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Review/Ch1 & Lecture of CHEM 101 : Week 1

by: AnnaBanana

Review/Ch1 & Lecture of CHEM 101 : Week 1 CHEM 101

Marketplace > Texas A&M University > Chemistry > CHEM 101 > Review Ch1 Lecture of CHEM 101 Week 1
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Howdy Y'all! This is Week 1 of our Chem 101 Course. This is free.
Chemistry 101
Dr. Ryan Bethel
Class Notes




Popular in Chemistry 101

Popular in Chemistry

This 7 page Class Notes was uploaded by AnnaBanana on Monday September 5, 2016. The Class Notes belongs to CHEM 101 at Texas A&M University taught by Dr. Ryan Bethel in Fall 2016. Since its upload, it has received 89 views. For similar materials see Chemistry 101 in Chemistry at Texas A&M University.

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Date Created: 09/05/16
CHEM 101: BOOK&LECTURE (REV/CH1) Highlight – Important Concepts Highlight – Important Principles Highlight – Key Terms ReviewChapter Intro: Observations are crucial to scientists. Quantitative – measurement (number and scale) Qualitative – immeasurable qualities “very blue” or “extremely cold” R-1 1. Standard Systems of Measurement ensure seamless world travel, even in areas with different languages. a. SI-System: metric-system based measurement method. Prefixes used to change the unit of measurement. b. Volume: Derived from length, not its own SI measurement. Needed in most science experiments. 2. Mass vs. Weight a. Mass: measure of the resistance of an object to a change in its state of motion. In English – force needed to make an object move. b. Weight: force that gravity exerts upon object. R-2(UncertaintyinMeasurement) 1. When measurements are taken, they often vary. Ex. 21.16 21.1  These #s don’t change. They’re certain. 21.15 21.17 .6, .7, .5  These #s do change. They’re uncertain. a. Measurements’ certainty relies on the precision of the measuring device. Ruler vs. yardstick. b. These differences between measurements’ certainty are called significant figures. 2. Significant Figures a. Always record data to correct significant figure decimal places. b. When a pipette/beaker/etc. is measured to +/- 0.01 mL, always mark it with the most detailed unit measured. In this case, it is two decimals, so twenty-five mL would be 25.00 mL, not 25 mL. 3. Precision and Accuracy a. Accuracy – the agreement of a particular value with the true value. b. Precision – the degree of agreement of several measurements of the same quantity. *reflects how easy it is to reproduce the same results in another experiment* c. Random Error – measurement has an equal probability of being high or low. i. Indeterminate - error caused when one (over or under) estimates the value of the last digit. d. Systematic Error – always occurs in the same direction each time. i. Determinate – usually due to a problem with calculations. ** Scientists often assume that the average value of a series of precise measurements is the “best” or “most true” value. However, the average value is only valid if there are no systematic errors in the calculations! R-3(Significant FiguresandCalculations) Rules for Counting Significant Figures 1. Non-Zero Integers Always 42.316 2. Zeroes: (three types) a. Leading Zeroes Never 0.0025 b. Captive Zeroes Always 1.008 c. Trailing Zeroes Decimal 1.00 x 10^6 3. Exact Numbers Infinite bc Exact. 1 in = 2.54 cm MathematicalOperationsforSignificantFigures 1. Multiplication or Division a. You round your answer to the # of SF in the least precise measurement. i. 4.56 x 1.4 = 6.38  6.4 because of SF (only 2 in least precise, so that’s what we keep) 2. Addition or Subtraction a. You round your answer to have the same number of decimal places in the least precise measurement. i. Obviously round by 5 (up) and less than 5 (down). ** Rounding: Dr. Bethel says to: Make sure to round ONLY at the end of a problem, and not during any other time. By keeping your measurements unrounded and un SF’ed, you will gain the most accurate measurement at the end of the problem. R-4(Learning toSolveProblemsSystematically) 1. Approaching a Project: a. What is the goal of this project? b. What do I know about what I’m studying? c. Where do I begin and how do I do it? R-5(DimensionalAnalysis) 1. Converting from one-unit system to the other is best by using the “unit factor method”. a. Dimensional Analysis – AKA “unit factor method” because you’re reducing/increasing units by factors of each other in order to change units. Ex. Inches to lbs. or mL to Kg. Example : I have a pin that is 2.85 cm. How long is it in inches? 1 in = 2.54 cm  Unit Factor 1 ???????? 