125 Topic 5 NOTES 126 Average Bond Enthalpy and Enthalpy of Reaction (ΔH) 1. Define bond enthalpy. 2. When a bond is formed energy is absorbed/released. (choose one) 3. When a bond is broken energy is absorbed/released. 4. In a chemical reaction bonds are both broken and formed. The difference in the energies of the bonds is known as the heat of reaction, ΔH. If energy is released during a chemical reaction it is said to be endothermic/exothermic. 5. If energy is absorbed the reaction is said to be endothermic/exothermic. 6. If the heat of reaction is negative the reaction is said to be endothermic/exothermic. 7. If the heat of reaction is positive the reaction is said to be endothermic/exothermic. 8. In an exothermic reaction the products/reactants have greater energy. 9. In an endothermic reaction the products/reactants have greater energy. 10. In an exothermic reaction the surrounding’s temperature goes up/down. 11. In an endothermic reaction the surrounding’s temperature goes up/down 12. In the formation of H2O from its elements, what bonds must be broken? (look at the visual representation of the reaction) 2H2 + O2 → 2H2O 13. What bonds are formed in the equation above? 14. The H-H bond energy is 432 kJ/mol. The O=O bond energy is 495 kJ/mol. How much energy is required to break up the reactants in this reaction? 15. The O-H bond energy is 467 kJ/mol. How much energy will be realeased in the formation of the products of this reaction? 16. Calculate the H for this reaction. ΔH = sum of bonds broken - sum of bonds formed. 17. Is this reaction endothermic or exothermic? How do you know? 18. Which are more stable, the reactants or products? 127 128 Enthalpy of Reaction Worksheet 1. a. Write a balanced equation for the combustion of ethyne (C2H2). b. Draw dot diagrams for each substance in your balanced equation. c. Determine the bond enthalpies of the bonds broken in the reaction. (use the table on the preceding page) d. Determine the bond enthalpies of the bonds formed. e. Determine the enthalpy of reaction, ΔH, for the reaction. f. All combustion reactions are (exothermic/endothermic). g. Which have more energy, the reactants or products? h. Which are more stable, the reactants or products? 2. Determine the enthalpy of reaction for the following reaction: N2(g) + 3 H2(g) → 2 NH3(g) 129 Enthalpy Level Diagrams Answer the questions by referring to the diagrams of the potential energy of a reaction. Potential Energy Diagram #1 Potential Energy Diagram #2 75 D XY 50 C 25 XY 50 X+ Y A D A E B B X+Y 25 C 0 E 0 Reaction Coordinate (X + Y XY) Reaction Coordinate (XY X + Y) Identify the letter that describes the following for each: Diagram #1 Letter Diagram #2 Value Value Letter (kJ) (kJ) Total Energy of the Reactants Total Energy of the Products Total Energy of the Activated Complex Heat of Reaction Indicate if this reaction is endothermic or exothermic 1. If a catalyst were added, would the activation energy increase, decrease, or remain the same? 2. If a catalyst were added, would the heat of reaction increase, decrease, or remain the same? Use the axes below to draw a potential energy diagram for the following conditions. The potential energy of the reactants is 100 kJ. The potential energy of the products is 50 kJ. The activation energy is 25 kJ. Potential Energy Potential Energy (kJ) 75 100 50 Reaction Coordinate 1. Is this reaction exothermic or endothermic? 2. What is the value of the heat of reaction? 3. Draw a dotted line representing the effect of adding a catalyst. 130 Enthalpy of Reaction Worksheet 2Al + Fe2O3 → Al2O3 + 2Fe 1. a. b. c. d. e. ΔH = -850 kJ Is this reaction exothermic or endothermic? Is heat a reactant or product in this reaction? What type of reaction is this? How many moles of aluminum are required to produce 850 kJ? How much heat energy is produced when 20.0 g of aluminum are reacted? f. How much heat energy is produced when 20.0 g of aluminum oxide are produced? g. What mass of iron is produced if 20.0 g of iron (III) oxide are consumed? N2 + 2O2 → 2NO2 2. a. b. c. d. e. ΔH = 67.7 kJ Is this reaction exothermic or endothermic? Is heat a reactant or product in this reaction? What type of reaction is this? Which are more stable in this reaction, the reactants or products? How much heat is involved in the production of 10.0 g of NO2? f. How many molecules of O2 are required to produce 10.0 g of NO2? 131 Energy and Calorimetry 1. What is the SI unit of energy? 2. The calorie is a non-SI unit of energy commonly used. How are the calorie and the joule related? 3. Convert the following units: a. 425 J = _____ cal b. 275 cal = _____ J 4. How does a calorie relate to a nutritional Calorie? 5. How many joules are contained in a slice of chocolate cake with 255 Calories? 6. Define specific heat. 7. What is the specific heat of water in joules and in calories? (Remember to use all appropriate units) 8. In the equation q = m CP T, what do each of the letters represent? (include units) q = _____________ m = _____________ CP = _____________ T = _____________ Solve the following problems using the specific heat data listed in the Appendix of your textbook. 9. 40.0 grams of water are heated from 10.0C to 30.0C. Determine the heat absorbed in joules. 10. 25.0 grams of aluminum are heated from 15.0C to 35.0C. Determine the heat absorbed. 11. 1.5 kg of lead is cooled from 136.4C to 21.7C. Determine the heat released. 12. What mass of copper can be heated by 3.00 x 103 J, if the temperature increases by 35.0C? 13. What is the change in temperature of 200.0 g of mercury that absorbs 4.20 x 103 J? 14. What is the final temperature of the mixture when 25.0 g of 15.0C water is mixed with 74.0 g of 95.0C water? 15. What is the final temperature of the mixture when a 55.0 g iron bar at 115C is placed into 155 g of water at 25.0C? 132 Heat of Reaction Problem Set 1. A 0.250 mol dm-3 solution of HCl is added to 1.00 mol dm-3 NaOH. 100.0 ml of each solution is used. A 13.0 degree temperature increase is noted. Calculate the energy absorbed or released in kJ mol-1 of HCl. 2. A sample of 3.50 g of NaOH is added to 125.0 ml of 1.00 mol dm-3 HCl. The temperature increase was observed to be 5.00 degrees. Calculate the energy absorbed or released in kJ mol-1 of HCl. 3. A 10.0 g sample of NH4Cl is added to 100.0 ml of water. The temperature decrease was observed to be 5.00 degrees. Calculate the energy absorbed or released in kJ mol-1 of NH4Cl. 4. A 1.50 g sample of zinc is added to 50.0 ml of 0.250 mol dm-3 copper (II) sulfate solution. The temperature increase was observed to be 7.00 degrees. Calculate the energy absorbed or released in kJ mol-1 of zinc. 5. Equal volumes of 6.00 mol dm-3 HNO3 (nitric acid) and 1.00 mol dm-3 KOH are mixed. The total volume of mixture is 500.0 cm3. The temperature increase was observed to be 15.0 degrees. Calculate the energy absorbed or released in kJ mol-1 of KOH. 133 6. A 1.25 mol dm-3 solution of HCl is added to 2.00 mol dm-3 NaOH. 200.0 ml of each solution is used. A 17.0 degree temperature increase is noted. Calculate the energy absorbed or released in kJ mol-1 of HCl. 7. A sample of 10.5 g of NaOH is added to 125.0 ml of 1.00 mol dm-3 HCl. The temperature increase was observed to be 5.00 degrees. Calculate the energy absorbed or released in kJ mol-1 of HCl. 8. A 20.0 g sample of NH4NO3 is added to 100.0 ml of water. The temperature decrease was observed to be 8.00 degrees. Calculate the energy absorbed or released in kJ mol-1 of NH4NO3. 9. A 1.50 g sample of zinc is added to 150.0 ml of 0.250 mol dm-3 copper (II) sulfate solution. The temperature increase was observed to be 10.0 degrees. Calculate the energy absorbed or released in kJ mol-1 of copper (II) sulfate. 10. Equal volumes of 1.00 mol dm-3 HNO3 (nitric acid) and 1.00 mol dm-3 Ca(OH)2 are mixed. The total volume of mixture is 1000.0 cm3. The temperature increase was observed to be 25.0 degrees. Calculate the energy absorbed or released in kJ mol-1 of HNO3. 134 Hess’s Law 1. Calculate the heat of reaction for the combustion of nitrogen monoxide gas, NO, to form nitrogen dioxide gas, NO2, as given in the following thermochemical equation. 2 NO (g) + O2 (g) 2 NO2 (g) You are given the following heat of formation data: N2 (g) + O2 (g) 2 NO (g) H0 = 180.58 kJ N2 (g) + 2 O2 (g) 2 NO2 (g) H0 = 66.4 kJ 2. Calculate the H for the following reaction 2 N2 (g) + 5 O2 (g) 2 N2O5 (g) Use the following date in your calculations: 2 H2 (g) + O2 (g) 2 H2O (l) N2O5 (g) + H2O (l) 2 HNO3 (l) N2 (g) + 3 O2 (g) + H2 (g) 2 HNO3 (l) H0 = -571.6 kJ H0 = -76.6 kJ H0 = -348.2 kJ 3. Calculate the heat of formation for pentane, C5H12. 5 C (s) + 6 H2 (g) C5H12 (g) Use the following data for your calculations: C (s) + O2 (g) CO2 (g) 2 H2 (g) + O2 (g) 2 H2O (l) C5H12 (g) + 8 O2 (g) 5 CO2 (g) + 6 H2O (l) H0 = -393.5 kJ H0 = -571.6 kJ H0 = -3535.6 kJ 135 4. Calculate the heat of formation for sulfur dioxide, from its elements. Write the balanced equation you are looking for: Use the following data for your calculations: 2 S (s) + 3 O2 (g) 2 SO3 (g) 2 SO2 (g) + O2 (g) 2 SO3 (g) 5. Calculate the ΔH for the process A 2C + E Use the following numbered processes: 1. A 2B ΔH1 2. B C +D ΔH2 3. E 2D ΔH3 Express your answers in terms of ΔH1, ΔH2, and ΔH3. 