Chemistry for Medicine MPRM-0103 Chapter 1 PDF

CHEMISTRY FOR MEDICIME MPRM-0103 CHAPTER 1 Measurements and Properties of matter Dr. Amal Al Sabahi  1.3 Measurement Session Learning outcomes: SLO#1: To distinguish between a qualitative and quantitative observation (a measurement). SLO#2: T...

CHEMISTRY FOR MEDICIME MPRM-0103 CHAPTER 1 Measurements and Properties of matter Dr. Amal Al Sabahi  1.3 Measurement Session Learning outcomes: SLO#1: To distinguish between a qualitative and quantitative observation (a measurement). SLO#2: To choose the appropriate device for a certain measurement. SLO#3: To match the physical quantities with their SI system units. SLO #4: to determine the density of a substance. SLO #5: To perform conversions among different temperature units.  1.3 Measurement Two types of observation collected: ▪ A qualitative observation: e.g. a white precipitate formed, gas evolved or color disappear. ▪ A quantitative observation (a measurement): e.g ethanol boils at 78C, the mass of CaCO3 is 0.10 g…etc. Class Activity: Which of the following is a measurement? - His body temperature is 39 C. - She has a high blood pressure. - The tips of your fingers are greenish. - You are obese, your BMI is 32.5.  1.3 Measurement What is a measurement? A measurement tells us about a property of something, and gives a number to that property. Different devices or instruments are used to measure the properties of a substance. Mass Current Volume Time Length thermometer  1.3 Measurement In any measured quantity both the number and unit should be written The unit indicates the scale of the measurement e.g. 25.01 mL 0.99 g/cm3 121 s 298 K 33 m 1.87 g  1.3 Measurement SI Units an international agreement set up a system of units called the International System of units (SI system) which is based on the metric system. Base quantity Units Symbol Mass kilogram kg Length meter m Time second s Temperature kelvin K Amount of substance mole mol Electric Current Ampere A Luminous Intensity candela cd  1.3 Measurement prefix symbol meaning Exponential femto f 0.000000000000001 10-15 pico p 0.000000000001 10-12 nano n 0.000000001 10-9 micro  0.000001 10-6 milli m 0.001 10-3 centi c 0.01 10-2 deci d 0.1 10-1 kilo k 1,000 103 Mega M 1,000,000 106 Giga G 1,000,000,000 109 Tera T 1,000,000,000,000 1012  1.3 Measurement Mass: describes the quantity of matter in an object /substance. Weighing balance is used to measure the mass. Common units: kg , g , mg , g , ng. Length: Common units: km , m , mm , g , ng…etc Examples: - Majority of lengths in chemistry ranges 10-9 to 10-12 m - Proteins diameter are commonly (1 to 3 nm). - Angstrom (Å)= 10-10 m is used to measure bond length.  1.3 Measurement Volume: amount of space occupied by the sample. Measured with burette , pipette , volumetric flask , graduated cylinder. Syringe  1.3 Measurement Volume: Units of volume are SI derived units: examples: m3 , cm3 , dm3 , mm3 …etc 1 cm3 = (110-2 m)3 = 110-6 m3 1 dm3 = (110-1 m)3 = 110-3 m3 1 cm3 = 110-3 dm3 Other common (not SI): L , mL , L … 1 L = 1 dm3 1 mL = 1 cm3  1.3 Measurement Density: Density: the mass of a substance occupied in unit volume. 