2.85 2.85 cm x 2.54 ????????x 2.54???????? = 1.12 in R-6(Temperature) 1. The three methods of measuring temperature (Kelvin, Celsius, and Fahrenheit) are used in different areas of science. a. Kelvin and Celsius are used in Chemistry. (But we still need to know Fahrenheit) 2. We must convert from one temperature scale to another with ease. a. Kelvin and Celsius are easy! (just add or subtract 273.15!) i. K to C  Temp(Kelvin) = Temp(Celsius) + 273.15 ii. C to K  Temp(Celsius) = Temp(Kelvin) – 273.15 b. Fahrenheit is harder. We use a ratio of 9/5. i. F to C  (Temp(Fahrenheit) – 32F) x (5C/9F) = Temp(Celsius) ii. C to F  (Temp(Celsius) + 32F) x (9F/5C) = Temp(Fahrenheit) ** Dr. Bethel says make sure to know how to do it and why, and not to just memorize the formula. R-7(Density) 1. Density – the mass of a substance per unit volume of the substance. In English: How thick/thin a substance is, measured by whatever measuring technique the substance uses. ???????????????????????????? = ???????????????? ???????????????????????? R-8(ClassificationofMatter) 1. Matter – anything occupying space and having mass a. Three States: i. Solid – rigid, fixed volume shape ii. Liquid – definite volume, no specific shape, assumes shape of container. iii. Gas – no fixed volume, takes on shape and volume of container. 2. Properties of different substances change with molecule arrangement. a. Mixtures – a substance’s combination that has a “variable composition” i. Homogenous mixture – having parts you can’t tell apart ex. Wine, air, cake batter ii. Heterogeneous mixture – having parts you can tell apart ex. Cement (clicker Q), Fruit Salad 3. Pure Substances – one or more “types” of atoms that combine to make a substance with its own well-defined physical and chemical properties. a. Elements: Cannot be further broken down into simpler substances by chemical or physical means. Gold, Silver, Carbon b. Compounds: substance with constant composition that can be broken into elements by chemical processes. Water, Carbon Dioxide 4. Physical Properties a. Measurable qualities of a substance, observable without changing the substance itself. i. Extensive properties – depend on amount of substance ii. Intensive properties – same, regardless of amount of substance. b. Physical forces cause substances to have qualities. How to separate certain things: Physical Property Method Density Centrifugation Boiling Point Distillation State of Matter Filtration Intramolecular Forces Chromatography Vapor Pressure Evaporation Magnetism Magnets Solubility Filtration R-9(Energy) Vocabulary Energy – the ability to do work or produce heat Kinetic Energy – energy of motion. This happens when something is moving. Potential Energy – stored energy due to position. When something is about to move or go, it has PE. Heat – energy flowing between two objects due to a temperature differential. Work – force acting over a distance. ** Dr. Bethel wants you to know that: the total energy content of the universe is constant. Joules – kilogram meter squared per second squared. AKA: (???????? ∗ ???? 2 ) 1 ???????????????????? = 2 ???? R-10(Moles) 1. What are moles? a. Moles are a placeholder for a large number of atoms. Atoms cannot be weighed individually and are small. So, we use a placeholder that lets us count with ex. 1 mole of CO etc. instead of 6.022 x 10^23 atoms CO. Much simpler and easier math that way! Chapter1 1-1(Chemistry:TheAtomsFirstApproach) Important Vocab - Atom – smallest unit of an element with all its properties - Molecule – smallest unit of a pure substance with all its properties - Ion – atom or a group of atoms that has an electric charge (lost or gained electron) Molecular Formula – Shows molecule composition 1. More than one molecule has the same formula (ex. Glucose and galactose). It’s just differently placed in 3D, even though it’s written the same way. Ionic Compounds have formula units ONLY. Reaction – one substance changes to another by changing how it is arranged. DIATOMIC MOLECULES!! - When atoms decide to be BFFs with each other. They make a 7 on the periodic table (except Hydrogen): These include: o Hydrogen o Nitrogen o Oxygen o Fluorine o Chlorine o Bromine o Iodine 1-2(Scientific Method)& 1-3 Scientific Method – a way of conducting research that makes predictions & then tests them out. 1. Making observations (and collecting data) 2. Suggesting a possible explanation (formulating a hypothesis) 3. Doing Experiments to test the possible explanation (testing hypothesis) Scientific Models - Theory – aka model. Series of tested hypotheses that gives an overall explanation for some natural phenomena. o INTERPRETATION of why nature is a certain way. o Theories keep being made, even if something already seems to account for a certain phenomenon. - Natural law – observation that applies to many different systems. - Observation – something that can be witnessed and recorded. - Law of Conservation of Mass – total mass of materials isn’t affected by a chemical change in those materials. **Dr. Bethel wants you to know: A law summarizes what happens; a theory (model) attempts to explain why it happens. **Dr. Bethel also wants you to know: in this class, HISTORY DOESN” T MATTER!!! Yay! Less notes!! But know experiments. 