136 H0 = -790.2 kJ H0 = -198.2 kJ The Enthalpy of Decomposition of Hydrogen Peroxide Purpose: To experimentally determine the enthalpy of reaction using a calorimeter. To use Hess’ Law to verify experimental data. Equipment: Graduated cylinder PC interfaced thermometer H2O2 solution (3.00% by mass) stirring rod or magnetic stirrer calorimeter (Styrofoam cup!) 0.500 M Fe(NO3)3 solution There is no single instrument that can directly measure heat in the way a balance measures mass or a thermometer measures temperature. However, it is possible to determine the heat change when a chemical reaction occurs. The change in heat is calculated from the mass, temperature change and specific heat of the substance which gains or loses heat. The equation that is used to calculate heat gain or loss is: q = (grams of substance ×(specific heat) × (ΔT) Where q = the heat gained or lost and ΔT is the change in temperature. Since ΔT = (final temperature minus initial temperature), an increase in temperature will result in a positive value for both ΔT and q, and a loss of heat will give a negative value. A calorimeter itself will participate in the transfer of heat. It will be assumed that this heat transfer will be minimal and will be neglected in calculations. ΔH is related to the amount of substance involved in a reaction. To determine ΔH from q, you will need to divide the value of q by the moles of substance. In this lab you will be determining the enthalpy of reaction (in kJ/mol) for the decomposition of hydrogen peroxide which slowly decomposes into water and oxygen gas. Procedure Add 50.0 mL of 3.0% H2O2 solution into the calorimeter. Place a temperature probe into the calorimeter and secure it with a clamp so it doesn’t overturn. Connect the probe to a computer and set it up to record the temperature for 15 minutes. (Your teacher will help you with this.) Start the data collection and stir continuously. After two minutes have elapsed, add 10.0 mL of 0.50 M Fe(NO3)3 solution to the solution. Let the computer continue to collect data for the remaining time. Continue to stir! Print a copy of the resulting temperature vs. time graph. Calculations (show your work and uncertainty) 1. Determine the moles of H2O2 present in the 50.0 mL sample. (Assume the density of the solution is the same as water, 1.00 g/mL.) 2. Determine the heat transferred to the solution in the calorimeter, qcalorimeter. (The specific heat of the solution can be assumed to be the same as that of water.) 137 3. From the law of conservation of energy, qreaction = - qcalorimeter. Determine qreaction. 4. Determine the heat of reaction, ΔH, for the decomposition of hydrogen peroxide in joules per mole. 5. Write the balanced equation for the decomposition of hydrogen peroxide. Use ΔH notation. Post Lab 1. How does the graphical temperature analysis improve the accuracy of your data? 2. What is the purpose of the Fe(NO3)3 solution? Why is it not included in the chemical equation for the decomposition of hydrogen peroxide? 3. 2H2 + O2 → 2H2O ΔH = -572 kJ Look at the equation written directly above. Look at the equation you wrote in calculation number 5. Using these two equations and Hess’ Law, calculate the ΔH for the following overall equation: H2 + O2 → H2O2 ΔH = ? Hint: rearrange the equations to make the third equation. 4. The true value of ΔH for reaction H2 + O2 → H2O2, is -187.5 kJ. Calculate your percent error. 138 Energy Change in Phase Changes 1. Define "heat (enthalpy) of vaporization"-2. Define "heat (enthalpy) of fusion"-3. What are the symbols used for heat of vaporization and heat of fusion? 4. When a substance is strongly attracted to other molecules in the liquid state, it will have a (high/low) heat of vaporization. 5. Which would you expect to have a higher heat of vaporization, H2O or Cl2? (Think about the type of intermolecular force holding the molecules together) Explain. 6. What are the values of the heat of vaporization and the heat of fusion of water? Remember to use the proper units. 7. How much energy does it take to vaporize 2.5 moles of water at 100ºC? 8. How much energy does it take to melt 35.0 g of water at 0ºC? 9. What happens to the temperature (kinetic energy) during a phase change? 10. List 3 phase changes that are exothermic. 10. Draw a heating diagram for heating 45.0 g of ice from a temperature of -25.0º C to steam at 110ºC. temperature (ºC) time 11. From the information given in problem number 10, how much energy is involved in heating the ice from -25.0º C to steam at 110ºC? Hint: this will take five steps! 139 Study Guide: Topic 5 Energetics 1. According to the enthalpy level diagram below, what is the sign for H and what term is used to refer to the reaction? H reactants products reaction progress H reaction A. positive endothermic B. negative exothermic C. positive exothermic D. negative endothermic (Total 1 mark) 2. Which statement is correct about the reaction shown? 2SO2(g) + O2(g) 2SO3(g) H = –196 kJ A. B. C. D. 196 kJ of energy are released for every mole of SO2(g) reacted. 196 kJ of energy are absorbed for every mole of SO2(g) reacted. 98 kJ of energy are released for every mole of SO2(g) reacted. 98 kJ of energy are absorbed for every mole of SO2(g) reacted. (Total 1 mark) 3. Which statement is correct for an endothermic reaction? A. The products are more stable than the reactants and H is positive. B. The products are less stable than the reactants and H is negative. C. The reactants are more stable than the products and H is positive. D. The reactants are less stable than the products and H is negative. (Total 1 mark) 4. The following equation shows the formation of magnesium oxide from magnesium metal. 2Mg(s) + O2(g)2MgO(s) HӨ = –1204kJ Which statement is correct for this reaction? A. 1204 kJ of energy are released for every mol of magnesium reacted. B. 602 kJ of energy are absorbed for every mol of magnesium oxide formed. C. 602 kJ of energy are released for every mol of oxygen gas reacted. D. 1204 kJ of energy are released for every two mol of magnesium oxide formed. (Total 1 mark) 140 5. Which of the quantities in the enthalpy level diagram below is (are) affected by the use of a catalyst? Enthalpy I II III Time A. B. C. D. I only III only I and II only II and III only (Total 1 mark) 6. When the solids Ba(OH)2 and NH4SCN are mixed, a solution is produced and the temperature drops. Ba(OH)2(s) + 2NH4SCN(s) → Ba(SCN)2(aq) + 2NH3(g) + 2H2O(l) Which statement about the energetics of this reaction is correct? A. B. C. D. The reaction is endothermic and H is negative. The reaction is endothermic and H is positive. The reaction is exothermic and H is negative. The reaction is exothermic and H is positive. (Total 1 mark) 7. Which statements about exothermic reactions are correct? I. They have negative H values. II. The products have a lower enthalpy than the reactants. III. The products are more energetically stable than the reactants. A. B. C. D. I and II only I and III only II and III only I, II and III (Total 1 mark) 8. Consider the specific heat capacity of the following metals. Specific heat capacity / J kg–1 K–1 Cu 385 Ag 234 Au 130 Pt 134 Which metal will show the greatest temperature increase if 50 J of heat is supplied to a 0.001 kg sample of each metal at the same initial temperature? A. Cu B. Ag C. Au D. Pt (Total 1 mark) Metal 141 9. When 40 joules of heat are added to a sample of solid H2O at –16.0°C the temperature increases to –8.0°C. What is the mass of the solid H2O sample? [Specific heat capacity of H2O(s) = 2.0 J g–1K–1] A. 2.5 g B. 5.0 g C. 10 g D. 160 g (Total 1 mark) 10. The mass m (in g) of a substance of specific heat capacity c (in J g–1 K–1 ) increases by t°C. What is the heat change in J? A. mct B. mc(t + 273) C. D. mct 1000 mc (t 273) 1000 (Total 1 mark) 11. Calculate the enthalpy change, H4 for the reaction C + 2H2 + 1 2 O2 CH3OH H4 using Hess’s Law and the following information. CH3OH + 1 12 O2 CO2 + 2H2O H1 = 676 kJ mol1 C + O2 CO2 H2 = 394 kJ mol1 H2 + 1 2 H3 = 242 kJ mol1 O2 H2O .............................................................................................................................................. .............................................................................................................................................. .............................................................................................................................................. .............................................................................................................................................. .............................................................................................................................................. .............................................................................................................................................. .............................................................................................................................................. .............................................................................................................................................. (Total 4 marks) 142 12. The data below is from an experiment used to measure the enthalpy change for the combustion of 1 mole of sucrose (common table sugar), C12H22O11(s). The time-temperature data was taken from a data-logging software program. The sugar was burned and used to heat water whose temperature is shown below. Mass of sample of sucrose, m = 0.100 g Mass of water heated by the combustion of sucrose, m = 300.0 g Heat capacity of water c= 4.18 J g–1 C–1 (a) Calculate ΔT, for the water, surrounding the chamber in the calorimeter. ..................................................................................................................................... ..................................................................................................................................... (1) (b) Determine the amount, in moles, of