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CHEM 1315 Exam 1 Study Guide

by: Christa Boettiger

CHEM 1315 Exam 1 Study Guide CHEM 1315 - 002

Marketplace > University of Oklahoma > Chemistry > CHEM 1315 - 002 > CHEM 1315 Exam 1 Study Guide
Christa Boettiger
GPA 4.0

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A comprehensive study guide of listing and defining the major the topics covered for Exam 1. Compiled from lectures and the textbook!
General Chemistry
Fares Z Najar
Study Guide
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This 6 page Study Guide was uploaded by Christa Boettiger on Sunday February 7, 2016. The Study Guide belongs to CHEM 1315 - 002 at University of Oklahoma taught by Fares Z Najar in Spring 2016. Since its upload, it has received 50 views. For similar materials see General Chemistry in Chemistry at University of Oklahoma.


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Date Created: 02/07/16
CHEM 1315 Exam 1 Study Guide Unit 0 (Appendix & Section 1.3, 2.3, 2.5, 2.6, 2.7)  Know how to perform unit conversions  Understand scientific notation  SI Units o Length = meter (m) o Mass = Kilogram (kg) o Time = second (s) o Temperature = Kelvin (K) o Amount = mole (mol)  Conversion Factors to Know: o 1kg = 1000g o 1g = 1000mg o 1L = 1000mL o 0 C = 273.15 K o 1 mol = 6.022 x 10 23atoms o 1mL = 1cm 3 o Converting F to C 5  ( 9 ) x (__ F – 32) = C o Converting C to F  ( 9 x __ C ) + 32 = F 5 o Converting Celsius to Kelvin  ___ C + 273.15 = K Unit 1 (Chapters 1 & 2)  Element: Most simple form, cannot be chemically broken down into a simpler substance ex: Hydrogen (H) o Solid: Definite shape and volume, least kinetic energy, cannot be compressed  Solids and liquids of the same substance have close to the same density o Liquid: Definite volume but NOT definite shape, cannot be compressed o Gas: No definite volume or shape, least dense, most kinetic energy, can be compressed  Compound: a molecule composed of two or more elements in a fixed definite proportion ex: Water (H2O) o Can only be separated by a chemical reaction  Mixture: Substance composed of two or more parts of varying proportions o Homogeneous: Same throughout ex: salt water, air o Heterogeneous: Can differentiate the parts ex: package of skittles  Physical Change: Only alters the state or appearance (atoms/molecules do not change)  Chemical Change: Alters the composition (atoms rearrange)  Bonds: o Ionic: One atom must gain or lose an electron  Metals and non metals often form ionic bonds  Metals gain an electron  Non metals lose an election o Covalent: Atoms share an electron  Scientists: o Leucippus  o Newton  Believed atoms were invisible particles in air o Antoine Lavoisier  Law of Conservation of Mass (1789)  In a chemical reaction, matter is not created nor destroyed o Joseph Proust  Law of Definite Proportions: All samples of a given compound, will have the same proportion of their elements (ex: a water molecule will always have 2 H and O) o John Dalton  Law of Multiple Proportions: When two elements form two different compounds, the masses of element A that combine with 1g of element B can be expressed as a ratio  Atomic Theory  1. Each element is made of atoms  2. All atoms of a given element have the same mass and other properties that distinguish them from other atoms  3. Atoms combine in whole number ratios and form compounds  4. Atoms of one element cannot change into atoms of another element (only can be bound together) o Robert Millikan  Oil drop experiment  Determined the charge of an electron (1.602x10 -19 Coulombs)  Found mass of an electron = 9.10x10 g -28 o JJ Thomson  Cathode Ray experiment  Found atoms are composed of sub atomic particles  “Plum Pudding Model” o Ernest Rutherford  Gold foil experiment  Discovered atoms have a nucleus o James Chadwick  Discovered neutrons  Atoms: o Protons: Positively charged particles in an atom, inside nucleus o Electrons: Negatively charged particles in an atom, outside nucleus o Neutrons: Neutrally charged particles in an atom, inside nucleus o Atomic Mass Unit (amu): 1/12 the mass of a carbon atom. The mass of a proton and neutron are 1 amu o Atomic Number: The number of protons and neutrons in an atom o Mass Number: Sum of an atoms protons and neutrons o Ions: Atoms that are positively or negatively charged due the gain or loss of an electron  Cations: Positively charged ions (lose an electron) Ex: Li + -  Anion: Negatively charged ions (gain an electron) Ex: Fl  Isotopes: o A variation of an atom in which it still has the same number of protons, but the number of neutrons varies o Atoms will have a different number of naturally occurring isotopes o Average Atomic Mass: the number