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General Chemistry 1, ch. 1-4 Study Guide

by: Kristen Manda

General Chemistry 1, ch. 1-4 Study Guide CHEM 1110 - 09

Marketplace > University of Tennessee - Chattanooga > Science > CHEM 1110 - 09 > General Chemistry 1 ch 1 4 Study Guide
Kristen Manda
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About this Document

This study guide covers everything we need to know for the upcoming test on chapters 1-4.
General Chemistry 1
Rebecca E. Stimson
Study Guide
matter, chemical property, Energy, wavelength, frequency, electron, periodic table
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This 5 page Study Guide was uploaded by Kristen Manda on Friday September 16, 2016. The Study Guide belongs to CHEM 1110 - 09 at University of Tennessee - Chattanooga taught by Rebecca E. Stimson in Fall 2016. Since its upload, it has received 61 views. For similar materials see General Chemistry 1 in Science at University of Tennessee - Chattanooga.


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Date Created: 09/16/16
Chemistry Chapters 1­4 Study Guide CHAPTER 1 Chemistry: the study of matter and changes that matter undergoes.  Scientific Method: observations, hypothesis, experiment, theory Hypothesis: a testable explanation which explains the observations Theory: explains data from accumulated experiments, used to predict other results Matter: anything that occupies space and has mass Substance: form of matter that has a constant composition throughout and has distinct properties.  Mixture: combination of 2 or more substances in which their identities remain distinct.  Homogenous mixtures: composition is the same throughout, sugar, water, iced tea Heterogeneous mixtures: composition is different throughout. Sand and iron fill, trail mix Physical Properties: can be observed and measured without changing its identity Chemical Properties: exhibited when interacting with other substances. EX) corrosiveness, flammability Physical Changes: state changes but identity doesn’t. EX) melting, boiling, freezing Chemical Change: change in composition; original no longer there. EX) combustion, digestion, oxidation Significant Figures: 1) Any non­zero digit is significant 2) 0s between non zero digits is significant 3) 0s to the left of a non­zero digit aren’t 4) zeros to the right of the last nonzero digit are significant if decimal is present **When multiplying and dividing, the number of significant figures in the final product or quotient is  determined by the original number that has the smallest number of significant figures. **When adding and subtracting, the answer cannot have more digits to the right of the decimal point than any  of the original numbers.  SI Units/Prefixes Length: meter (m)  Amount of Substance: mole (mol) Mass: kilogram (kg) luminous intensity: candela (cd) Temperature: kelvin (K) electric current: ampere (A) CHAPTER 2 Atom: smallest quantity of matter that still retains the properties of matter Element: substance that cannot be broken down into two or more simpler substances. Alpha Rays: positively charged particles Beta Rays: skews away from negative field Gamma Rays: highest intensity Atomic Number: (z) the number or protons and also the # of  electrons. Mass Number: (A) protons + neutrons  Isotopes: atoms that have the same atomic number but different mass number. A different number of neutrons  and isotopes have similar chemical properties.  Atomic Mass: mass of an atom (amu). 1 amu= ½ the mass of a carbon 12 atom Average Atomic Mass: average mass of the naturally occurring mixture of isotopes.  (isotopic mass)(natural abundance) + (isotopic mass)(natural abundance)  **change all percent’s into decimals Mole: the amount of a substance that contains as many elementary entities as there are atoms in exactly 12g of  carbon 12.  Avogadro’s Number: 6.022e23. There are 6.022e23 atoms in 1 mol.  Molar Mass: the mass in grams of one mole of the substance (g/mol).  