sucrose. ..................................................................................................................................... ..................................................................................................................................... ..................................................................................................................................... ..................................................................................................................................... (1) 143 (c) (i) Calculate the enthalpy change for the combustion of 1 mole of sucrose. ........................................................................................................................... ........................................................................................................................... (1) (ii) The true value of the enthalpy of combustion of sucrose is -5644 kJ mol-1. Calculate the percentage experimental error based on the data used in this experiment. ........................................................................................................................... ........................................................................................................................... (1) (d) A hypothesis is suggested that TNT, 2-methyl-1,3,5-trinitrobenzene, is a powerful explosive because it has: • a large enthalpy of combustion • a high reaction rate • a large volume of gas generated upon combustion Use your answer in part (c)(i) and the following data to evaluate this hypothesis: Equation for combustion Relative rate of combustion Sucrose C12H22O11(s) + 12O2(g) 12CO2(g) + 11H2O(g) Low TNT 2C7H5N3O6(s) 7CO(g) + 7C(s) + 5H2O(g) + 3N2(g) High Enthalpy of combustion / kJ mol–1 3406 ..................................................................................................................................... ..................................................................................................................................... ..................................................................................................................................... ..................................................................................................................................... ..................................................................................................................................... ..................................................................................................................................... (3) (Total 7 marks) 144 13. The standard enthalpy change of formation values of two oxides of phosphorus are: P4(s) + 3O2(g) P4O6(s) HӨf= –1600 kJ mol–1 P4(s) + 5O2(g) P4O10(s) HӨf= –3000 kJ mol–1 What is the enthalpy change, in kJ mol–1, for the reaction below? P4O6(s) + 2O2(g) P4O10(s) A. B. C. D. +4600 +1400 –1400 –4600 (Total 1 mark) 14. The equations and enthalpy changes for two reactions used in the manufacture of sulfuric acid are: S(s) O2(g) SO2(g) HӨ = –300 kJ 2SO2(g) + O2(g) 2SO3(g) HӨ = –200 kJ What is the enthalpy change, in kJ, for the reaction below? 2S(s) + 3O2(g) 2SO3(g) A. B. C. D. –100 –400 –500 –800 (Total 1 mark) 15. For the reaction 2H2(g) + O2(g) 2H2O(g) the bond enthalpies (in kJ mol–1) are H–H x O=O y O–H z Which calculation will give the value, in kJ mol–1, of HӨ for the reaction? A. 2x + y –2z B. 4z – 2x – y C. 2x + y – 4z D. 2z –2x – y (Total 1 mark) 145 16. But–1–ene gas, burns in oxygen to produce carbon dioxide and water vapour according to the following equation. C4H8 + 6O2 4CO2 + 4H2O (a) Use the data below to calculate the value of HӨ for the combustion of but-1-ene. Bond Average bond enthalpy / kJ mol– CC C=C CH O=O C=O O–H 348 612 412 496 743 463 1 .................................................................................................................................... .................................................................................................................................... .................................................................................................................................... .................................................................................................................................... .................................................................................................................................... .................................................................................................................................... (3) (b) State and explain whether the reaction above is endothermic or exothermic. .................................................................................................................................... .................................................................................................................................... (1) (Total 4 marks) 146 Topic 6: Kinetics 6.1 Rates of Reaction 6.1.1 Define the term rate of reaction. 6.1.2 Describe suitable experimental procedures for measuring rates of reactions. 6.1.3 Analyze data from rate experiments. 6.2 Collision Theory 6.2.1 Describe the kinetic theory in terms of the movement of particles whose average energy is proportional to temperature in kelvins. 6.2.2 Define the term activation energy, Ea. 6.2.3 Describe the collision theory. 6.2.4 Predict and explain, using the collision theory, the qualitatively effects of particle size, temperature, concentration and pressure on the rate of a reaction. 6.2.5 Sketch and explain qualitatively the Maxwell-Boltzmann energy distribution curve for a fixed amount of gas at different temperatures and its consequences for changes in reaction rate. 6.2.6 Describe the effect of a catalyst on a chemical reaction. 6.2.7 Sketch and explain Maxwell-Boltzmann curves for reactions with and without catalysts. Topic 7: Equilibrium 7.1 Dynamic Equilibrium 7.1.1 Outline the characteristics of chemical and physical systems in a state of equilibrium. 7.2 The position of Equilibrium 7.2.1 Deduce the equilibrium constant expression from the equation for a homogeneous reaction. 7.2.2 Deduce the extent of a reaction from the magnitude of the equilibrium constant. 7.2.3 Apply Le Chatelier’s principle to predict the qualitative effects of changes of temperature, pressure, and concentration on the position of equilibrium and on the value of the equilibrium constant. 7.2.4 State and explain the effect of a catalyst on an equilibrium reaction. 7.2.5 Apply the concepts of kinetics and equilibrium to industrial processes. 147 Topic 6 & 7 NOTES 148 Rates of Reaction 1. Define “rate of reaction” 2. All rates are measured as a change in something over a change in time. What are the units of rate of reaction? 3. Below is a graph of concentration vs time for reaction: molarity time a. Does the graph show the change in reactants or products? How do you know? b. The slope of the graph at any time gives the instantaneous rate. Is the rate (slope) constant? c. When is the magnitude of the slope the greatest? d. What is the sign of the slope? 4. Given the following reaction: A + 2B → C The rate of reaction in terms of A is -0.5 mol dm-3 sec-1. a. What is the rate in terms of B? b. What is the rate in terms of C? 5. Molarity 1.0 Time Shown is the graph for the reaction A + 2B → C. The line on the curve represents the concentration of A over time. Draw lines representing the concentrations of B and C over time and label each on the graph. 6. We do not have a device that directly measures molarity. Name 4 things (that are directly related to concentration) you can measure in the lab that would allow you to determine a rate? 7. What methods could you use to measure the rate of reaction for the following: a. Calcium carbonate slowly dissolves in acid to make carbon dioxide gas and a dissolved salt. b. A sample of mineral ore slowly dissolves in water to make aqueous ions. c. A piece of zinc is placed in a solution of copper (II) sulfate. (You did this in your types of reactions lab.) 149 Kinetics Worksheet 1. Kinetics is the study of the ___________ of a reaction. 2. In order for particles to react, they must ___________. 3. In order for collisions to be successful and produce products, the particles must possess what two characteristics? a. b. 4. Name four ways to speed up a reaction. a. b. c. d. 5. Name two reasons increasing the temperature speeds up a reaction. Put a star next to the primary reason. 6. Draw a potential energy diagram for an exothermic reaction. Label the reactants, products and activated complex. PE Reaction coordinate 7. On the diagram above, draw a dotted line representing the effect of a catalyst. 8. On the diagram above, draw a double line representing the effect of an inhibitor. 9. Explain why the addition of a catalyst increases the rate of reaction. 10. Draw a curve of kinetic energy vs. number of particles for a substance at a given temperature (MaxwellBoltzmann curve). Draw a line to represent the minimum threshold energy for a reaction to occur. Draw a dotted line that shows how the distribution would change if the temperature were raised. Draw a double line to show the effect of a catalyst on the distribution. Number of Particles kinetic energy 150 Discovering Equilibrium Activity To learn a little bit about equilibrium observe the demonstration done by your teacher which involves two 1-L beakers. One is completely filled with water and represents reactants (A). The second beaker represents products (B). We will represent the reaction with the following equation. Notice the double headed arrow. This means the reaction is “reversible”. That means not only can the reactants make products, but the products can turn around and make reactants. A ↔ B Colorless red Now let’s start the reaction. (Remember both the forward and reverse reaction occur at the same time!) To simulate the reaction we will use two 50 mL beakers. Dip each little beaker into our products and reactants. You should be able to fill the little beaker in the reactants (water) and only partially fill the beaker in the products (red water). Now start the reaction—turn the reactants into products by pouring the contents of the little beaker into the products. At the same time pour the products into the reactants. 