CHEM 101 L1 Matter- Its Properties, Measurements, & Calculations in Chemistry PDF

Matter- Its Properties, Measurements, & Calculations in Chemistry Dr Disah Mpadi Office: Block 1, Room 16 Email: [email protected] 1 Outline ▪ The scientific method ▪ Properties of matter ▪ Classification of matter ▪ SI...

Matter- Its Properties, Measurements, & Calculations in Chemistry Dr Disah Mpadi Office: Block 1, Room 16 Email: [email protected] 1 Outline ▪ The scientific method ▪ Properties of matter ▪ Classification of matter ▪ SI (Metric) Units & uncertainty of scientific measurement ▪ Significant figures & calculations ▪ Dimensional analysis ▪ Temperature & density: Their use in problem solving. 2  Chemistry: Is the science that deals with the composition & properties of matter. Properties: Physicochemical The International Union of Pure & Applied Chemistry (IUPAC) is recognized as the world authority on chemical nomenclature, terminology, standardised methods for measurement, atomic mass, & more. 3  The Scientific Method ▪ The scientific method is the mixture of observation, experimentation, & the devising of laws, hypotheses, & theories. Theory established: Observation: Hypothesis: Theory or Model: Unless later observations Natural or Tentative Amplifies hypothesis & or experiments show experimental explanation gives predictions inadequacies of model Experiments Experiments to designed to test test predictions hypothesis of theory Revise hypothesis: If Modify theory: If experiments show that experiments show that it is inadequate. it is inadequate. 4  Properties of Matter Properties ▪ Matter is anything that occupies space & displays the properties of mass & inertia. Physical property: Chemical property: A A sample of matter sample of matter is displays a property converted to a new ▪ Composition refers to the parts or without changing its type of matter with components of a sample of matter & composition different composition their relative proportions: i.e., H2O is (change in composition) through a chemical composed of 11.19% hydrogen & 88.81% Colour: Copper is reaction or change oxygen. brown, Sulfur is yellow Burning of wood, Brittle: Sulfur Firewood (to C, CO2 ▪ Properties are those qualities or Malleable & & H2O) ductile: Copper attributes that we can use to distinguish Zinc reacts with HCl one sample of matter from others; Physical acid to produce H2 & generally grouped into two broad change(state): ZnCl2 Liquid water, categories: physical & chemical. Freezing (ice), 5 Boiling (vapour)  Classification of Matter All Matter Variable composition? No Yes ▪ Matter is made up of atoms & each different type of atom is the building block of a different Pure Mixture: chemical element (contains one kind of atom). Substance: Contains various types Visibly distinguishable parts? of atoms? No Yes No ▪ Compound is a substance comprising atoms of two or more elements joined together. Yes Element Homogeneous/Solution ▪ Substances refers to elements & compounds: whose composition & properties are uniform Compound Heterogeneous throughout a given sample. Air, Sucrose solution, Fe, S, Au Seawater ▪ Mixture: Substances can vary in composition & properties from one sample to another. H2O, CuSO4 Sand & water; concrete ▪ Molecule: Is the smallest entity having the same A classification scheme for matter. proportions of the constituent atoms as does the compound as a whole. Molecules are covalently bonded, i.e., H2O, CH3CH2OH. 6 Cont’d All Matter Can it be separated by physical means? No Yes Substance Mixture Can it be decomposed Is it uniform throughout? by chemical process? No Yes Element Yes Homogeneous No Compound Heterogeneous A classification scheme for matter. States of Matter: ▪ Solid: Atoms or molecules are in close contact & normally form crystal. ▪ Liquid: Atoms or molecules can flow becoz their loosely held. ▪ Gas: Atoms or molecules can flow becoz their loosely held than in liquid. ▪ Plasma: Ionised gas (ions + electrons). 