Chapter 1 2024 Chemistry PDF

1  Why study Chemistry ? 1. Water/soil testing 2. Atmospheric science 3. Climate change 1. Materials science 1. Human and Vet. Medicine 2. Solar energy conversion 2. Agriculture 3. Information...

1  Why study Chemistry ? 1. Water/soil testing 2. Atmospheric science 3. Climate change 1. Materials science 1. Human and Vet. Medicine 2. Solar energy conversion 2. Agriculture 3. Information technology Environmental 3. Pharmaceuticals Science Physics Biology Chemistry Geology Engineering 1. Mining 1. Manufacturing 2. Mineral recovery 2. Power plants 3. Oil and gas industry 3. Metallurgy 2 Chapter 1: Particles of Matter My expectations: I consider the first parts of Chapter 1 to be review material, and I won’t be going over it. You are expected to be able to: → Distinguish between pure substances and mixtures (Section 1.2) → Distinguish physical properties from chemical properties and physical processes from Chemical processes (1.5) → Recognize physical states of matter and changes in state 3 (1.6)  The Scientific Method OBSERVATION the careful noting and recording of natural phenomena HYPOTHESIS EXPERIMENT THEORY a tentative explanation a general explanation of a test of a hypothesis of a single or small natural phenomena or theory number of natural phenomena LAW a generally observed natural phenomenon Creative Commons Attribution-Share Alike 4  1.8 Six SI Base Units - used in many base calculations/unit definitions Measurement Abbrev. Unit Length m meter Temperature K kelvin Time s second Amount of a substance mol mole Mass kg kilogram Eg : Newton (N; force) :: kg·(m/s2) == kg·m·s-2 : Joule (J; energy) :: kg·(m2/s2) == kg·m2·s−2 (or N·m == kg·m2·s-2) 5  Prefixes for SI Units Scientific Name Symbol Value Notation Giga G 1,000,000,000 109 Mega M 1,000,000 106 Kilo k 1,000 103 Hecto h 100 102 Deka da 10 101 -----------------------------------Unit name-----------------------------100 Deci d 0.1 10-1 Centi c 0.01 10-2 Milli m 0.001 10-3 Micro μ 0.000,001 10-6 Nano n 0.000,000,001 10-9 Pico p 0.000,000,000,001 10-12 6 Derived Units Definition : Combinations of the base quantities of length, mass, time, temperature, etc… 𝑙𝑒𝑛𝑔𝑡ℎ (𝑚) Examples : velocity = 𝑡𝑖𝑚𝑒 (𝑠) volume = length x length x length (m3) ::: cm3 [ 1 mL = 1 cm3 ] [ 1 L = 1000 cm3 = 1 dm3 ≠ 1 m3 (= 1000 L) ] 𝑚𝑎𝑠𝑠 (𝑔) 𝑔 density = ::: 𝑣𝑜𝑙𝑢𝑚𝑒 (𝑚3 ) (𝑐𝑚3 ) 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑐𝑜𝑚𝑝𝑜𝑛𝑒𝑛𝑡 % composition = x 100 % 𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 7 Ex 1-1. Calculating Density A man receives a platinum ring from his fiancée. Before the wedding, he notices that the ring feels a little light for its size and decides to determine its density. He places the ring on a balance and finds that it has a mass of 3.15 grams. He then finds that the ring displaces 0.233 cm3 of water. Is the ring made of platinum? Density of some common substances (20 C) g/cm3 Ice 0.917 (0 C) Water 1.00 ( at 4C) Aluminum 2.70 Iron 7.87 Copper 8.96 Lead 11.3 Mercury 13.5 Gold 19.3 Platinum 21.5 Silver 10.5 Tin 7.37 8 Ex 1-1. Calculating Density (aside) (Note: The volume of irregularly shaped objects is often measured by the displacement of water. To use this method, the object is placed in water and the change in volume of the water is measured. This increase in the total volume represents the volume of water displaced by the object and is equal to the volume of the object.) 