1 L beaker 50 mL beaker A ↔ B What volume of A did you put in the little beaker? ______ mL What volume of B did you put in the other little beaker? _____ mL Is this system at equilibrium? How do you know? If we continue dipping the little beakers simultaneously into the big beakers and emptying them, what will eventually happen to the water levels in the big beakers? Write your hypothesis below: 151 Let’s test our hypothesis. Continue the reaction by dipping the little beakers into the big beakers and empty their contents into the opposite side. Repeat this for 20 times. Stop the reaction and look again. What volume of A did you put in the last little beaker? ______ mL What volume of B did you put in the other little beaker? ______ mL Is the system at equilibrium? How do you know? Equilibrium exists when the rate of the forward reaction equals the rate of the reverse reaction. This means you are making as many reactants as you are using up. The result is that the amounts of products and reactants do not change. It does NOT mean that you have equal amounts of products and reactants. Summarize your understanding of equilibrium by answering the following true or false questions. _____ 1. At equilibrium the reaction stops. _____ 2. To be an equilibrium system, the reaction must be reversible. _____ 3. At equilibrium the forward and reverse reactions occur at the same rate. _____ 4. At equilibrium the concentrations of products and reactants become equal. _____ 5. At equilibrium the concentrations of products and reactants become constant. Which of the following systems are examples of equilibrium systems where both forward and reverse reactions occur at the same rate? Place a checkmark in front of any equilibrium system. _____ 6. Grass growing in a field. _____ 7. Water evaporating and condensing in a sealed jar. _____ 8. Wood burning in a fire. _____ 9. A saturated solution in contact with undissolved solute. 152 Equilibrium Constants The law of mass action represents a mathematical relationship in equilibrium systems. It is expressed by the ratio of product concentrations divided by reactant concentrations. Note that only aqueous and gaseous substances are included in the expression. To write the equilibrium constant expression, know as K or Keq, look at the following example: 2A(g) + 3B(s) ↔ C(g) + 3D(g) [C] · [D]3 K= [A]2 Notice that products are in the numerator and reactants are in the denominator. The brackets represent molarities. Coefficients become exponents. The reactant, B, was not included because it is a solid and its molarity is a constant. 1. Write the equilibrium constant expressions for the following equations. N2O4(g) ↔ 2NO2(g) 3. N2(g) + 3H2(g) ↔ 2NH3(g) 2. PbCl2(s) ↔ Pb2+(aq) + 2Cl-(aq) 4. 3A(g) + B(g) ↔ 2C(s) + D(g) What does the value of K represent? 5. If the value of K > 1, there is a greater concentration of (products/reactants) at equilibrium. 6. If the value of K< 1, there is a greater concentration of (products/reactants) at equilibrium. Solve the following problems using this equilibrium system: N2(g) + 3H2(g) ↔ 2NH3(g) 7. If the concentration of N2 is 0.50 M, the concentration of H2 is 0.40 M and the concentration of NH3 is 1.0 M, what is the value of K? (You must use the equilibrium constant expression from problem number 3 above.) 8. At 500 K, the equilibrium constant, K, equals 1.0 × 10-3 M-2. If [H2] = 0.10 M and [NH3] = 0.10 M, what is the concentration of N2 at this equilibrium position? 153 Le Chatelier's Principle 1. State Le Chatelier's Principle 2. Define equilibrium 3. State three different types of stress you can apply to an equilibrium system. 4. An increase in temperature always favors the endothermic/exothermic reaction. 5. An increase in concentration of a substance in an equilibrium system will shift the reaction towards/away from that substance. 6. An increase in pressure will shift an equilibrium system towards the side with the fewest/most moles of gaseous substances. 7. Given the following hypothetical chemical reaction: 2A(g) + B(g) 2C(g) + heat a. b. c. d. e. f. 8. Is this reaction endothermic or exothermic? If we increase the concentration of A, which direction will the reaction shift? (towards the reactants or towards the products) If we increase the concentration of C, which direction will the reaction shift? If we increase the pressure of the system, which direction will it shift? If we increase the temperature of the system, which direction will it shift? If we increase the temperature of the system, what will happen to the value of K? Given the following hypothetical chemical reaction: heat + A(g) + 3B(g) 5C(g) a. Is this reaction endothermic or exothermic? b. If we decrease the concentration of A, which direction will the reaction shift? (towards the reactants or towards the products) c. If we decrease the concentration of C, which direction will the reaction shift? d. If we decrease the pressure of the system, which direction will it shift? e. If we decrease the temperature of the system, which direction will it shift? f. If we decrease the temperature of the system, what will happen to the value of K? 154 Lab: Equilibrium and LeChatelier’s Principle LeChatelier’s principle states that if an equilibrium system is subjected to stress, the system will react to relieve the stress. To relieve a stress, the system can do one of two things: form more products using up reactants, or reverse the reaction and form more reactants using up products. In this experiment you will observe several equilibrium systems. Then by putting different stresses on the systems, you will observe how equilibrium systems react to a stress. Pre-Lab Questions 1. Define saturated, unsaturated and supersaturated. Which of these represents an equilibrium system? 2. What types of stress can be placed on an equilibrium system? 3. Write dissociation reactions for the following: a. AgNO3 b. NaOH c. HCl d. KSCN (SCN- is a polyatomic ion known as thiocyanate) e. Na3PO4 Materials Equilibrium solutions (1) Saturated NaCl solution (colorless) (2) Acid/Base Indicator solution (green) (3) FeSCN2+ solution (reddish-orange) (4) Cobalt complex solution (purple) **HCl, 12M (extremely corrosive solution… use only in fume hood) HCl, 0.1 M AgNO3, 0.1 M NaOH, 0.1 M Test tubes KSCN solid Beaker Na3PO4 solid Hot plate 155 Procedure 1. Equilibrium in a saturated solution. NaCl(s) ↔ Na+(aq) + Cl-(aq) Obtain a small quantity of saturated NaCl solution in a test tube. Note its appearance. In the fume hood, add a few drops of 12 M HCl. (HCl is a source of both H+ ions and Cl- ions but only the Cl- are part of the equilibrium system.) Note any changes in the test tube. Empty the contents in the waste container in the fume hood and wash the test tube. 2. Acid/Base indicator equilibrium. HX(aq) ↔ H+(aq) + X-(aq) Yellow blue An indicator is a weak acid, HX, where X is typically a complicated organic ion. Its formula is not important for this concept. What is important is that HX is one color and X- is another color. This allows us to detect shifts in the equilibrium. Place a small quantity of acid/base indicator solution in a test tube. Note its appearance. Add 5 drops of 0.1 M HCl solution and stir. Note any color change. To the same test tube add 0.1 M NaOH until no further color change occurs. Again, note the color. (Adding OH- ions causes the H+ concentration to decrease as the ions combine to form water molecules.) See if you can add the right amount of acid and/or base to this test tube to cause the solution to be green in color after it is stirred. Rinse and clean the test tube. You may dispose of the solution by washing it down the drain with water. 3. Complex ion equilibrium. Fe3+(aq) + SCN-(aq) ↔ FeSCN2+(aq) colorless colorless red-brown Obtain a small quantity of FeSCN2+ solution and divide it into three small test tubes. The first test tube will be your reference. Note its color. To the second test tube add 2-3 crystals of KSCN. To the third test tube add a few crystals of Na3PO4. (PO43- ions have the ability to form complex ions with Fe3+, which has the same effect as removing Fe3+ ions from solution.) Note the change. Rinse and clean the test tubes. You may dispose of the solutions by washing them down the drain with water. 4. Cobalt complex equilibrium. Co(H2O)62+ + 4 Cl- → CoCl42- + 6 H2O ΔH = + 50 kJ/mol Pink blue Obtain a small quantity of Cobalt complex equilibrium solution and divide it into four small test tubes. The first test tube will be your reference. Note the color. To the second test tube add a few drops of H2O. (Water is not the solvent in this solution it is a product!) Note the color. To the third test tube add a few drops of 12 M HCl (do this in the fume hood). Note the color. To the fourth test tube add a few drops of 0.1 M AgNO3 solution. Note the color. Obtain a sealed pipet containing the cobalt complex equilibrium solution. Note its color. Place the pipet into an ice water bath for a few minutes. Shake the pipet to stir the contents. Note the color. Then place the pipet into a hot water bath (~60ºC) while gently shaking the pipet. Note the color. 