7 Cont’d Separating mixtures: ▪ Physical process: Chemical cpd retains its identity. ▪ Filtration: Process of separating a solid from a liquid in which is suspended, e.g., sand in water (Heterogeneous mixture). ▪ Distillation: A liquid is condensed from the vapour given off by a boiling solution, e.g., copper(II) sulfate solution, (Homogeneous mixture). There is simple & fractional distillation. ▪ Chromatography: separation of ink components on a paper using water as a solvent. Decomposing Compound: ▪ Chemical process: ▪ A chemical change result in a chemical cpd being decomposed into its constituent elements, e.g., Extraction of iron from iron oxide. ▪ Electrolysis of water. 8  SI (Metric) Units & Uncertainty of Scientific Measurement Multiple Prefix Physical Quantity Unit Symbol 1018 exa (E) 1015 peta (P) Length metre m 1012 Tera (T) Mass kilogram kg 109 giga (G) Time second s 106 mega (M) Temperature kelvin K 103 kilo (k) Amount of substance mole mol 102 hecto (h) 101 deka (da) Electric current ampere A 10-1 deci (d) Luminous intensity candela cd 10-2 centi (c) 10-3 milli (m) ▪ SI (International System of Units) 10-6 micro () base/fundamental quantities & 10-9 nano (n) prefixes. 10-12 pico (p) 10-15 femto (f) 10-18 atto (a) ▪ All other physical quantities have 10-21 zepto (z) units that can be derived from these seven 9 10-24 yocto (y) Cont’d Measurement of Matter: ▪ Measurement-quantitative observation: The product of a number & a unit. Involve some uncertainty. Uncertainty is indicated by using significant figures. ▪ The Unit indicates the standard against which the measured quantity is being compared. ▪ Chemistry is quantitative science: We can measure a property of a substance & compare it to a standard having a known value of the property. The development of science requires careful quantitative measurement. ▪ Nonnumerical information is qualitative, such as the colour red. 10 Cont’d Examples: ▪ Mass: Describes the quantity of matter in an object. In SI, standard mass is 1 kilogram (kg). ▪ Weight is the force of gravity on an object: 𝑊 = 𝑔 𝑥 𝑚 ▪ An object has a fixed mass (𝑚). ▪ Laboratory device for measuring mass: Balance ▪ Time: The second is defined as the duration of 9,192,631,770 cycles of a particular radiation emitted by certain atoms of the element cesium (cesium-133). ▪ In SI, standard of time is the second (𝑠). ▪ Second, minutes, hours, & years can be used in scientific work depending on whether we deal with short intervals or long ones. 11 Cont’d ▪ Derived Units: When measured properties are expressed as combinations of the ▪ Temperature: Temperature fundamental or base quantities. scales: Celsius, Fahrenheit, & Kelvin. ▪ e.g., 𝒗𝒆𝒍𝒐𝒄𝒊𝒕𝒚 = 𝒅𝒊𝒔𝒕𝒂𝒏𝒄𝒆 (𝒖𝒏𝒊𝒕: 𝐦 𝒔−𝟏 ) 𝒕𝒊𝒎𝒆 ▪ Some deirived units have special names: ▪ The SI temperature scale: Kelvin ▪ e.g., The combination kg m-1 s-2 is called the scale. pascal (pressure). ▪ In the lab, temperature is ▪ The combination kg m2 s-2 is called the joule measured in Celsius degrees (oC), (energy). then must be converted to the ▪ Though SI standard unit of volume is m3 (cubic Kelvin scale (unit: K). metre); but commonly used units are cm3 & L (litre). ▪ Temperature conversions: ▪ Non-SI Units: United States still use non- ▪ Kelvin from Celsius: SI. 𝑇(K) = 𝑡(oC) +273.15 ▪ Mass is given in pounds (lb). ▪ Distance or room dimensions in feet. ▪ Fahrenheit from Celsius: ▪ But there is a standard unit conversion between 9 non-SI & SI. 𝑡(oF) = 𝑡(oC) + 32 5 12 Cont’d Uncertainties in Scientific Measurements: ▪ All measurements are subject to error or uncertainty. The uncertainty of a measurement depends on the precision of the measuring device. ▪ Types of errors: ▪ Systematic or Determinate errors: ▪ Occurs in the same direction each time; it is either always high or always low. ▪ Systematic errors influence the accuracy of a measurement. ▪ Accuracy refers to how close a measured value is to the accepted, true, or actual value. ▪ Random or Indeterminate errors: ▪ Means that a measurement has an equal probability of being high or low. Occurs when estimating last digit of a measurement. ▪ Random errors are linked to the precision of measurements. ▪ Precision refers to the degree of reproducibility of a measured quantity; thus, the closeness of agreement (scatter in the data) among several measurements of the same quantity. 