Chapter 1 General Chemistry
Chapter 1 General Chemistry CHEM 1411
Lone Star College-CyFair
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This 9 page Class Notes was uploaded by Cassandra Danhof on Wednesday September 21, 2016. The Class Notes belongs to CHEM 1411 at Lone Star College-CyFair taught by Prof. Chakranarayan in Fall 2016. Since its upload, it has received 10 views. For similar materials see General Chemistry I in Chemistry at Lone Star College-CyFair.
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Date Created: 09/21/16
General Chemistry Chap. 1: Chemistry and Measurement Learning Objectives Important Terms 1.1 Modern Chemistry: A Brief Glimpse Provide examples of the contributions of chemistry to humanity ● Discovery of fire began the notice of the way certain rocks/minerals react ○ Ex. Phoenicians extracting purple dye from snails ● Later in the 18th c. chemists began to discover atoms and molecules, so they could synthesize larger molecules into smaller ones ● LCD's in things like computers and phones ● Developing new materials to replace old, inefficient ones ● Genes/DNA in biochemistry/microbiology 1.2 Experiment and Explanation Describe how chemistry is an experimental science Experiment: An ● Chemistry is an experimental science observation of natural because it is based on the scientific method to phenomena carried out help make discoveries and improve one's already in a controlled manner found. so that the results can Understand how the scientific method is an approach to be duplicated and performing science rational conclusions ● The scientific method is an approach to made performing science that is based off of hypothesis, Law: A concise experiments, and conclusions. You first have to statement or make a reasonable, testable hypothesis. Then, you mathematical equation develop and perform your experiment. You do this about a fundamental several times, so that your results are similar. After relationship or this, you reach a conclusion, and other scientists regularity of nature try to test your hypothesis as well. If it is tested(Generalization that enough and proven correct, it is called a theory. Ifsummarizes the data) it is concise, and has been proven correct over a Hypothesis: A tentative long period of time about nature, it is a law. explanation of some regularity of nature Theory: A tested explanation of basic natural phenomena 1.3 Law of Conservation of Mass Explain the law of conservation of mass Mass: The quantity of ● The law of conservation of mass is a law matter in a material that ensures that mass can never either be gained Matter: Whatever or loss. In a chemical reaction, the mass of the occupies space and can before products and the after products will be the be perceived by our same. The only reason it might be different is senses (also has a because of lab/experimental error. mass) Apply the law of conservation of mass (example 1.1) Law of Conservation of ● You heat 2.53 g of metallic mercury in air, Mass: States that the which produces 2.73 g of a red-orange residue. total mass remains Assume the chemical change is the reaction of constant during a metal in the air. chemical change (mass ● Mercury + Oxygen --> red-orange residue of reactants = mass of ● What is the mass of oxygen that reacts? products) Mass of substances before reaction = mass of substances after reaction. 2.53 g mercury + _____ g oxygen --> 2.73 g red-orange residue 2.73 - 2.53 = 0.2 grams oxygen 1.4 Matter: Physical State and Chemical Constitution Compare and contrast the three common states of Solid: The form of matter: solid, liquid, and gas matter characterized ● A solid has a fixed shape and volume; it is by rigidity characterized by molecules that are packed tightly Liquid: The form of together. matter that is a ● A liquid has a fluid shape but a fixed relatively compressible volume; it is characterized by molecules that can fluid move freely Gas: The form of ● A gas has a fluid shape and volume; it is matter that is an easily characterized by rapidly-moving molecules compressible fluid Describe the classifications of matter: elements, States of Matter: three compounds, and mixtures (heterogeneous and forms of matter - solid, homogenous) liquid, and gas ● Elements are substances in their pure state. Physical Change: A These cannot be broken down into a simpler change in the form of substance matter bvc23ut not in ○ Ex. Oxygen its chemical identity ● Compounds are the combination of 2 or more molecules which have been fused Chemical Change (Chemical Reaction): A chemically. In order to break down a compound, it change in which one or needs to be chemically separated. more kinds of matter ○ Ex. Water are transformed into a ● Mixtures are materials that have been new kinds of matter or brought together by physical means. This means several new kinds of you can also separate them by physical means. matter ○ Ex. Water and Sugar Understand the difference between chemical changes Physical Property: A characteristic that can (chemical reactions) and physical changes be observed for a ● The difference between chemical and physical changes is whether or not they can material without changing its chemical change the property(ies) of that substance. A identity chemical change can, but a physical change can't. ○ Ex. Physical Change: Cutting Chemical Property: A characteristic of a an onion material involving its ○ Ex. Chemical Change: Burning an onion chemical change Distinguish between chemical properties and physical Substance: A kind of matter that cannot be properties ● A physical properties is something that can separated into other be observed without changing the property of the kinds of matter by any physical process substance. A chemical properties does take a change of the substance in order to observe. Element: A substance ○ Ex. Physical Properties: Color, that cannot be decomposed by any size, shape ○ Ex. Chemical Properties: The chemical reactions into ability to rust. simpler substances Compound: A substance composed of two or more elements chemically combined Law of Definite Proportions (Law of Constant Composition): A pure compound, whatever its source, always contains definite or constant portions of the elements by mass Mixture: A material that can be separated by physical means into two or more substances Heterogeneous Mixture: A mixture that consists of physically distinct parts, each with different properties Homogenous Mixture (Solution): A mixture that is uniform in its properties throughout given samples Phase: One of several different homogeneous materials present in the portion of matter under study 1.5 Measurement and Significant Figures Define and use the terms precision and accuracy when Unit: Fixed standard of describing measured quantities measurement ● Precision of the closeness of a measured Precision: The quantity and accuracy is the closeness of a single closeness of the set of measurement to the true data. For example. If the values obtained from true data was supposed to show 7.02: similar measurements ○ Precise, Accurate Data: 7.02, of an identical quantity 7.02, 7.03, 7.01 Accuracy: The ○ Precise, Non-Accurate Data: closeness of a single 7.10, 7.11, 7.10, 7.09 measurements to its Learn the rules for determining significant figures in true value reported measurements Significant Figures 1. All digits are significant except zeros at the (SigFigs): Digits in a beginning (leading zeros) of the number and measured number (or a possible terminal zeros (one or more zeroes at the result of a calculation end of a number). Thus, 9.12 cm. 0.912 cm, and with measured 0.00912 cm all contain three significant figures numbers) that include 2. Zeros between nonzero digits are all certain digits plus a significant. (or confined zeros) Ex. 205, 2.05, 2.005final digit having some 3. Terminal (Trailing) zeros ending at the right uncertainty of the decimal point are significant. Each of the Number of Significant following has three significant figures: 9.00 cm, Figures: Number of 9.10 cm, 90.0 cm digits reported for the 4. Terminal zeros in a number without an explicit decimal point may or may one be value of a measured or calculated quantity, significant. If someone gives a measurement as indicating the precision 900 cm, you don’t know whether one, two, or of the value (Number of three significant figures are intended. If the person sig figs = all certain writes 900. cm (note the decimal point), the zeros digits + one uncertain are significant. More generally, you can remove digit) any uncertainty in such cases by expressing the Scientific Notation: The measurement in scientific notation. representation of a Know how to represent numbers using scientific notation number in the form ● A x 10^n where A is a number ● A = coefficient with a nonzero digit to ● 10n = Exponential term the left of the decimal ● n = Exponent point and n is an ○ Ex. 3.00 x 10^8 m/s (rather integer or whole than 300,000,00) m/s number Apply the rules of significant figures to reporting Exact Number: A calculated values number that arises 1. Multiplication and Division: When when you count items multiplying or dividing measured quantities, give or something when you as many Significant figures in the answer as there defines a unit. are in the measurement with the least number of Rounding: The significant figures procedure of dropping 2. Addition and Subtraction: When adding or nonsignificant digits in subtracting measured quantities, give the same a calculation result and number or decimal places in the answer as there adjusting the last digit are in the3 measurement with the least number or reported decimal places Be able to recognize exact numbers and uncertainties ● Inexact - From measurement; has a lestain degree of uncertainty ● Exact - No uncertainty - conversions of sig figs do not apply ● Exact numbers are something that cannot be rounded. For example, 1 dozen eggs cannot be 11.8 eggs. Same goes for the number 9 as oppose to the number 9.0. Know when and how to apply the rules for rounding ● When the number you are rounding has a number on the right of it that is lower than (or equal to) 4, you keep the number ● When the number you are rounding has a number on the right of it that is higher than 4, you round up the number. 