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General Chemistry I 177

by: Megan Spiegel

General Chemistry I 177 Chem 177

Marketplace > Iowa State University > Chemsitry > Chem 177 > General Chemistry I 177
Megan Spiegel
GPA 3.85

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The notes uploaded are all of the lectures from week 1(Monday, Wednesday, Friday) in Dr. Anderson's Chem 177 Course. These notes have been formulated to be easy on the eyes and concise yet detai...
General Chemistry I
Dr. Anderson
Class Notes
General Chemistry
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This 12 page Class Notes was uploaded by Megan Spiegel on Saturday August 27, 2016. The Class Notes belongs to Chem 177 at Iowa State University taught by Dr. Anderson in Fall 2016. Since its upload, it has received 119 views. For similar materials see General Chemistry I in Chemsitry at Iowa State University.


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Date Created: 08/27/16
Chem 177: Introduction to Matter and Measurement (Chapter One)   ❏ What is Chemistry?­ the study of the properties and behavior of matter.   ❏ What is Matter?­ anything that has mass and takes up space   ­ composed of atoms     ◼Methods of Classification   ❏ The States of Matter   a. Solid   b. Liquid   c. Gas  d. Plasma  ❏ Composition   e. Heterogenous Mixtures­ not uniform throughout    f. Homogenous Mixtures­ uniform throughout   ­ Can be solutions and mixtures   ❏ What is a Substance?­ has distinct properties and composition that does not vary  between samples  ­ 2 substances= elements and compounds     ◼The Law of Constant Composition (or Definite Proportions): which states  that all samples of a given chemical compound have the same elemental  composition by mass.  e.g. oxygen makes up about ​ 8/9​of the mass of any sample of pure water, while  1​ hydrogen makes up the remaining ​ /9​of the mass.    ◼Types of Properties  ❏ Physical: observed without changing the substance   e.g. “aluminum is shiny” ­ being shiny is a physical property   ❏ Chemical: observations made when the substance is changed   e.g. when a metal starts to rust­ rust is a different substance than metal­ when  nickel is transitioning from a metal to rust.   ­ Intensive Properties: independent of the amount of substance there is   e.g. boiling pt. Of an element or substance   ­ Extensive Properties: dependent on the amount of substance there is   e.g. mass     ◼Types of Changes   ❏ Physical: types of changes that do not change the composition of the substance   ❏ Chemical: types of changes that do change the composition of the substance   i.e. usually involves energy/combustion  e.g. turning Hydrogen Peroxide H​ O​ 2​2​o just elements, H and O (involves heat  energy)     ◼Segregating Mixtures  ❏ Filtration: solids separated from liquid solutions   ❏ Distillation: uses differences in boiling points between elements to separate  homogenous mixtures.   ❏ Chromatology: separates substances on the basis of adhering to solid surfaces.   e.g. to separate ink from water you dip in paper because ink adheres to paper     ◼Numbers and Chemistry   ­ Many topics in Chemistry are quantitative   ❏ Concepts of Numbers in Science   ❏ Units of measurement   ❏ Quantities measured and calculated   ❏ Significant Figures   ❏ Uncertainty in measurements   ❏ Dimensional Analysis (Dr. Anderson is very strict on this topic­ be sure to perfect it!)   ­ Useful knowledge for Stoichiometry   ❏ SI Units: the universal code of measurement in Chemistry   ­ A different base unit is used for each quantity   → Mass has a base unit of Kilograms   → Length has a base unit of Meters   → Volume has a 2 base units of Liters (1 decimeter [1 cubic decimeter]) and  Millimeters (1 centimeter [1 cubic centimeter])  ❏ Temperatures   ­ In general heat flows from an object with a higher temperature to an object with a  lower temperature   ­ The celsius and kelvin scale are most often utilized   Kelvin = 273  Instant celsius conversions to know!:   0°C = 32°F (freezing pt. Of water)   100°C = 212°F (boiling point. Of water)     ◼Density: a physical property of a substance   3​ ­ Units are derived from mass and volume (g/mL or g/cm​ )  ­ Equation: D=M/V    ◼Numbers Encountered in Science  ❏ Exact Numbers: counted or always given by definition   e.g. 12 individual eggs is given by the term “one dozen”  ❏ Inexact Numbers: (A measured number) usually depends on how the number was  determined   ­ All measured numbers have some degree of human error or uncertainty         ◼Accuracy v. Precision     ❏ Accuracy: refers to the proximity of a measurement to the true value (how close are you  to the correct value).  ❏ Precision: refers to the proximity of many measurements to each other (how close in  value 3 measurements are to each other → ​ measurements are within 0.3 of each other.)           ◼Significant Figures   ­ Refers to digits that were measured   ­ Help to not overstate the accuracy of our answers  ❏ Checklist for Significant Figures                                               ◼Dimensional Analysis: converting one measurement to another   e.g. 1 inch = 2.54 cm     ❏ Example         Atomic Theory of Matter (Chapter Two)   ◼John Dalton (19th century)   ­ Came up with the idea that atoms are the fundamental building blocks of matter     ◼Dalton’s Postulates   1. Elements are composed of atoms   2. Atoms of one element are the same; atoms of two different elements are different   I.e. the formation of compounds   3. Atoms are not changed in chemical reactions   4. Atoms are neither created nor destroyed; they always were   5. Compounds that are the same have the same ratio of atoms     ◼Law of Conservation of Mass: total mass of substances present at the  end of a chemical process of the same amount of mass that it was at the  beginning (before the chemical process occurred)  e.g. S ​(32g) ​ + Fe ​(56g)​ → SFe ​(88g)   [32+56=88g]    ◼Law of Multiple Proportions: when combining elements weights the ratio will always be that of  a small whole number.   E.g. O​ 2​f NO → O​  o2​NO​  {2​st balance} → 1.143 of 2.286 → 1.143/1.143 of  2.286/1.143 = 1:2 ratio    ◼Discovery of Subatomic Particles   ❏ Dalton believed atoms were composed of even smaller particles   i.e. protons, neutrons and electrons   ­ Electron Cathode Ray Experiment (by J.J. Thompson) assured the  existence of subatomic particles called electrons     ◼Radioactivity: spontaneous emission of high energy radiation by an atom  ­  Henri Becquerel was responsible for this discovery followed by Marie and Pierre  Curie  ­ Their experiments proved that there was more energy and subatomic particles to  an atom that originally thought.   ❏ Half Lives of Elements: Radioactive deterioration of an element (more detail later in  chapters)   ❏ Three Types of Radioactivity   1. Alpha Particles [ ???? ] (positively charged)   2. Beta Particles [ ???? ] (negatively charged)   3. Gamma Rays [ ???? ] (uncharged)   ­ Some Gamma Rays have been known to be deflected by protective glass    ◼The Atom (circa 1900)  ❏ Thompson’s “plum pudding” Description of an Atom: “A positive sphere with electrons  embedding within it.                                   ◼The Nuclear Atom: has a nucleus, the rest is open empty space.   ❏ Rutherford’s Theory: developed the planetary model of the atom which put all the  protons in the nucleus and the electrons orbited around the nucleus like planets around  the sun  ❏ Later, protons and neutrons were discovered   ❏ Protons (+) neutrons (0) and electrons (­)   ­ Protons and Neutrons masses= 1 → found within the nucleus   ­ Electron mass= 0 → found orbiting the nucleus   ❏ Atoms are weighed AMU’s (Atomic Mass Units)   ­ 1 AMU = 1.66054 x 10​ g  ­24​ ­ weight and mass can be measured by a mass spectrometer   ­ On a mass spectrometer the mass of an atom is compared to the mass of a  carbon atom which has 6 neutrons and 6 protons.     ◼Diagram of an Element               ◼Isotopes: atoms of the same element with  different masses   ❏ Carbon is the most widely used isotope   ❏ Isotopes have the same number protons  but different number of neutrons     ◼Periodic Table     ❏ Rows = “periods”   ❏ Columns = “groups”   ❏ Elements that are in the same group have the same chemical properties   HELPFUL HINT FOR EXAMS: know and memorize the 7 elemental groups and be able to  name a few elements from each group!   1. Alkali Metals​ → Lithium, Sodium, Potassium, Rubidium, Cesium, Francium  2. Alkaline Earth Metals​ → Beryllium, Magnesium, Calcium, Strontium, Barium, Radium  3. Transition Metals ​ → Scandium, Titanium, Vanadium, Chromium, Manganese, Iron,  Cobalt, Nickel, Copper, Zinc, Titanium, Zirconium, Niobium, Molybdenum, Technetium,  Ruthenium, Rhodium, Palladium, Silver, Cadmium, Hafnium, Tantalum, Tungsten,  Rhenium, Osmium, Iridium, Platinum, Gold,Mercury, Rutherfordium, Dubnium,  Seaborgium, Bohrium, Hassium, Meitnerium, Ununnilium, Unununium, Ununbium  4. Metalloids or Semimetals​ → Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium,  Polonium  5. Nonmetals​ → Hydrogen, Carbon, Nitrogen, Phosphorous, Oxygen, Sulfur, Selenium,  Fluorine, Chlorine, Bromine, Iodine, Astatine  ** To help you in future classes, please consider printing out the following Periodic Table that I created and bring it to every Chem class and recitation**


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