of isotopes that exist for the element multiplied by the % abundance of each isotope o Natural Abundance: the percentages of each variation of isotope for a particular atom  Avogadros Number o One mole of anything is 6.022 x 10 23 23 o A mole of an element contains 6.022 x 10 atoms of that element o The lighter the atom the less mass in one mole of atoms o Atomic Weight  1) Multiply decimal fraction of each isotope by the mass number  2) Add together  Ex: 23.99 amu with 78.99% abundance = (23.99 x 0.7899) Unit 2 (Chapters 3 & 4)  Schrodinger’s cat: An experiment in which scientist Erwin Schrodinger claimed that when covered up, his cat in a gaseous chamber could be both alive AND dead since until its true state was discovered. Relates to how light is both particle and wave  Wave: o Amplitude: vertical height of a crest  Determines the brightness of a light o Wavelength (): The distance between two crests (units = meters)  Determines the color of light o Frequency (): The number of cycles a wavelength passes over a given time (units = s or Hertz (Hz) c o Equation: v= (frequency = speed of light/ wavelength) λ  Speed of light is 3.00 x 10 m/s  1nm = 10 m (convert wavelength from nanometers to meters)  Particle: o Photoelectric Effect: Observation that metals will emit electrons when light of a certain frequency shines on it o The higher the frequency the greater the energy (more damaging) of each photon o Photon: a stream of light particles o Equations:  Planck’s constant is h = 6.626 x 10 -34J s hc  E= (energy of one photon) or E=hv (if given λ frequency)  Atomic spectroscopy: The study of electromagnetic radiation absorbed and emitted by atoms o Atoms absorb energy and emit is as light o Each element emits light of a different color combination  Bohr Model o Electrons move in a circular orbit around the nucleus of an atom o Electrons can move between energy levels  When it drops energy levels (ex: from n = 2 to n = 1) it emits light  The distance dropped determines the color (n = 5 vs n = 2 to n = 1) o Lowest energy state = ground state (ex: n = 1) o Higher energy state = excited state (ex: n = 2 or more)  Brogdile Relation: h -34 o λ= mv (h = 6.626 x 10 ) o Uncertainty Principle: We cannot know an electrons position and velocity at the same time  The more accurately we know the position the less we know of the velocity and momentum (and vice versa) o Exclusion Principle: Two neutrons cannot occupy the same quantum state at the same time  Quantum Mechanics: Mathematical model in which the wave and particle nature of matter are used together o When electrons release energy the change in energy is NEGATIVE o When electrons absorb energy the change in energy is POSITIVE Orbitals o Describe the distribution of electron density in space (the probability of where an electron is) o Orbitals with the same value of n form a shell o Different orbital types within a shell are subshells 2 o The total number of orbitals in a shell is n o Quantum Numbers:  n = principle quantum number  The energy level of the orbital  Size of the orbital (a larger n = farther from the nucleus)  l = angular momentum  defines the shape of the orbital Value 0 1 2 3 of l Type s p d f  m l magnetic quantum number  describes the placement of the orbital  on any energy level there can be 1 s orbital, 3 p orbitals, 5 d orbitals, and 7 f orbitals  m s spin quantum number  describes the electrons magnetic field  always +/- ½ o Pauli’s Exclusion Principle  No two electrons in the same atom can have the same four quantum numbers  Periodic Table of Elements o Organization  Groups = vertical columns  Periods = horizontal rows  Diatomic Elements (exist as 2 atoms)  Br 2 I2, N2, C2 , 2 , 2 , 2  When atoms lose an electron their charge becomes more positive  Noble Gases: elements on the far right side that fill an entire energy level o Metals (on left of periodic table)  Tend to lose electrons during a chemical change (form + ion)  Most elements are metals  High thermal conductivity  High electrical conductivity  Malleability  Ductility o Non Metals (on the right of the periodic table)  Atoms tend to gain electrons during a chemical change (form – ion)  Brittle  Powdery  Solids or gases o Metalloids  Have properties of both metals and non metals  Electron Configuration 5 o Ex: 4p (4 is the energy level, p is the orbital letter, 5 is the number of electrons on that energy level) o The total number of electrons is equal to the sum of the superscripts o S orbitals can hold 2 electrons, p orbitals can hold 6, d orbitals can hold 10, and f orbitals can hold 14 o A half full level is better than partially full o Effective nuclear charge (Z ) eff


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