1-4 (FundamentalChemicalLaws) Law of Conservation of Mass - total mass of materials isn’t affected by a chemical change in those materials. Law of Definite Proportion – Compounds have constant proportions of elements by mass. Law of Multiple Proportions - when 2 elements form a series of compounds, the ratio of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers. Methane 1g Carbon = 0.336 g H Ethane 1 g Carbon = 0.252g H ** Any set of carbon/hydrogen molecules has a whole # ratio 1-5 (Dalton’sAtomicTheory) 1. Dalton’s Atomic Theory a. How to relate masses to many different atoms? (elements weight different amounts!) b. Make a chart – so Dalton made a chart, but it wasn’t right because he didn’t know how to weigh atoms right and he didn’t know about the diatomic molecules. Dalton’s Atomic Theory 1. Each element consists of tiny particles: Atoms. 2. All atoms of a same element are identical. Atoms of other elements are different. 3. Chemical compounds form when different elements combine. 4. Chemical reactions reorganize atoms, not change atoms themselves. 2. Joseph Gay-Lussac & Avogadro a. Measure/arrange by volume, not mass b. Volumes of gases at = temperature and pressure have the same number of particles. Avogadro came up with the mole, but nobody listened. 1-6 (AtomViews) 1. J.J. Thompson a. Invented cathode ray tube which glowed when high voltage charge passed through the tube. The negative side made a ray that another negative point repelled. The ray was actually a stream of electrons, and it didn’t like the negative charge. b. Same for every atom and electron sample. c. Found charge-to-mass ratio of electrons: i. ???? = −1.76 ???? 10 8???? ???? ???? 2. Millikan a. Charged oil drop experiments i. Electrify oil. ii. Drop between magnets. iii. Find charge of single electron. ???????????????? ???????? ???????????????????????????????? = 9.11 ∗ 10 −31 ???????? ???? ????ℎ???????????????? ???????? ???????????????????????????????? = 1???? = −1.602 ∗ 10 −19 ???? 3. Rutherford in 1911 a. Alpha particles shot through foil to hit detector. b. Some deflected, which meant that there was something positive that they hit which made them bounce off. i. Revolutionary because nobody knew there was a central nuclear blob in an atom. c. Nucleus – positive blob in center of atom. We now know that a nucleus is the center of an atom and consists of protons and neutrons in varying degrees, depending upon the element and the isotope. 1-7 (Atoms) −24 1. Protons – Positive charge, Nucleus, 1.67 ∗ 10 ???? Mass : 1.007u 2. Neutrons – No Charge, Nucleus, 1.009u 3. Electrons – Negative Charge, Outside Nucleus. 0.0005u Atomic Mass - Atomic mass is relative to other atomic mass of other elements. o O atom = H atom x 16 - Today’s standard definition is Carbon-12 o Mass of 6 protons and 6 neutrons is set to exactly 12u. o Mass of one Carbon 12 atom = 1.99265 x 10^-25 g o 1u = (mass of one carbon-12 atom / 12) = 1 Dalton Mass Number - Always whole #! - Nucleons = Protons + Neutrons - Mass # is number of nucleons. AKA protons and neutrons added together. Terms: - Atomic Number o is the number of protons. Always the same for every element. Determines where the element is on the periodic table. Mass #   Atomic Symbol Atomic #  M as # Isotopes - When neutron # changes in the SAME element (and everything else stays the same). Because it’s neutrons, it has the same charge (nothing because protons and electrons cancel out). - Carbon – 13 Can be written as 12 OR 12 and is still the same element. C C 6 - Most atoms don’t change very much after having a neutron changed, but atoms like Hydrogen (which have very little mass at all) change dramatically when you add a neutron. (you’re doubling its size). o Regular H20 becomes “heavy water” which you shouldn’t drink! Ions - Adding a charge to an atom. o By removing an electron, you make an ion. o Remove e (charge up)  Positive ions are CATIONS o Add e (charge down)  Negative ions are ANIONS. Howdy, Y’all! Thanks for using my notes! I hope they help you study for CHEM 101. Please note that these are my notes and you shouldn’t be passing them off as yours. Find me on StudySoup! (There, you can buy all my notes & study guides for the whole semester!) Please let me know if you have any questions! Thanks, Anna


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