13 Cont’d Top-pan Balance Analytical Examples: Balance Three 10.5, 10.4, 10.6 g 10.4978, 10.4979, measurements 10.4977 g Their average 10.5 g 10.4978 g Reproducibility ±0.1 g ±0.0001 g Precision Low or poor High or good ▪ Thus, the results obtained by using top-pan balance have lower or poorer precision than those for the analytical balance. 14 Cont’d Large random errors (Poor technique) & large systematic error Small random errors but a large systematic error Precision vs Accuracy: ▪ Can a set of measurements be precise without being accurate? Can the average of a set of measurements be accurate & the True or individual measurements be accepted imprecise? Explain. value ▪ N.B: In quantitative work, precision is often used as an indication of Large random errors Small random errors accuracy, only if systematic errors & no systematic error & no systematic error are absent. 15  Significant Figures & Calculations ▪ Consider the measurements from a low- Rules to follow to determine the number of precision balance: 10.4, 10.2, & 10.3 g: significant figures in a numerical quantity: ▪ Their average: 10.3 g. 1. All nonzero digits are significant. 2. Zeros are also significant when are between nonzero numbers or at the end of a ▪ Interpreting this average means that the 1st two number to the right of decimal point, but digits (10) are known with certainty & the last with two important exceptions for quantities digit after decimal point (3) is uncertain less than one: (estimated). ▪ Any zeros (a) preceding the decimal point, or (b) following the decimal point & preceding the 1st nonzero digit, are ▪ The mass is only known to the nearest 0.1 g & we NOT significant. can express the result as 10.3 ± 0.1 g, where ± 3. The case of terminal zeros that precede the 0.1 is the uncertainty of the measuring device. decimal point in quantities greater than one is ambiguous. ▪ The number of significant figures (s.f.) in a measured quantity provide a clue about the capabilities of the measuring device & the precision of the measurements. 16 Cont’d ▪ Rule 3: ▪ The digit below has seven significant Examples: figures, quantity greater than 1: ▪ Rules 1 & 2: ▪ The digit below has seven significant ▪ Consider 7500 m: figures, quantity less than 1: ▪ Do we mean measured to nearest metre or nearest 10 metre? ▪ If all zeros are significant, we write, Significant: 7500. m. Fulfill Rule 2. All zeros Not significant: between nonzero numbers. Fulfills Rule 2(a). ▪ The best approach to deal with above Significant: question is to use scientific notation Fulfill Rule 2. Zeros or exponential notation: 0.003006700 at the end of a number ▪ The coefficient establishes the number to the right of decimal of significant figures, & the power of point. ten locates the decimal point: Not significant: Significant: 2 s.f. 3 s.f. 4 s.f. Fulfill Rule 2(b); The zeros only locate Fulfill Rule 1. 7.5 x 103 m 7.50 x 103 m 7.500 x 103 m the decimal point. All nonzero integers. 17 Cont’d Significant figures (s.f.) in numerical calculations: ▪ Precision must neither be gained nor be lost in calculations involving measured quantity. Simple rules involving s.f. in a calculation: ▪ Multiplication & Division: The result of multiplication or division may contain only as many significant figures as the least precisely known quantity or quantity with least number of s.f. in the calculation, e.g.: 14.79 cm x 12.11 cm x 5.05 cm = 904 cm3 (4 s.f.) (4 s.f.) (3 s.f.) (3 s.f) ▪ Addition & Subtraction: The result of addition or subtraction must be expressed with the same number of decimal places as the least precise measurement or the absolute error in the result can be no less than the absolute error in the least precisely known quantity, e.g.: 15.02 g + 9,986.0 + 3.518 g = 10, 004.5 g 18 Cont’d N.B: Exact numbers can be considered to have an infinite or unlimited number of s.f., & are not obtained using measuring devices: ▪ i.e., resulting of counting, 8 molecules; ▪ Or by definition: 1 min = 60 s, 1 inch = 2.54 cm (conversion factors). Rounding off numerical results: ▪ To round off a number to the required number of significant figures: ▪ increase the final digit by one unit if the digit dropped is 5, 6,7, 8, or 9. e.g., To 3 s.f., 16.34 rounds off to 16.3. 