𝒎𝒂𝒔𝒔 𝒎 𝒅= = 𝒗𝒐𝒍𝒖𝒎𝒆 𝑽 Density units are SI derived units, examples: kg/m3 , g/cm3 , g/mL ,.. Density doesn’t depend on the quantity of mass present. What does that mean? It is affected by temperature variation. How? Sometimes used as an identification tag for a substance. How? Gas density  liquid density  solid density. Why?  1.3 Measurement Density: Temperature ℃ Density of water g/cm3 5 1.00000 20 0.99829 40 0.99225 60 0.98313 80 0.97763 80 0.97160 100 0.95808 Artificial Intelligence with solid fill  1.3 Measurement Density: Example: a piece of platinum metal with a density of 21.5 g/cm3 has a volume of 4.49 cm3. What is its mass? Solution: 𝑚 𝑑= 𝑉 𝑚 =𝑑×𝑉 𝑔 3 𝑚 = 21.5 × 4.49 𝑐𝑚 𝑐𝑚3 𝑚 = 96.5 𝑔 Artificial Intelligence with solid fill  1.3 Measurement Density: Exercise : An unknown colorless substance Density liquid has a mass of 27.9 g and a (g/cm3) volume of 35.4 mL at 20 oC. Ethanol 0.789 Identify the substance (among Benzene 0.880 the 3 in the given table) water 0.998 Artificial Intelligence outline  1.3 Measurement Density: Table 1.3 Densities of some substances at 25 ℃ Substance Density g/cm3 Substance Density g/cm3 Hydrogen (g) 0.0000899 Aluminum (s) 2.70 Graphene 0.00016 Diamond (s) 3.5 aerogel Air (g) 0.001 Iron (s) 7.9 Ethanol (l) 0.79 Lead (s) 11.3 Water (l) 0.97763 Mercury (l) 13.6 Graphite (s) 0.97160 Gold (s) 19.3 Artificial Intelligence with solid fill Table salt (s) 0.95808 Osmium (s) 22.6  1.3 Measurement Density: Class exercise : Calculate the volume of 25.0 g of each of the following substances at 25C. i. Hydrogen gas (d= 0.0000899 g/cm3) ii. Water (d= 1.00 g/cm3) iii. Iron (d= 7.87 g/cm3) Artificial Intelligence outline  1.3 Measurement Temperature Scales Temperature is measured with a thermometer Three systems for measuring temperature are given: ▪ The Celsius scale (C) ▪ The Kelvin scale (K) (SI base unit). ▪ The Fahrenheit scale (F).  1.3 Measurement Temperature Scales Units conversion 180 oF difference = 100 oC difference 180 𝑜 𝐹 9 𝑜𝐹  = 100 𝑜 𝐶 5 𝑜𝐶 0 oC = 32 oF = 273 K  1.3 Measurement Temperature Scales, Units conversion 𝟗 𝒐𝑭 𝒐𝑭 ℃ 𝒕𝒐 ℉ 𝑻𝑭 = × 𝑻𝒄 + 𝟑𝟐 𝟓 𝒐𝑪 𝟓 𝒐𝑪 ℉ 𝒕𝒐 ℃ 𝑻𝒄 = 𝑻𝑭 − 𝟑𝟐 𝒐 𝑭 × 𝒐 𝟗 𝑭 𝑻𝑲 = 𝑻𝑪 + 𝟐𝟕𝟑. 𝟏𝟓 ℃ 𝒕𝒐 𝐾 TK /TF conversion TK  Tc  TF TF  Tc  TK  1.3 Measurement Temperature Scales Example Convert 327.5 C (the melting point of lead) to: i) kelvin ii) degrees Fahrenheit Solution: i) 𝑇𝐾 = 𝑇𝐶 + 273.15 = 327.5 + 273.15 = 600.65 K = 600.7 K 9 𝑜𝐹 ii) 𝑇𝐹 = 𝑜 × 𝑇𝑐 + 32 ℉ 5 𝐶 9 𝑜𝐹 = 𝑜 × 327.5 ℃ + 32 ℉ 5 𝐶 = 1113.2 ℉  1.3 Measurement Temperature Scales Example: The boiling point of N2 is 77 K, convert it to F? Solution: 𝑇𝐶 = 𝑇𝐾 − 273.15 = 77 − 273.15 = −𝟏𝟗𝟔 ℃ 9℉ 𝑇𝐹 = × 𝑇𝑐 + 32 ℉ 5℃ 9 𝑜𝐹 = × −196 ℃ + 32 ℉ = −𝟑𝟐𝟏 ℉ 5 𝑜𝐶 Example: A person has a body temperature of 102.5 F. What is his temperature in kelvin scale? A) 39.2 B) 36.5 C) 312.4 D) 308.5  1.3 Measurement Assigned textbook questions: 14e: 1.5 , 1.6 , 1.8 , 1.10 , 1.12 , 1.13 , 1.14 , 1.59 , 1.62 , 1.63 , 1.64 , 1.67 , 1.69 , 1.70 , 1.77 13e: 1.19 , 1.20 , 1.22 , 1.24 , 1.26 , 1.27 , 1.28 , 1.55 , 1.58 , 1.59 , 1.60 , 1.63 , 1.65 , 1.66 , 1.73  1.4 Handling Numbers Session Learning outcomes: SLO #1: To Apply significant figures rules in calculations. SLO #2: To employ scientific notation when handling numbers. SLO #3: To discriminate between precision and accuracy. SLO #4: To differentiate between systematic errors and random errors.  