156 Data and Analysis 1. Saturated Solution NaCl(s) ↔ Na+(aq) + Cl-(aq) Initial observation of system: Change observed when HCl added: Stress: increasing the concentration of ClDirection of Shift to relieve the stress: (right or left) 2. Acid/Base indicator HX(aq) ↔ H+(aq) + X-(aq) Yellow blue Initial observation of system: Change observed when HCl added: Stress : Direction of shift: Change observed when NaOH added: Stress: Direction of shift: 3. Complex ion equilibrium. Fe3+(aq) + SCN-(aq) ↔ FeSCN2+(aq) colorless colorless red-brown Initial observation of system: Change observed when KSCN added: Stress: Direction of shift: Change observed when Na3PO4 added: Stress: Direction of shift: 157 4. Cobalt complex equilibrium. Co(H2O)62+ + 4 Cl- → CoCl42- + 6 H2O Pink blue ΔH = + 50 kJ/mol Initial observation of system: Change observed when H2O added: Stress: Direction of shift: Change observed when HCl added: Stress: Direction of shift: Change observed when AgNO3 added: Stress: Direction of shift: Change observed when placed in ice water: Stress: Direction of shift: Change observed when placed in hot water: Stress: Direction of shift: Questions 1. When the concentration of a reactant was added to an equilibrium system, in what direction did the system shift? How was the effect different if a product was added? 2. In the cobalt complex equilibrium, neither Ag+ nor NO3- ions were part of the equilibrium equation yet its addition changed the system. Explain what happened when the silver nitrate solution was added and why it affected the equilibrium system. 3. Look at the cobalt complex equilibrium equation. Is this reaction endothermic or exothermic? Which direction did the equilibrium shift when the temperature was increased? Explain why increasing the temperature of the system shifted the system in the direction it did. 158 Study Guide: Topics 6 & 7 Kinetics and Equilibrium 1. Which quantities in the enthalpy level diagram are altered by the use of a catalyst? Enthalpy I II III Time A. B. C. D. I and II only I and III only II and III only I, II and III (Total 1 mark) 2. Which statement is correct for a collision between reactant particles leading to a reaction? A. Colliding particles must have different energy. B. All reactant particles must have the same energy. C. Colliding particles must have a kinetic energy higher than the activation energy. D. Colliding particles must have the same velocity. (Total 1 mark) 3. Which changes increase the rate of a chemical reaction? I. Increase in the concentration of an aqueous solution II. Increase in particle size of the same mass of a solid reactant III. Increase in the temperature of the reaction mixture A. I and II only B. I and III only C. II and III only D. I, II and III (Total 1 mark) 159 4. The sequence of diagrams represents the system as time passes for a gas phase reaction in which reactant X is converted to product Y. Diagram 1 t = 7 seconds Diagram 2 t = 5 minutes Diagram t = 10 minutes Diagram 4 t = 5 days Time, t X= Y= Which statement is correct? A. At t = 5 days the rate of the forward reaction is greater than the rate of the backward reaction. B. At t = 7 seconds the reaction has reached completion. C. At t = 10 minutes the system has reached a state of equilibrium. D. At t = 5 days the rate of the forward reaction is less than the rate of the backward reaction. (Total 1 mark) 5. Which statement concerning a chemical reaction at equilibrium is not correct? A. The concentrations of reactants and products remain constant. B. Equilibrium can be approached from both directions. C. The rate of the forward reaction equals the rate of the reverse reaction. D. All reaction stops. (Total 1 mark) 6. What is the equilibrium constant expression, Kc, for the reaction below? N2(g) + 2O2(g) 2NO2(g) NO2 A. Kc = N 2 O 2 2NO 2 B. Kc = 3N 2 O 2 C. Kc = D. Kc = NO2 2 N 2 O 2 2 NO2 2 N 2 O 2 2 (Total 1 mark) 160 7. Sulfur dioxide and oxygen react to form sulfur trioxide according to the equilibrium. 2SO2(g) + O2(g) 2SO3(g) How is the amount of SO2 and the value of the equilibrium constant for the reaction affected by an increase in pressure? A. The amount of SO3 and the value of the equilibrium constant both increase. B. The amount of SO3 and the value of the equilibrium constant both decrease. C. The amount of SO3 increases but the value of the equilibrium constant decreases. D. The amount of SO3 increases but the value of the equilibrium constant does not change. (Total 1 mark) 8. What will happen to the position of equilibrium and the value of the equilibrium constant when the temperature is increased in the following reaction? Br2(g) + Cl2(g) 2BrCl(g) ∆H = +14 kJ A. B. C. D. Position of equilibrium Shifts towards the reactants Shifts towards the reactants Shifts towards the products Shifts towards the products Value of equilibrium constant Decreases Increases Decreases Increases (Total 1 mark) 9. In the reaction below N2(g) + 3H2(g) 2NH3(g) ∆H = –92 kJ which of the following changes will increase the amount of ammonia at equilibrium? I. Increasing the pressure II. Increasing the temperature III. Adding a catalyst A. B. C. D. I only II only I and II only II and III only (Total 1 mark) 10. The manufacture of sulfur trioxide can be represented by the equation below. 2SO2(g) + O2(g) 2SO3(g) ∆Hο = –197 kJ mol–1 What happens when a catalyst is added to an equilibrium mixture from this reaction? A. The rate of the forward reaction increases and that of the reverse reaction decreases. B. The rates of both forward and reverse reactions increase. C. The value of ∆Hο increases. D. The yield of sulfur trioxide increases. (Total 1 mark) 161 11. The reaction between two substances A and B A+BC+D has the following rate expression: rate = k [A] Draw the graphical representation of: [A] against time [B] against time [B] [A] time time rate against [A] rate against [B] rate rate [A] [B] (Total 3 marks) 12. When excess lumps of magnesium carbonate are added to dilute hydrochloric acid the following reaction takes place. MgCO3(s) + 2HCl(aq) → MgCl2(aq) + CO2(g) + H2O(l) (a) Outline two ways in which the rate of this reaction could be studied. In each case sketch a graph to show how the value of the chosen variable would change with time. …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… 162 (4) (b) State and explain three ways in which the rate of this reaction could be increased. …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… (6) (c) State and explain whether the total volume of carbon dioxide gas produced would increase, decrease or stay the same if (i) more lumps of magnesium carbonate were used. …………………………………………………………………………………… …………………………………………………………………………………… …………………………………………………………………………………… (2) (ii) the experiments were carried out at a higher temperature. …………………………………………………………………………………… …………………………………………………………………………………… …………………………………………………………………………………… (2) (Total 14 marks) 163 13. (i) Draw a graph that shows the distribution of molecular energies in a sample of a gas at two different temperatures, T1 and T2, such that T2 is greater than T1. (2) (ii) Define the term activation energy. (1) (iii) State and explain the effect of a catalyst on the rate of an endothermic reaction. (2) (Total 5 marks) 14. The table below gives information about the percentage yield of ammonia obtained in the Haber process under different conditions. Temperature/°C Pressure/ atmosphere 10 100 200 300 400 600 (a) 200 50.7 81.7 89.1 89.9 94.6 95.4 300 14.7 52.5 66.7 71.1 79.7 84.2 400 3.9 25.2 38.8 47.1 55.4 65.2 500 1.2 10.6 18.3 24.4 31.9 42.3 From the table, identify which combination of temperature and pressure gives the highest yield of ammonia. ………………………………………………………………………………………. (1) (b) The equation for the main reaction in the Haber process is N2(g) + 3H2(g) ∆H is negative 2NH3(g) Use this information to state and explain the effect on the yield of ammonia of increasing (i) pressure: …………………………….……………………………………….. ……………………………………………………………..…………………. ……………………………………………………………………………….. ……………………………………………………………………………….. (2) 164 (ii) temperature: …………………………………………………………………. …………………………………………………………………………….…. ……………………………………………………………………………….. ……………………………………………………………………………….. ……………………………………………………………………………….. (2) (c) In practice, typical conditions used in the Haber process are a temperature of 500 °C and a pressure of 200 atmospheres. Explain why these conditions are used rather than those that give the highest yield. ………………………………………………………………………………………. ………………………………………………………………………………………. ………………………………………………………………………………………. ………………………………………………………………………………………. (2) (d) Write the equilibrium constant expression, Kc, for the production of ammonia. ………………………………………………………………………………………. ………………………………………………………………………………………. (1) (Total 8 marks) 15. (a) An industrial gas mixture is produced by the catalytic reforming of methane using steam. CH4(g) + H2O(g) H = +206 kJ CO(g) + 3H2(g) By circling the appropriate letter(s) below, identify the change(s) that would shift the position of equilibrium to the right. A increasing the temperature B decreasing the temperature C increasing the pressure D adding a catalyst E decreasing the pressure F increasing the concentration of H2 (2) 165 (b) The following graph represents the change of concentration of reactant and product during a reaction. 