15,675 rounds off to 1.58 x 104 ▪ leave the final digit unchanged if the digit dropped is 0, 1, 2, 3, or 4: e.g., To 3 s.f., 16.35 rounds off to 16.4. 19 Cont’d Applying significant figure rules: Addition/Subtraction Applying significant figure rules: ▪ Show the result of the following calculation with the correct number of significant figures. Multiplication/Division 2.06 𝑥 102 + 1.32 𝑥 104 − 1.26 𝑥 103 =? ▪ Show the result of the following calculation with the correct number of significant figures. ▪ Analyse: 0.255 𝑥 0.0035 To determine the correct number of s.f., identify =? the largest quantity & then write the other 2.16 𝑥 10−2 quantities with the same power of ten as appears in the largest quantity. The answer can ▪ Analyse: have no more digits beyond the decimal point than the quantity having the smallest number of By inspecting the three quantities, the least precise such digits. known quantity, 0.0035 has two s.f. & the result should contain only 2 s.f. ▪ Solve: 2.06 𝑥 102 + 1.32 𝑥 104 − 1.26 𝑥 103 ▪ Solve: = 0.0206 𝑥 104 + 1.32 𝑥 104 − 0.126 𝑥 104 0.255 𝑥 0.0035 = 0.041319444 = 0.0206 + 1.32 − 0.126 𝑥104 2.16 𝑥 10−2 = 1.2146 𝑥104 = 𝟎. 𝟎𝟒𝟏 or 𝟒. 𝟏 𝒙 𝟏𝟎−𝟐 = 𝟏. 𝟐𝟏 𝒙𝟏𝟎𝟒 20 Cont’d Good Practice: ▪ In a series of calculations, carry the extra digits through to the final result, then round off. ▪ Round off to the correct number of s.f. only in the final answer. 21  Key point on Significant figures: To obtain the correct number of s.f. in a calculation or a digit, the best way is to express any given digit using scientific notation or exponential notation. e.g., 0.003006700 = 3.006700 x 10-3 (7 s.f.) 22  Dimensional Analysis Unit conversions II: ▪ Sometimes is Unit conversions I: The speed limit on many highways in necessary to convert the US is 55 mi/h. Calculate the A pencil is 7.00 in long. Compute number that would be posted in a given result from its length in centimeters. kilometres per hour. one system of units to another (i.e., non- Analyse: Analyse: SI to SI units). We want to convert from inches to We want to convert from mi/h to centimeters. We must use the unit km/h. We must use the following unit factor: 1 in = 2.54 cm. factors: 1 km = 1000 m; 1 m = 1.094 yd; 1760 yd = 1 mi. ▪ This method of unit Solve: conversion is Solve: referred to as 7.00 in 𝑥 2.54 cm = 7.00 2.54 cm 55 mi 1760 yd 1m 1 km dimensional = 𝟏𝟕. 𝟖 𝐜𝐦 1 in h 𝑥 1 mi 𝑥 𝑥 1.094 yd 1000 m analysis or unit = 𝟖𝟖 𝐤𝐦/𝐡 factor method. 23 Cont’d Good Practice: ▪ When solving chemistry problems, always include the units for the quantities used. 24  Temperature & Density: Their use in problem solving ▪ Density: Is the ratio of mass to volume mass (𝑚) density 𝜌 = volume (𝑉) ▪ Despite SI base units of mass & volume, the most encountered unit of density is g/cm3 or g/mL. ▪ Density is a function of temperature becoz volume varies temperature, whereas mass remains constant. ▪ The state of matter affects the density of a substance: generally, solids are denser than liquids & both denser than gases. ▪ Mass & volume are both extensive properties. ▪ Extensive property: Is dependent on the quantity of matter observed, e.g., mass, volume. ▪ Intensive property: Is independent of the amount of matter observed, e.g., density. ▪ Intensive properties can often be used to identify substances. 25 Cont’d Density can be used as conversion factor to determine the object’s mass or volume: Example: The density of osmium is 22.59 g/cm3. Calculate the mass of a cube of osmium that is 1.25 inches on edge (1 in = 2.54 cm). Note: 𝑉 = 𝑙 3. 𝑚=𝜌𝑥𝑉 3 2.54 cm 22.59 g 𝑚 𝑜𝑠𝑚𝑖𝑢𝑚 = 1.25 in 𝑥 𝑥 1 in 1 cm3 = 𝟕𝟐𝟑 𝐠 𝐨𝐬𝐦𝐢𝐮𝐦 26 Cont’d ▪ Temperature conversion II: Temperature: ▪ Normal body temperature is 98.6 oF. ▪ Temperature scales: Celsius scale, Convert this temperature to the Kelvin scale, Fahrenheit scale. Celsius & Kelvin scale. ▪ Physical sciences use Celsius scale, ▪ Celsius from Fahrenheit: Kelvin scale. 5 ▪ Engineering sciences use Fahrenheit 𝑡(oC) = [𝑡(oF) – 32] 9 scale. 5 𝑡(oC) = [98.6– 32]= 37.0 oC 9 ▪ Temperature conversion I: ▪ Kelvin from Celsius: Convert 300.00 K to the Celsius scale. 𝑇(K) = 𝑡(oC) + 273.15 ▪ Celsius from Kelvin: 𝑇(K) = 37.0 + 273.15 = 310.2 K 𝑡(oC) = 𝑇(K) -273.15 𝑡(oC)= 300.00 − 273.15 = 𝟐𝟔. 𝟖𝟓 oC 27 The EnD Of Slides 28