9 Creative Commons Attribution-Share Alike 4.0 1.9 Dimensional analysis Definition : Conversion of measurements and quantities from one unit to another Conversion factor General formula: desired unit given unit ´ = desired unit given unit Example : Regina is 263.2 km from Saskatoon How many miles ? Conversion factor : 1 km = 0.6214 mi 0.6214 𝑚𝑖 263.2 km x = 163.55 mi 1 𝑘𝑚 10 Ex 1-2 Unit conversion Show that 1 m3 = 1000 L, not 1 L 11 SI vs US system (Imperial) The US is one of 3 countries that have not officially adopted SI units as their standard units of measure (Liberia, Myanmar) Canada and UK actually use a mixed system (mostly SI) The is no expectation that you know the conversion factors between US units and SI units Must be able to convert within a unit ie : cm – m – km – nm 12 1.10 Evaluating and Expressing Experimental Results How can we tell how reliable a number is that is reported ? Since we have sources of errors, how can we show this reliability when talking about scientific measurements ? By reporting numbers with decimal places, number of units etc. we can convey the reliability = Significant figures 13 1.10 Significant Figures All measurements contain uncertainty. Depends on instruments used to make measurement A digit that must be estimated is called uncertain (last recorded digit). ± 0.01 g ± 0.0001 g 14 Exact numbers These are numbers that have no uncertainty 1. Counting of discrete objects 2. Defined quantities (eg 1 foot = 12 inches) 3. Integral parts of equations radius = diameter/2 Exact numbers have an infinite number of significant figures Significant Figures: the rules! 1. All nonzero digits are significant 2. Interior zeroes (between two non-zero digits) are significant 3. Leading zeroes (to the left of the first non-zero digit) are not significant. 16  Sig. Figs: cont. 4. Trailing zeroes (located at the end of the number): ❑ Trailing zeroes after a decimal point are always significant ❑ Trailing zeroes before a decimal point (and after a non- zero number) are always significant ❑ Trailing zeroes before an implied decimal point are ambiguous – avoid by using scientific notation. 17  Ex 1-3 Counting Sig. Figs. How many significant figures are in the following numbers? 1) 0.04550 g 2) 100 lb 3) 1625 m 4) 101.05 mL 5) 350.0 g 18 Carrying Significant Figures in Calculations When multiplying or dividing measurements, the result has the same number of significant figures as the measurement with the fewest number of significant figures When adding or subtracting measurements, the result has the same number of decimal places as the measurement with the fewest number of decimal places 19 Significant Figures in Mathematical Operations (cont.)1) ◼ * Multiplication/Division: Number of sig. figs. in the result = Number in the least precise measurement used in the calculation 6.38  2.0 = 12.76 16.84 = 6.6299 2.54 20 Significant Figures in Mathematical Operations (cont. 2) * Addition/Subtraction: Number of significant figures in the result depends on the number of decimal places in the least accurate measurement. 6.8 + 11.934 = 18.734 37.657 − 2.1 = 35.557 = 35.6 (1 dec. place) 21 Significant Figures in Mathematical Operations (cont. 3) * Calculations that involve multiple steps: To determine the number of significant figures, we must first perform individual operations to determine the number of significant figures at each step. 1.23 𝑔 −0.567 𝑔 = 1.924975 g/cm3 0.34442 𝑐𝑚3 The first step involves performing the subtraction in the numerator. 1.23 g – 0.567 g = 0.663 g = 0.66 g 22 Significant Figures in Mathematical Operations (cont. 4) * With only two significant figures in the numerator there can only be two significant figures in the final answer. 23 Types of Error in Measurements Error Type Definition Inherent error. For example, poor instrument Systematic calibration. Can be identified and eliminated ! Random Limitations in the skill of experimenter Accuracy - How close your experimental measurements are to the correct answer. Precision - Refers to the degree of reproducibility of a measured quantity. 