1Use significant figures in calculations (example 1.2) 4.18 - 58.16 X (3.38 - 3.01) = -17.3392 2 SF = -17 1.6 SI Units Become familiar with the SI (metric) system of units, International System of including the SI prefixes Units (SI): A particular ● The SI system goes by a system of 10 for choice of metric units the length, mass, and time. It is different for SI Base Units: The SI temperature units from which all others can be derived Quantity Unit Symbol SI Prefix: A prefix used in the International Length meter m System to indicate a Mass kilogram kg power of 10 Meter (m): the SI base Time second s unit of length Temperature kelvin K Kilogram (kg): The SI base unit of mass Amt of Mole mol substance Second (s):The SI base unit of time Celsius Scale: The Prefix Multiple Symbol temperature scale in general scientific use mega 10^6 M Kelvin (K): Absolute temperature kilo 10^3 k hecto 10^2 h deca 10^1 da 10^0 m,g,l deci 10^-1 d centi 10^-2 c milli 10^-3 m micro 10^-6 μ nano 10^-9 n pico 10^-12 p Convert from one temperature scale to another (example 1.3) The hottest place on record in North America is Death Valley in California. It reached a temperature of 134 degrees Fahrenheit in 1913. What is this temperature reading in degrees Celsius? Kelvins? In Celcius: tc= 5∘C x tf−32∘F¿=¿ 5∘C x (134∘F−32∘F) 9∘F 9∘F = 56.7∘C In Kelvins: Tk = ( tc x 1K ) + 273.15 K = ( 56.7∘C x 1∘C 1K 1∘C ) + 273.15 K = 329.9 K 1.7 Derived Units Define and provide examples of derived units SI Derived Units: A unit ● SI Unit of Area = (SI Unit of Length) x (SIderived by combining Unit of Length) SI base units ● Area = length x width Liter (L): A unit of ● Volume = length x width x height volume equal to a cubic ● 1 L isa volume equal to a cube whose side is 1 dm decimeter 3 Density: Mass per unit ● 1 L = 1 dm x1 dm x1 dm = 1 dm volume ○ 1 dm = 10 cm ○ 1 L = 10 cm x 10 cm x 10 cm ○ 1 L = 1000 3 3 cm ● 1000 ml = 1000 cm ○ 1 mL = 1 cm3 Calculate the density of substance (example 1.4) ● Common units: ○ Solids g/ cm3 ○ Liquids g/mL ○ Gasses g/L ● Each substance has a characteristic density ● To identify the unknown substance, a chemist determined its density. By pouring a sample of the liquid into a graduated cylinder, she found that the volume was 35.1 mL. She also found that the sample weighted 30.5 g. What was the density of the liquid? What was the substance? D = m = 30.5g = 0.8969 g/mL v 35.1mL Use density of relate mass and volume (example 1.5) An experiment requires 43.7 g of isopropyl alcohol. Instead of measuring out the sample on a valence, a chemist dispenses the liquid into a graduated cylinder. The density of isopropyl alcohol is 0.785 g/mL. What volume of isopropyl alcohol should he use? V = m = 43.7g = 55.7 mL d 0.785g/mL 1.8 Units and Dimensional Analysis (Factor-Label Method) Apply dimensional analysis to solving numerical Dimensional Analysis problems (Factor-Label Method): 3 3 1L 10 cm The method of 10 cm3 = 10 cm 3 =1 calculation in which one carries along the Convert from one metric unit to another metric unit (example 1.6) units for quantities (a general problem ● Unit Relationship −2 solving method) ○ 1 cm = 10 m or 1 cm = 0.01 m Conversion Factor: A factor equal to 1 that ● Each conversion factor (cf) has a value of 1 converts a quantity 1cm 10 m ○ −2 = −2 , expressed in one unit 10 m 10 m the quantity expressed 1cm = 1 in another unit. 10 m ● Nitrogen gas is the major component of air. A sample of nitrogen gas in a glass bulb weights 243 mg. What is the mass in SI base units of mass (kg)? Since 1 mg = 10 g , you can write −3 10 g −1 243 mg X 1mg = 2.43 X 10 g Then, because the prefix kilo- means 103 , you write 1 3 kg = 10 g And −1 1kg −4 2.43 X 10 g x 10 g = 2.43 x 10 kg Note, however, that you can combine the two conversion steps as follows: −3 10 G 1kg −4 243 mg X 1mg x 10 g = 2.43 x 10 kg Types of Conversion Factors ● Metric to Metric = exact ● English to English = exact ● Metric to English or English to Metric = may or may not be exact Convert from one metric volume to another metric volume (example 1.7) The world’s oceans contain approximately 1.35 x 9 3 of water. What is this is liters? 10 km 1.35 x 10 kim3 Convert from any unit to another unit (example 1.8) How many centimeters are there in 6.51 miles ¿ 1∈¿ 1= 5280 ft 1 = 12∈ 1 ft 1= 1mi 2.54cm ¿ ¿ 5280 ft 12∈ ¿ 1∈¿ 6.51 mi X x 1 ft x 2.54cm = 1.05 1mi ¿ ¿ 6 x 10 cm
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