1.4 Handling Numbers Significant Figures ▪ Are the meaningful digits in a measured or calculated quantities. ▪ Significant figures are the certain digits + one uncertain digit in the measurements. e.g. In measuring the volume of a solution using a burette 24.19 mL 3 certain digits (24.1) 1 uncertain digit (9) 4 significant figures (24.19)  1.4 Handling Numbers Significant Figures e.g. In measuring the mass of a substance using an electronic balance 10.4978 g 5 certain digits (10.497) 1 uncertain digit (8) 6 significant figures (10.4978)  1.4 Handling Numbers Significant Figures e.g. In measuring the temperature using a thermometer. 2 certain digits (36) 36.0 C 1 uncertain digit (0) 3 significant figures (36.0) Example: How many significant figures are in the following measurements? ▪ d= 2.2 g/cm3 2 significant figures ▪ m= 112.7 g 4 significant figures ▪ Cppm= 23.956 ppm 5 significant figures  1.4 Handling Numbers Significant Figures, Guidelines: ▪ Nonzero digits are all significant figures. (e.g. 12.67) ▪ Zeros between nonzero are also significant figures. (e.g. 3006 , 20.1 , 888.02 and 707) ▪ Zeros to the left of the first nonzero digit are not significant. (e.g. 0.013, 0.0000453 , 0.999) ▪ All zeros written to the right of nonzero with a decimal is significant. (e.g. 3.30 , 768.000 , 0.4310) ▪ Zeros to the right of a nonzero if no decimal is present are not significant. (e.g 320 , 69600) ▪ Exact numbers have infinite number of significant figs obtained by definition, counting, etc  1.4 Handling Numbers Significant Figures: Examples a. 0.00250 g → 3 significant figures b. 9000 A → 1 significant figures c. 12 cm → 2 significant figures d. 1098 s → 4 s.f e. 2.001 x 10-3 → 4 s.f f. 1.0100 x 10-5 → 5 s.f g. 1000. m → 4 s.f (because of the decimal point). h. 22.04030 → 7 s.f  1.4 Handling Numbers Significant Figures: Rules for calculations: Multiplication and division: The number of significant figures in the result is the same as that in the quantity with the smallest number of significant figures. Example: 2.8 x 4.5039 = 12.61092 → 13 (2 s.f) (5 s.f) (2 s.f) The product should have only two significant figures since 2.8 has two significant figures (apply rounding). Example: 223.51 𝑔 𝑑𝑒𝑛𝑠𝑖𝑡𝑦 = = 𝟕. 𝟖𝟕 𝑔/𝑚𝐿 𝟐𝟖. 𝟒 𝑚𝐿  1.4 Handling Numbers Significant Figures: Rules for calculations: Addition and Subtraction: The answer cannot have more digits to right of the decimal place than the original numbers added/subtracted. Example: 10.217 mg 3 decimal places + 1.29 mg 2 decimal places + 0.9 mg 1 decimal place 12.407 mg 1 decimal place (round to 12.4) Example: 9.89 x 103 − 1.1100 x 103 9.89 x 103 2 decimal places − 1.1120 x 103 4 decimal places 8.77(8) x 103 2 decimal places (round it)  1.4 Handling Numbers Significant Figures: Rules for calculations: Addition and Subtraction: Example: find the answer to the following operation and write it to the correct number of significant figures: 3.58 x 10-2 + 7.8 x 10-3 + 4.917 x 10-1 3.58 x 10-2 2 decimal places + 0.78 x 10-2 2 decimal places + 49.17 x 10-2 2 decimal places 53.53 x 10-2 2 decimal places Unify the powers and units before starting the calculation  1.4 Handling Numbers Writing scientific notation When working with very large or very small numbers, it is more practical to use scientific notation. In the form: N.NN  10n Example: 0.000000250 ➔ (scientific notation) ➔ 2.50 x10-7 57000000000 ➔(scientific notation) ➔ 5.7 x1010 When writing a number in scientific notation, the number of significant figures should not be changed  1.4 Handling Numbers Mixed arithmetic operations: 𝟐𝟖.