24  Accuracy vs Precision –Example 1 Not accurate, not precise == RANDOM Error 25 Accuracy vs Precision –Example 2 Precise but not accurate == SYSTEMATIC error 26 Accuracy vs Precision –Example 3 Precise & accurate -- Note : still some error present 27 Calculations for Precision and Accuracy (Information only) Mean ( X ) An average calculated by summing a set or related values and dividing the sum by the number of values in the set  i ( xi ) x= n  i ( xi ) Represents the sum of all the individual xivalues n Represents the number of values 28  Calculations for Precision and Accuracy (Information only) Standard deviation(s) A measure of the amount of variation, or dispersion, in a set of related values s= (  i xi − x )2 n −1 ( ) 2Represents the sum of all the individual  i xi − x numbers minus the average squared n Represents the number of values 29 29  1.11 Temperature Measure of the average amount of kinetic energy (higher T = greater average kinetic energy) Celsius Scale, C used in most countries (fun aside : Celsius’ original scale had boiling at 0 and freezing at 100) Kelvin Scale, K absolute scale (no –ve numbers) 0 K = absolute zero (temperature at which molecular motion “stops” ) 30 Creative Commons Attribution-Share Alike 4.0 1.11 Temperature (cont.) Converting temperatures: The 3 scales are related to each other: 5 C = (F – 32) 9 the 32 is an exact number K = C + 273.15 the 273.15 is exact 31 Example: Temperature Conversions The lowest temperature measured on the Earth is –128.6 F, recorded at Vostok, Antarctica, in July 1983. What is this temperature in C and in Kelvin? Collect and Organize: We are given the temperature in F and asked to convert it to C and Kelvin. Temperature = −128.6 F 5 C = (F – 32) K = C + 273.15 9 32 Practice: Temperature Conversions (cont. 1) The lowest temperature measured on the Earth is –128.6 F, recorded at Vostok, Antarctica, in July 1983. What is this temperature in C and in Kelvin? Analyze (set up the different equations) 5 C = (F – 32) 9 5 = (-128.6 F – 32) 9 K = C + 273.15 5 = (-128.6 F – 32) + 273.15 9 33 Practice: Temperature Conversions (cont. 2) The lowest temperature measured on the Earth is –128.6 F, recorded at Vostok, Antarctica, in July 1983. What is this temperature in C and in Kelvin? Solve (do the calculations) 5 C = (-128.6 F – 32) 9 = -89.2 C (one decimal place : subtraction) 5 K = (-128.6 F – 32) + 273.15 9 = 183.9 K (one decimal place : subtraction) 34 Practice: Temperature Conversions (cont. 3) lowest temperature measured on the Earth is –128.6 F, The recorded at Vostok, Antarctica, in July 1983. What is this temperature in C and in Kelvin? Think About It (The smell test) Since the Celsius scale has smaller degrees than the Fahrenheit scale, a value that is less negative on the Celsius scale seems reasonable. Adding 273.15 to this produces a positive value on the Kelvin scale, which is reasonable since you cannot have a value less than absolute zero on the Kelvin scale. 35 Solving Chemical Problems Sort the information from the problem Collect and Organize Identify the given quantity and unit, the quantity and unit you want to find, any relationships implied in the problem Given Find Know Design a strategy to solve the problem Analyze Sometimes may want to work backwards Each step involves a conversion factor or equation (Find) Apply the steps in the concept plan Solve Check that units cancel properly Multiply terms across the top and divide by each bottom term Check the answer Think About It Double check the set-up to ensure the unit at the end is the one you wished to find Check to see that the size of the number is reasonable 36  Problem Solving (cont.) Arrange conversion factors to cancel units until desired units obtained May string conversion factors So we do not need to know every relationship, as long as we can find something else the given and desired units are related to 37 Ex 1-4: A 747 airplane is fueled with 173,231 L of jet fuel. If the density of the fuel is 0.768 g/cm3, what is the mass of the fuel in kg? 38