0.7 0.6 0.5 Product 0.4 [reactant] or [product] / 0.3 mol dm –3 Reactant 0.2 0.1 0.0 0 (i) 10 20 30 Time / s 40 50 60 Calculate the average rate of reaction over the first 15 s, stating the units. ............................................................................................................................ ............................................................................................................................ ............................................................................................................................ ............................................................................................................................ ............................................................................................................................ ............................................................................................................................ (3) (ii) After 19 s the concentrations of the reactant and product do not change. State what this indicates about the reaction. ............................................................................................................................ ............................................................................................................................ (1) (Total 6 marks) 166 Topic 8: Acids and Bases 8.1 Theories of Acids and Bases 8.1.1 Define acids and bases according to the Bronsted-Lowery and Lewis theories. 8.1.2 Deduce whether or not a species could act as a Bronsted-Lowery and/or a Lewis acid or base. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Bronsted-Lowery base (or acid). 8.2 Properties of Acids and Bases 8.2.1 Outline the characteristic properties of acids and bases in aqueous solution. 8.3 Strong and Weak acids and Bases 8.3.1 Distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity. 8.3.2 State whether a given acid or base is strong or weak. 8.3.3 Distinguish between strong and weak acids and bases, and determine the relative strengths of acids and bases using experimental data. 8.4 The pH Scale 8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale. 8.4.2 Identify which of two or more aqueous solutions is more acidic or alkaline using pH values. 8.4.3 State that each change of one pH unit represents a 10- fold change in the hydrogen ion concentration [H+(aq)]. 8.4.4 Deduce changes in [H+(aq)] when the pH of a solution changes by more than one pH unit. 167 Topic 8 NOTES 168 Acids and Bases Worksheet 1. Properties of Acids a. Acids have a pH value less than ___________. b. Acids have a ________ taste. c. Acids react with metal to form ____________ gas. d. Acids react with carbonate compounds to form ______________ gas. e. When a strong acid reacts with a strong base it forms ____________ and ____________. f. Litmus turns ___________ in the presence of an acid solution. 2. Properties of Bases a. Bases have a pH value greater than ____________. b. Bases have a ____________ taste. c. Bases feel ______________ . d. Litmus turns __________ in the presence of a basic solution. e. Phenolphthalein turns _____________ in the presence of a basic solution. 3. Definitions a. Define an acid and base according to Bronsted-Lowry. 4. b. How does a conjugate acid/base pair differ from one another? c. Water undergoes "autoionization." What does this mean? Write a chemical equation showing this process. d. What is a hydronium ion? Write its formula. e. Define "electrolyte". Give an example of a nonelectrolyte, strong electrolyte and a weak electrolyte. f. Name the three common strong acids. g. What is the difference between a strong and weak acid? The pH scale. a. Write the mathematical equation for pH. b. What is the pH of a solution with a [H+] concentration of 10-5? c. What is the pH of a solution with a [H+] concentration of 10-11? d. What is the [H+] of a solution with a pH of 8? e. A solution has a pH of 2. Is it basic, acidic, or neutral? f. A solution has a pH of 12. Is it basic, acidic, or neutral? g. A solution has a pH of 3. A second solution has a pH of 5. Which is more acidic? h. In part "g" how many times more acidic (in terms of [H+] is the most acidic solution compared to the least acidic solution? 169 Brønsted-Lowry Acids and Bases 1. Label the Brønsted-Lowry acids and bases in the following reactions and show the direction of proton transfer. H+ H2O + Cl- ↔ OH- + HCl Acid base base acid a. NH3 + H2O ↔ NH4+ + OHb. H2O + H2O ↔ H3O+ + OHc. HCO3- + HSO4- ↔ H2CO3 + SO42d. H3PO4 + F- ↔ H2PO4- + HF e. NO2- + H2O ↔ HNO2 + OH- 5. Fill in the table of the following acid base conjugate pairs. Acid H2SO4 Base Acid HPO42- NO3- Base NH3 SO32- H2O HCO3- H2O 6. Name the following acids. a. HCl b. H2SO4 c. HBr d. e. f. HNO3 HNO2 H3PO4 7. Write formulas for the following acids. a. hydroiodic acid b. sulfurous acid c. carbonic acid d. e. f. acetic acid hydrosulfuric acid chloric acid 170 pH and pOH 1. Define pH. 2. Define pOH. 3. pH + pOH = __________ 4. If the pH of a solution is 3, the solution is 5. If the pOH of a solution is 5, the solution is 6. Kw is called the "ionization constant" for water. What is the equation for Kw and what is its numerical value at 25C? 7. In the symbol [H+], what do the brackets represent? 8. If the [H+] of a solution is 10-8, what is the [OH-]? 9. If the pOH is 9, what is the [OH-]? 10. If the pH of a solution is 3, what is the pOH? the [H+]? the [OH-]? 11. In a strong acid solution, the [H+] concentration is (greater than, equal to, less than) the original concentration of the acid. 12. In a weak acid solution, the [H+] concentration is (greater than, equal to, less than) the original concentration of the acid. 13. If 1.0% of a 0.10 M solution of weak acid HF ionizes, what is the concentration of the [H+]? What is the pH? 14. 5% of a 0.20 M weak base ionizes. What it the [OH-] of the solution? What is the pOH and the pH? 15. What is the pH of a 10-2 M solution of HCl? What is the pOH? (acidic/basic/neutral)__ . 171 (acidic/basic/neutral)__ . Lewis Acids and Bases 1. Define an acid according to Lewis theory. 2. Define a base according to Lewis theory. 3. Define dative bond. 4. Draw Lewis dot diagrams for NH3 and BH3. 5. Using the diagrams you drew in #4, use an arrow to show the electron pair that will be used to make the dative bond. (Draw the arrow from the base to the acid to show the electrons involved.) 6. Identify the Lewis acids and bases in the reactants of the following reactions: a. Ag+ + 2NH3 → [Ag(NH3)2]+ b. B(OH)3 + H2O → B(OH)4- + H+ c. Fe3+ + 6CN- → [Fe(CN)6]3- 7. Boron trifluoride, BF3, and ammonia, NH3, react to produce a product. A dative bond is formed between the boron atom on BF3 and the nitrogen atom on NH3. Write the equation for this reaction, using Lewis electron-dot formulas. Label the Lewis acid and the Lewis base. Determine how many grams of product are formed when 10.0g of each reactant are placed in a reaction vessel, assuming that the reaction goes to completion. 172 pH and pOH pH is an indication of how acidic or basic a solution is. pH = -log[H+]. A logarithm (log) is just the exponent to which ten is raised to obtain a given value. For example if [H+] = 10-5, then the log of 10-5 = -5 and the pH would be 5. pOH = -log[OH-]. pH + pOH = 14 [H+][OH-] = 10-14 Using the relationships above, fill in the following table. [H+] 10-4 [OH-] pH pOH 8 -4 10 4 -11 10 7 -2 10 1 10-6 Calculate the pH of the following solutions. 1. 0.01 M HCl 2. 0.001 M NaOH 3. 0.10 M HNO3 4. 0.050 M Ca(OH)2 5. 0.01 M HF (assume 1% ionization) 6. 0.0001 M HC2H3O2 (assume 10% ionization) 7. 2.00 M NH4OH (assume 5% ionization) 173 Acidic or basic EXTEND- Titration to Determine the Molarity of Vinegar Procedure Student’s trial run– Put approximately 10 mL of vinegar (RECORD EXACT AMOUNT) and 15 mL of distilled water into an Erlenmeyer flask. Add 3 drops of phenolphthalein indicator. Titrate to the color change with the standardized NaOH solution. Once the solution has changed color, stop the titration. Wait for all of the lab groups to complete their first run. Compare your solution with the class. Remember: Do not let NaOH go below the 50mL mark on the buret Data Collection – Collect data for three additional trials. Record data and observations in the table below. Data REMEMBER UNCERTAINTIES!! Volume of Initial volume of Trial Vinegar (mL) 0.500 M NaOH* (mL) Trial Run 1 2 3 *The NaOH solution is 0.500 ± 0.005 mol dm-3 Qualitative observations made during lab: 174 Final volume of 0.500 M NaOH* (mL) Volume of 0.500 M NaOH* used (mL) Use your data from the activity above to answer the following questions. 1. Calculate the number of moles of NaOH added in each trial. Show a sample calculation and record all results in the table. 2. At the equivalence point in any titration, how do the moles of base compare to the moles of acid? 3. Based on your answer to number 2, record the number of moles of acid present in each trial in the table below. 4. Calculate the molarity of vinegar for all 3 trials. Show a sample calculation and record all results in the table. 5. Report the average molarity of vinegar. Trial Moles of NaOH Moles of vinegar Molarity of vinegar (M or mol/L or mol dm-3) 1 2 3 Average 6. Calculate the % error for your data. (Obtain theoretical value from teacher) 7. Describe at least two possible sources of error. Suggest one way that you could improve your lab techniques to lower your percent error the next time you do this lab. 8. One method of determining molarities during titration is to use the equation M1V1 = M2V2, where M1V1 equals moles of acid and M2V2 equals moles of base. Using your data from Trial 1 solve for the molarity of the vinegar, M1. Compare this value to your results in the table above. 9. Write a balanced chemical equation for the neutralization of sodium hydroxide and vinegar, HC2H3O2. 