𝟕𝟗 𝟔.𝟕𝟖 Example: − + 𝟏𝟏. 𝟖𝟗𝟕 Note: carry out 𝟑.𝟐𝟒 𝟐.𝟐 3 s.f 2 s.f 5 s.f extra digits, round = 8.88(58) − 3.0(8) + 11.897 off in the last step 2 d.p 1 d.p 3 d.p = 17.7028 = 17.7 (1 decimal place) Example: 𝟕. 𝟑𝟓 × 𝟏𝟑. 𝟔𝟎  (𝟐𝟖. 𝟕𝟑 − 𝟐𝟖. 𝟐𝟑) step 1: 28.73 − 28.23 0.50 2 s.f Step 2: 7.35  13.60  0.50 = 50. or 5.010 (2 s.f)  1.4 Handling Numbers Mixed arithmetic operations: Exercise: carry out the following arithmetic operations to the correct number of significant figures i) 1.0455 + 0.98  1.387 ii) 0.00565 +19.821 – 8.36 iii) 4.65 g + 5.111 g + 4.9 g iv) 3.754  103 + 4.12  102 + 15  103 v) 11 cm  413 mm  0.56 dm  1.4 Handling Numbers Accuracy and precision: Accuracy: determines how closely is a measurement to the true value of the quantity measured. Example: A patient body temperature =36.5C no. Mazin Hamed Qais 1 36.5 37.2 32.6 2 36.4 40.6 32.9 3 36.7 38.6 32.8 4 36.6 36.0 33.1 avg 36.5(5) 38.1 32.9 accurate inaccurate inaccurate  1.4 Handling Numbers Accuracy and precision: precision: refers to how closely measurements of the same quantity agree with each other. Example: a patient body temperature =36.5C no. Mazin Hamed Qais 1 36.5 37.2 32.6 2 36.4 40.6 32.9 3 36.7 38.6 32.8 4 36.6 36.0 33.1 precise imprecise precise  1.4 Handling Numbers Accuracy and precision: H2O(l) Asma Farah Dana 40.00 mL 38.87 mL 44.72 mL 40.03 mL 38.85 mL 49.35 mL 40.01 mL 38.86 mL 31.19 mL 40.02 mL inaccurate inaccurate accurate Poor accuracy Poor accuracy Good accuracy Precise imprecise Precise Good precision poor precision Good precision  1.4 Handling Numbers Accuracy and precision: 2 1 4 3  1.4 Handling Numbers Accuracy and precision: Measurement is associated with error; no measurement is 100% precise and accurate. 2 Types of errors: Systematic error: occurs in predictable manner, leading to measured values constantly off the true value. Reasons: imperfect instrument setup or method of observation.  1.4 Handling Numbers Accuracy and precision: 2 Types of errors: Random error: are not predictable, leading that measured values that vary greatly from the true value. Reasons: fluctuations in the instrumental measurement or incorrect interpretation of an instrument’s reading.  1.4 Handling Numbers Accuracy and precision: Asma Farah Dana H2O(l) 40.00 mL 38.87 mL 44.72 mL 40.03 mL 38.85 mL 49.35 mL 40.01 mL 38.86 mL 31.19 mL 40.02 mL inaccurate inaccurate accurate Poor accuracy Poor accuracy Good accuracy Precise imprecise Precise Good precision poor precision Good precision Systematic error Random error No error  1.4 Handling Numbers Accuracy and precision: Practice Question: 1) The volume of a liquid is 26.0 mL. A student measured the volume and found it to be 26.2 mL, 26.1 mL, 25.9 mL, and 26.3 mL in the 1st, 2nd, 3rd, and 4th trial, respectively. Which statement is true for his measurements? a. They are neither precise nor accurate. b. They have poor accuracy. c. They have good precision. d. They have poor precision.  