People:  John Dalton: said that atoms are tiny indivisible particles JJ Thompson: discovery of the electron; used the cathode ray tube RA Millikan: determined the charge of an electron; examined the motion of tiny oil drops suspended in an  electric field Wilhelm Rontgen: discovered x­rays Antoine Becquerel: discovered radioactivity Ernest Rutherford: gold foil experiment: the nucleus is positively concentrated James Chadwick: discovered the neutron Einstein: a beam of light is a stream of particles Schrodinger: equation for the probability of finding an electron in a certain area of space Louis de Broglie: electrons can act like waves Max Plank: radiant energy is only emitted or absorbed in discrete quantities Chapter 3 Energy: the capacity to do work or transfer heat Kinetic Energy: the energy that results from motion Potential Energy: energy possessed by an object by virtue of its position. Thermal Energy: form of kinetic energy, energy associated with the random motion of atoms and molecules Chemical Energy: the energy stored within the structural units of chemical substances Electrostatic Energy: is potential energy that results from the interaction of charged particles Law of Conservation of Energy: energy cannot be created or destroyed Joule: unit of energy, amount of energy possessed by a 2kg mass moving at a speed of 1m/s.  Wavelength: peak to peak of a wave Frequency: cycles per second Amplitude: height of wave, center to peak or troph Node: amplitude is zero Constructive interference: adding waves in a phase Destructive interference: adding waves out of phase Photoelectric Effect: a phenomenon in which electrons are ejected from the surface of a metal exposed to light  at minimum frequency Threshold Frequency: the minimum frequency of light required to eject an electron from the surface of a metal Bohr’s Theory of the Hydrogen Atom: Bohr made the atomic orbitals.  Heisenberg’s uncertainty principle: it is impossible to know simultaneously both the momentum and the  position of a particle with certainty.  Schrodinger: described electron density (        ) and atomic orbitals. Electron density is the probability where  the electron is most likely to be Principle Quantum Number (n): designates size, collection of orbitals, only in integral numbers Angular Momentum Number (l): designates shape, depends on n, can be anything from 0 to (n­1) l 0 1 2 3 Orbita s­1 p­3 d­5 f­7 l  desig. Magnetic Quantum Number (Ml): specifies orientation Ms: is used to specify an electrons spin. Either positive or negative ½  **(n), (l), and (ml) all required to describe an atomic orbital Pauli Exclusion Principle: no two electrons in an atom can have the same four quantum numbers.  Aufbau Principole: states that electrons are added to the lowest energy orbitals first.  Hund’s Rule: the most stable arrangement of electrons is the one in which the number of electrons with the  same spin is maximized.  Noble Gas Core: shortcut for electron configurations. Only use if asked.  **Exceptions on periodic table for electron configurations are ones ending in the d subshell like chromium and  copper. ** Equationsyoushouldknow: ConstantsIwillprovide: K=C+273.15 SpeedoflightC=3.00x10 m/s D=M/V Planck’sconstanth=6.63x10 J  s orKgm/s 2 23 c=λ NA=6.022x10 E=h 1”=2.54cm;1Km=0.6215mi;1lb=453.6g;1gal= 3.785L Chapter 4 John Newland: law of octaves, when elements arranged in order, every 8  element had similar properties Mendeleev: first periodic table Main Group Elements: s block and p block, 1a through 7a Noble Gases: group 8a and have a full p subshell Transition Metals: group 1B and 3B through 8B Lanthanides and Actinides: Make up the f block elements Valance Electrons: the outermost electrons of an atom that are involved in bonds and predict chemical  properties.  Metals: low ionization energies, cations, good conductors of heat and electricity, shiny, lustrous, malleable, and ductile Nonmetals: vary in color, not shiny, brittle, poor conductors of electricity and heat, have high electron  affinities, anions Metalloids: elements with properties intermediate between those of metals and nonmetals Ionic Radius: the radius of a cation or anion, when an atom loses an electron, radius decreases.  Isoelectronic Series: a series of two or more species that have identical electron configurations but different  nuclear charges.  Periodic Table Trends: **Study Periodic Table, Know first 18 Elements** Effective nuclear Charge: is the actual magnitude of positive charge that is “experience” by an electron in the  atom  Shielding: where an electron partially shielded from the positive charge of the nucleus by the other  electrons  Zeff increases from left to right Atomic Radius: the distance between the nucleus of an atom and its valence shell, or half the distance between  adjacent, identical nuclei  Increases from top to bottom and decreases from left to right Ionization Energy: is the minimum energy required to remove an electron from an atom in the gas phase  Increases from left to right  Takes more energy to remove core electrons Electron Affinity: is the energy released when an atom in the gas phase accepts an electron  Increases from left to right meaning that it becomes more favorable to accept an electron


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