175 EVALUATE Write balanced equations for the following neutralization reactions: 1. H3PO4 + KOH → 2. HI + Mg(OH)2→ Titrations 1. What is the formula for determining molarity using titration? 2. If we neutralize 15.0 cm3 of 1.00 mol dm-3 HCl with 12.5 cm3 of NaOH, what is the molarity of the NaOH? 3. What volume of a 0.50 M HCl solution do we need to neutralize 35 mL of 0.80 M NaOH? 4. What is the molarity of KOH if we used 50.0 mL of 0.100 mol dm-3 HNO3 to neutralize 25.0 mL of the KOH? 5. What is the purpose of an indicator solution in a titration? 176 Electrolytes Using a conductivity probe or meter, test the conductivity of the following substances in water solution and categorize into one of the following categories (your teacher may do this for you). Make sure you test pure water to see what its conductivity is before you test the solutions. Substance (0.1 M solution) Excellent Conductor Poor Conductor Nonconductor NaCl Sugar (C12H22O11) HCl Vinegar (HC2H3O2) NaOH NH3 Alcohol (C2H5OH) Based on the results above, answer the following questions. Salts, like NaCl, are strong/weak/non electrolytes. (circle one) HCl is a strong/weak acid and is a strong/weak electrolyte. Vinegar is a strong/weak acid and is a strong/weak electrolyte. NaOH is a strong/weak base and is a strong/weak electrolyte. NH3 is a strong/weak base and is a strong/weak electrolyte. Sugar and alcohols are strong/weak/non electrolytes. Sports drinks, like Gatorade, contain electrolytes. What ingredient do these drinks contain that makes them a good source of electrolytes? 177 178 Study Guide: Topic 8 Acids and Bases 1. Which is a conjugate acid-base pair in the following reaction? HNO3 + H2SO4 A. B. C. D. H2NO3+ + HSO4– HNO3 and H2SO4 HNO3 and H2NO3+ HNO3 and HSO4– H2NO3+ and HSO4– (Total 1 mark) 2. Which equation represents an acid-base reaction according to the Lewis theory but not the Brønsted-Lowry theory? A. NH3 + HCl NH4Cl H3O + OH– B. 2H2O C. NaOH + HCl D. CrCl3 + 6NH3 + NaCl + H2O [Cr(NH3)6]3+ + 3Cl– (Total 1 mark) 3. In which reaction is H2PO4–(aq) acting as a Brønsted-Lowry base? A. B. C. D. H2PO4–(aq) + NH3(aq) → HPO42–(aq) + NH4+(aq) H2PO4–(aq) + OH–(aq) → HPO42–(aq) + H2O(l) H2PO4–(aq) + C2H5NH2(aq) → HPO42–(aq) + C2H5NH3+(aq) H2PO4–(aq) + CH3COOH(aq) → H3PO4(aq) + CH3COO–(aq) (Total 1 mark) 4. An aqueous solution of which of the following reacts with magnesium metal? A. Ammonia B. Hydrogen chloride C. Potassium hydroxide D. Sodium hydrogencarbonate (Total 1 mark) 5. Which acids are strong? I. HCl(aq) II. HNO3(aq) III. H2SO4(aq) A. B. C. D. I and II only I and III only II and III only I, II and III (Total 1 mark) 179 6. Solutions of hydrochloric acid (HCl(aq)) and ethanoic acid (CH3COOH(aq)) of the same concentration reacted completely with 5.0 g of calcium carbonate in separate containers. Which statement is correct? A. B. C. D. CH3COOH(aq) reacted slower because it has a lower pH than HCl(aq). A smaller volume of CO2(g) was produced with CH3COOH(aq) than with HCl(aq). A greater volume of CO2(g) was produced with CH3COOH(aq) than with HCl(aq). The same volume of CO2(g) was produced with both CH3COOH(aq) and HCl(aq). (Total 1 mark) 7. When the following 1.0 mol dm–3 solutions are listed in increasing order of pH (lowest first), what is the correct order? A. B. C. D. HNO3 H2 CO3 NH3 Ba(OH)2 NH3 Ba (OH)2 H2 CO3 HNO3 Ba (OH)2 H2 CO3 NH3 HNO3 HNO3 H2 CO3 Ba (OH)2 NH3 (Total 1 mark) 8. Define the terms Brønsted-Lowry acid and Lewis acid. For each type of acid, identify one example other than water and write an equation to illustrate the definition. (Total 5 marks) 9. Identify one example of a strong acid and one example of a weak acid. Outline three different methods to distinguish between equimolar solutions of these acids in the laboratory. State how the results would differ for each acid. (Total 5 marks) 180 10. Vinegar has a pH of approximately 3 and some detergents have a pH of approximately 8. State and explain which of these has the higher concentration of H+ and by what factor. (Total 1 mark) 11. The pH values of solutions of three organic acids of the same concentration were measured. acid X acid Y acid Z (i) pH = 5 pH = 2 pH = 3 Identify which solution is the least acidic. (1) (ii) Deduce how the [H+] values compare in solutions of acids Y and Z. (2) (iii) Arrange the solutions of the three acids in decreasing order of electrical conductivity, starting with the greatest conductivity, giving a reason for your choice. (2) (Total 5 marks) 181 Topic 9: Oxidation and Reduction 9.1 Introduction to Oxidation and Reduction 9.1.1 Define oxidation and reduction in terms of electron loss and gain. 9.1.2 Deduce the oxidation number of an element in a compound. 9.1.3 State the names of compounds using oxidation numbers. 9.1.4 Deduce whether an element undergoes oxidation or reduction in reactions using oxidation numbers. 9.2 Redox Equations 9.2.1 Deduce simple oxidation and reduction half-equations given the species involved in a redox reaction. 9.2.2 Deduce redox equations using half-equations. 9.2.3 Define the terms oxidizing agent and reducing agent. 9.2.4 Identify the oxidizing and reducing agents in redox equations. 9.3 Reactivity 9.3.1 Deduce a reactivity series based on the chemical behavior of a group of oxidizing and reducing agents. 9.3.2 Deduce the feasibility of a redox reaction from a given reactivity series. 9.4 Voltaic Cells 9.4.1 Explain how a redox reaction is used to produce electricity in a voltaic cell. 9.4.2 State that oxidation occurs at the negative electrode (anode) and reduction occurs at the positive electrode (cathode). 9.5 Electrolytic Cells 9.5.1 Describe, using a diagram, the essential components of an electrolytic cell. 9.5.2 State the oxidation occurs at the positive electrode (anode) and reduction occurs at the negative electrode (cathode). 9.5.3 Describe how current is conducted in an electrolytic cell. 9.5.4 Deduce the products of the electrolysis of a molten salt. 182 Topic 9 NOTES 183 Oxidation - Reduction Rules for determining oxidation numbers: 1. Pure elements have an oxidation number = 0 2. The oxidation number of a monatomic ion = charge of the ion 3. Alkali metals in compounds have an oxidation number = +1 4. Alkaline earth metals in compounds have an oxidation number = +2 5. Fluorine in compounds has an oxidation number = -1 6. Oxygen in compounds has an oxidation number = -2 except in peroxides (-1) 7. Hydrogen has an oxidation number = +1 except in metal hydrides (-1) 8. The sum of all oxidation numbers in a compound = 0 9. The sum of all oxidation numbers in a polyatomic ion = charge of the ion Assign oxidation numbers to each element in the following. 1. K 2. RbCl 3. Na2O 4. NH3 5. CO2 6. N2O5 7. N2O 8. N2O3 9. CaCO3 10. KNO3 11. Na2Cr2O7 12. KMnO4 13. H3PO4 14. H2O2 15. FeCl3 16. Li2SO4 Define each of the following terms: 17. Oxidation 18. Reduction 19. Oxidizing Agent 20. Reducing Agent In the following equations, identify the substance oxidized, the substance reduced, the oxidizing agent, and the reducing agent. Ox. 21. CH4 + O2 → CO2 + H2O 22. Zn + AgNO3 → Zn(NO3)2 + Ag 23. Al + Cl2 → AlCl3 24. Na + H2O → H2 + NaOH 184 Red. OA RA Writing and Balancing Half Reactions For each of the reactions below, split the equation into half reactions and balance using the method demonstrated by your teacher. 1. Zn(s) + Al3+(aq) → Al(s) + Zn2+(aq) 2. MnO4-(aq) + Fe(s) → Fe2+(aq) + Mn2+(aq) (occurs in acidic solution) 3. Cr(s) + NO3-(aq) → Cr3+(aq) + NO(g) (occurs in acidic solution) 4. NO2-(aq) + Al(s) → NH3(g) + AlO2-(aq) (occurs in basic solution) 5. CN-(aq) + MnO4-(aq) → CNO-(aq) + MnO2(s) (occurs in basic solution) 185 Activity Series of Metals and Nonmetals Metals have different reactivities. For example, the alkali metals all react with water, most metals react with acid, while a small number react with neither. The metals can be placed in a reactivity series that helps identify whether one metal is capable of replacing another in a single replacement reaction. Reactivity Series of metals: Li, K, Ca, Na, Mg, Al, Zn, Fe, Ni, Sn, Pb, Cu, Hg, Ag, Au Most reactive least reactive Using the series it is possible to determine if aluminum is capable of replacing copper in the compound CuCl2. (Al + CuCl2 → ? ) Since Al is higher in the series it will be capable of replacing the Cu ions in CuCl2. Since Al is higher than Cu we say that it is more easily oxidized and is a better reducing agent than Cu. Predict if the following reactions will occur. For those that do, complete and balance the equation, and label the oxidizing agent. For those that do not, write “no reaction.” 