1.4 Handling Numbers Accuracy and precision: Practice Question: 2) The volume of a liquid is 20.5 mL. Which of the following sets of measurement represents the value with good accuracy? a. 18.6 mL, 17.8 mL, 19.6 mL, 17.2 mL b. 19.2 mL, 19.3 mL, 18.8 mL, 18.6 mL c. 18.9 mL, 19.0 mL, 19.2 mL, 18.8 mL d. 20.2 mL, 20.5 mL, 20.3 mL, 20.1 mL  1.4 Handling Numbers Assigned textbook questions: 14e: 1.17 , 1.18 , 1.20 , 1.21 , 1.24 , 1.25 , 1.26 , 1.83 , 1.105 13e: 1.31 , 1.32 , 1.34 , 1.35 , 1.38 , 1.39 , 1.40 , 1.79 , 1.105  1.5 Dimensional Analysis Session Learning outcomes: SLO #1: to recognize the relationship between units that expresses the same physical quantity SLO #2: To manipulate conversion factors in dimensional analysis problems SLO #3: To solve problems utilizing dimensional analysis.  1.5 Dimensional Analysis It is used to convert between units and in solving problems. by following these steps: 1. Find the appropriate equation/s that relate the units. two units. e.g. 1cm3 = 1 mL , 1 in= 2.54 cm , 1 kg= 103 g , etc 2. derive the appropriate conversion factor by looking at the direction of the required change. 1 𝑖𝑛 2.54 𝑐𝑚 e.g. or 2.54 𝑐𝑚 1 𝑖𝑛 3. Multiply the quantity to the converted unit factor to obtain the quantity with the desired units.  1.5 Dimensional Analysis Example: A roll of Aluminum foil has a mass of 1.07 kg. What is its mass in pounds (1 lb = 453.6 g)? Solution: 103 𝑔 1 𝑙𝑏 1.07 𝑘𝑔 × × = 2.36 𝑙𝑏 1 𝑘𝑔 453.6 𝑔 Example: The volume of a solution is 2.56  106 mm3 what is the volume in L? Solution: (1 𝑑𝑚) 3 1𝐿 6 3 2.56 × 10 𝑚𝑚 × 2 3 × 3 = 2.56 𝐿 (10 𝑚𝑚) 1 𝑑𝑚  1.5 Dimensional Analysis Example: the length of a leaf measured is 11.05 mm what is its length in inch (1 in = 2.54 cm Solution: 1 𝑐𝑚 1 𝑖𝑛 11.05 𝑚𝑚 × × = 0.4350 𝑖𝑛 10 𝑚𝑚 2.54 𝑐𝑚 Example: Ammonia gas under certain conditions has a density of 0.625 g/L. Calculate the density in g/cm3? Solution: 𝑔 106 𝜇𝑔 1𝐿 1 𝑚𝐿 0.625 × × 3 × 𝐿 1𝑔 10 𝑚𝐿 1 𝑐𝑚3 = 0.625 × 103 𝜇𝑔/𝑐𝑚3  1.5 Dimensional Analysis Example: a sheet of aluminum foil has a total area of 1.000 ft2 and mass of 3.636 g. What is the thickness of the foil in millimeter (density of Al = 2.699 g/cm3) (1 ft = 30.48 cm) Solution: 𝑚 3.636 𝑔 𝑉= = = 1.347 𝑐𝑚 3 𝑑 2.699 𝑔/𝑐𝑚3 𝑉 1.347 𝑐𝑚3 (1.000 𝑓𝑡)2 10 𝑚𝑚 𝑙= = × × = 𝐴 1.000 𝑓𝑡 2 (30.48 𝑐𝑚)2 1 𝑐𝑚 1.450 × 102 𝑚𝑚 3.636 𝑔 (1.000 𝑓𝑡) 2 1 𝑐𝑚 3 10 𝑚𝑚 Or.. × × × 1.000 𝑓𝑡 2 (30.48 𝑐𝑚) 2 2.699 𝑔 1 𝑐𝑚 = 1.450 × 102 𝑚𝑚  1.5 Dimensional Analysis Assigned textbook questions: 14e: 1.27 , 1.28, 1.29 , 1.33 , 1.36 , 1.38 , 1.39 , 1.40 , 1.82 , 1.94. 13e: 1.41 , 1.42 , 1.43 , 1.47 , 1.50 , 1.52 , 1.53 , 1.54 , 1.78 , 1.90. Unit equivalences between different systems English – SI equivalences length 1 in (inch) = 2.54 cm (centimeter) 1 m (meter) = 1.094 yd (yard) 1 mi (mile) = 1.609 km (kilometer) mass 1 oz (ounce) = 28.35 g (gram) 1 kg ( kilogram) = 2.205 Ib (pound) 1 Ib (pound) = 453.6 g (gram) volume 1 in3 (cubic inch) = 16.39 cm3 (cubic centimeter) Back with solid fill 1 ft3 (cubic foot) = 28.32 L (liter)