1. Zn + FeCl3 → 2. Ag + Cu(NO3)2 → 3. Na + KCl → 4. Ca + Mg(NO3)2 → 5. Al + MgF2 → 6. Zn + Pb(NO3)2 → The reactivity of the halogens is a matter of periodic trends, fluorine being the most reactive and iodine being the least reactive as you travel down the family. Predict whether or not the following reactions will occur. Complete and balance those that do. 7. Cl2 + NaI → 8. Br2 + CaF2 → 9. F2 + AlCl3 → 10. I2 + KBr → 186 Electrochemical Cells An electrochemical, galvanic, or voltaic cell converts chemical potential energy into electricity. It relies on the difference in activity between the two metals that make up the cathode and anode or the two metal ions in the solutions. The electrons flow as a result of a spontaneous reaction. The spontaneity can be somewhat predicted based on the activity series. Species higher up on the chart will act as the anode while species lower on the chart will act as cathodes. A more detailed prediction of the output voltage can be calculated using Standard Reduction Potentials (which we are not going to cover in this class). Factors that affect the voltage output of a cell are the activity difference in the two metals (emf), concentration, and temperature. A quantitative description is given by the Nernst equation (also something we’re going to skip). At the anode, oxidation takes place, the neutral metal atoms lose electrons which flow through the wire to the light bulb causing it to light up. The positive ions that form as a result of oxidation, dissolve in solution. At the cathode, reduction takes place, the positive ions in solution gain electrons that have flowed from the anode. The ions become neutral atoms and precipitate out of solution onto the cathode. Sometimes an inert but conducting material is used as an electrode when one or both of the species is non-conducting. (For example, if the following is one of the half-reactions in the cell: MnO4- + 5e- + 8H+ Mn2+ + 4H2O Neither MnO4- nor Mn2+ can serve as electrodes since they are not solid so a Platinum electrode could be used) Salt bridge to allow flow of ions so there is no charge build up in either solution. Charge build-up would shift the equilibrium and the reaction would cease e- Light or voltmeter ee- e- Anode: Oxidation, Negative electrode Cathode: Reduction, Positive electrode Electrolyte: allows for charges to move from one electrode to the other. 187 Electrochemical Cells V _ + Given the following redox reaction that takes place in the electrochemical cell above: 3 Cu2+ (aq) + 2 Al (s) → 3 Cu (s) + 2 Al3+ (aq) 1. 2. 3. 4. Which species is being oxidized? Write the equation for the oxidation half reaction. Oxidation takes place at which electrode? What is the charge of this electrode? 5. 6. 7. 8. Which species is being reduced? Write the equation for the reduction half reaction. Reduction takes place at which electrode? What is the charge of this electrode? 9. Electrons in the external circuit flow from which electrode to which electrode? 10. Towards which electrode do positive ions in the salt bridge flow? 11. Electrochemical cells convert _____________ energy into _____________ energy. 11. On the diagram above, label the following: salt bridge, anode, cathode, Cu, Al, Cu2+ solution, Al3+ solution, direction of electron flow in the external circuit, direction of ion flow in the salt bridge 188 Electrolytic Cells Electrolytic cells are the opposite of galvanic cells: They convert electrical energy into chemical potential energy. Electricity must be supplied in order to force a non-spontaneous reaction to occur. The spontaneity can be somewhat predicted based on the activity series. Current flows from the species being oxidized to the species being reduced. Major Differences: Direction of electron flow changes Oxidation still occurs at the anode and reduction still occurs at the cathode. Notice that cathode is now the negative electrode, it attracts cations Notice that the anode is now the positive electrode, it attracts anions Notice that there is a power supply (battery) instead of a light bulb or voltmeter e- Power Source ee- Salt bridge to allow flow of ions so there is no charge build up in either solution. e- Cathode: Reduction occurs here Negative Electrode Cations attracted to anode Anode: Oxidation occurs here Positive Electrode Anions are attracted to cathode Electrolyte: allows for charges to move from one electrode to the other. Examples of electrolysis in action are: making hydrogen and oxygen gas from water o 2H2O 2H2 + O2 o 2H+ + 2e- H2 (Reduction at negative electrode/cathode) o O2- O2 + 4 e- (Oxidation at positive electrode/anode) Making sodium and chlorine from molten salt o NaCl(l) Na + Cl2 o Cathode: Na+ + 1 e- Na o Anode: 2Cl- + 2 e- Cl2 Electroplating precious metals onto less expensive metals 189 Electrolytic Cells + - 1. Molten potassium iodide is placed into the reaction vessel in the above electrolytic cell. What are the two reactants in this reaction? 2. Write the half reaction for the oxidation that will take place. 3. At which electrode will oxidation take place? 4. What is the charge of this electrode? 5. Write the half reaction for the reduction that will take place. 6. At which electrode will reduction take place? 7. What is the charge of this electrode? 8. Label the anode and cathode on the diagram above showing what is produced at each electrode. 9. Write the overall balanced equation for the reaction. 10. Electrolytic cells convert _____________ energy into ______________ energy. 11. State two ways that electrochemical and electrolytic cells are similar. 12. State two ways that electrochemical and electrolytic cells are different. 190 Topic 9 NOTES 191 Study Guide: Topic 9 Oxidation and Reduction 1. Which are examples of reduction? A. B. C. D. I. Fe3+ becomes Fe2+ II. Cl– becomes Cl2 III. CrO3 becomes Cr3+ I and II only I and III only II and III only I, II and III (Total 1 mark) 2. In which change does nitrogen undergo oxidation? A. NO2 N2O4 B. NO3– NO2 C. N2O5 NO3– D. NH3 N2 (Total 1 mark) 3. What are the oxidation numbers of the elements in sulfuric acid, H2SO4? Hydrogen Sulfur Oxygen A. +1 +6 –2 B. +1 +4 –2 C. +2 +1 +4 D. +2 +6 –8 (Total 1 mark) 4. Consider the following reaction: H2SO3(aq) + Sn4+(aq) + H2O(l) → Sn2+(aq) + HSO4–(aq) + 3H+(aq) Which statement is correct? A. H2SO3 is the reducing agent because it undergoes reduction. B. H2SO3 is the reducing agent because it undergoes oxidation. C. Sn4+ is the oxidizing agent because it undergoes oxidation. D. Sn4+ is the reducing agent because it undergoes oxidation. (Total 1 mark) 192 5. Which statement is correct about an oxidizing agent in a chemical reaction? A. It reacts with oxygen. B. It reacts with H+ ions. C. It loses electrons. D. It undergoes reduction. (Total 1 mark) 6. A voltaic cell is made from magnesium and iron half-cells. Magnesium is a more reactive metal than iron. Which statement is correct when the cell produces electricity? A. Electrons are lost from magnesium atoms. B. The concentration of Fe2+ ions increases. C. Electrons flow from the iron half-cell to the magnesium half-cell. D. Negative ions flow through the salt bridge from the magnesium half-cell to the iron halfcell. (Total 1 mark) 7. What process occurs at the cathode in a voltaic cell and at the anode in an electrolytic cell? A. Cathode of voltaic cell Oxidation Anode of Electrolytic cell Reduction B. Oxidation Oxidation C. Reduction Oxidation D. Reduction Reduction (Total 1 mark) 8. Which statement is correct for the electrolysis of a molten salt? A. Positive ions move toward the positive electrode. B. A gas is produced at the negative electrode. C. Only electrons move in the electrolyte. D. Both positive and negative ions move toward electrodes. (Total 1 mark) 9. Iron in food, in the form of Fe3+, reacts with ascorbic acid (vitamin C), C6H8O6, to form dehydroascorbic acid, C6H6O6. (i) Write an ionic half-equation to show the conversion of ascorbic acid to dehydroascorbic acid in aqueous solution. ......................................................................................................................... ......................................................................................................................... (1) (ii) In the other ionic half-equation Fe3+ is converted to Fe2+. Deduce the overall equation for the reaction between C6H8O6 and Fe3+. ......................................................................................................................... ......................................................................................................................... (1) (Total 2 marks) 193 10. Deduce the change in oxidation number of chromium in the below reaction. State with a reason whether the chromium has been oxidized or reduced. Cr2O72 + 14H+ + 6Fe2+ 2Cr3+ + 6Fe3+ + 7H2O .............................................................................................................................................. .............................................................................................................................................. .............................................................................................................................................. .............................................................................................................................................. (Total 2 marks) 11. A part of the reactivity series of metals, in order of decreasing reactivity, is shown below. magnesium zinc iron lead copper silver If a piece of copper metal were placed in separate solutions of silver nitrate and zinc nitrate (i) determine which solution would undergo reaction. …………………………………………………………………………………………… (1) (ii) identify the type of chemical change taking place in the copper and write the half-equation for this change. …………………………………………………………………………………………… …………………………………………………………………………………………… (2) (iii) state, giving a reason, what visible change would take place in the solutions. …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… …………………………………………………………………………………………… (2) (Total 5 marks) 194 12. A current is passed through molten sodium chloride. Identify the substance formed at each electrode and write an equation to represent the formation of each substance. Determine the mole ratio in which the substances are formed. …………………………………………………………………………………………………… …………………………………………………………………………………………………… …………………………………………………………………………………………………… …………………………………………………………………………………………………… …………………………………………………………………………………………………… …………………………………………………………………………………………………… (Total 5 marks) 195 196 197 198 199