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CHEM 1331 First Exam Review

by: gypsgirl

CHEM 1331 First Exam Review 1331

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About this Document

Covers Chapters 1-4 Update: I was just informed that the test will not cover Ch. 4 and certain parts of Ch. 3, however I will leave it on here in case anyone wants to use it later on for studying.
Fundamentals of Chemistry
Mark A Smith
Study Guide
Chemistry, atoms, Molecules, Models
50 ?




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This 5 page Study Guide was uploaded by gypsgirl on Wednesday September 21, 2016. The Study Guide belongs to 1331 at University of Houston taught by Mark A Smith in Spring 2015. Since its upload, it has received 126 views. For similar materials see Fundamentals of Chemistry in Chemistry at University of Houston.


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Date Created: 09/21/16
Chem 1331 First Exam Review Ch. 1 - Scientific Method  Make observations  Formulate hypotheses  Perform experiments - Models/ Theories = explanations of nature’s behavior  Subject to medication over time and aren’t always correct - Quantitave Observations= Measurements  Consist of a number and a unit  Have some uncertainty (ex. sig figs)  SI is preferred - Fundamental Laws  Conservation of Mass- mass is neither created nor destroyed  Definite Proportion- a compound always has exactly the same proportion of elements by mass  Multiple Proportions- when 2 elements bond to create a series of compounds, the ratios of the masses of the second element can always be reduced to small whole numbers when combined with one gram of the first element; basically you can combine two elements and get different molecules ex. SO and SO2 - Dalton’s Atomic Theory  All elements are made of atoms  All atoms of an element are identical  Chem. Compounds formed by atoms combining  Atoms do not change in chem. reactions, but the way they bond does - Early Atomic Experiments and Models  Thomson Model- used cathode-ray tubes to discover the charge-to- mass ratio of electrons; plum pudding model- electrons are dispersed throughout a positive charge cloud  Millikan Experiment- used charged oil drops to determine the mass of an electron- 9.11x 10^-31 kg.  Rutherford Experiment- alpha particles were shot at a thin sheet of gold foil; many particles were deflected proving there was a nucleus in the atoms  Nuclear Model- positively charged nucleus at center of atom, surrounded by negatively charged electron cloud - Atomic Structure  Small, dense nucleus made of: protons (+ charge) and neutron (no charge)  Electrons (- charge, mass= 1/1840 of a proton) hang out around the nucleus in the large amount of remaining atomic volume  Isotopes have same atomic number as element but different mass number (# of protons (atomic number) + # of neutrons) Ch. 2 - Electromagnetic Radiation Defined by wavelength (upside down v with longer left leg) , frequency (v), and speed (c= 2.9979 x 10^8 m/s) (wavelength)v = c Viewed as a stream of “particles” (photons)  Each particle as an energy of hv, h is Planck’s constant (6.626 x 10^-34 J. s) - Photoelectric Effect Electrons are emitted when light strikes a metal surface Einstein suggested electromagnetic radiation can be viewed as a stream of “particles” based on his analysis of emitted electrons’ amounts and kinetic energy - Hydrogen Spectrum Emission spectrum of hydrogen shows discrete wavelengths Says that hydrogen has discrete energy levels - Bohr Model of the Hydrogen Atom  Bohr created a model using data from the hydrogen spectrum and angular momentum assumptions (that is quantized- energy can only occur in discrete units called quanta) in which electrons travel in circular orbits  Completely wrong, however it is still important - Wave (Quantum) Mechanical Model  Electrons are standing waves  Square of the wave function (aka an orbital) shows the probability distribution for an electron’s position  We can never know the EXACT position of an electron (consistent w/ Heisenberg’s uncertainty principle)  Heisenberg’s Uncertainty Principle- you can’t accurately know the position and momentum of a particle at the same time  Define orbital shapes with probability maps  Orbitals characterized by quantum numbers: n, l, ml  n- relates to the size and energy of an orbital, positive integer  l- relates to the shape of the atomic orbital, 0 to n-1  ml- relates to orientation of orbital relative to same l orbitals, integer values of -l to l, ex. l= 2 so ml can be -2, -1, 0, 1, or 2 - Electron Spin  Quantum number, ms (+ or – ½)  Pauli Exclusion principle: no two electrons can have the same values of the quantum numbers n, l, ml, and ms  An orbital holds two electrons that have opposite spins - Periodic Table  The aufbau principle explains the way the periodic table is organized  The Aufbau Principle (wave mechanical model)- protons are added one by one to the nucleus to build elements, electrons are similarly added  Atoms in a given group have the same valence electron configuration  Trends in the periodic table (ex. Ionization energies and atomic radii) are explained by the concepts of nuclear attraction, shielding, penetration, and electron repulsions Ch. 3 - Chemical Bonds  Hold atoms in molecules together  Occur when some atoms can lower their total energy by bonding  Types:  Ionic- electrons are transferred to make ions  (Nonpolar) Covalent- equal sharing of electrons  Polar Covalent- unequal sharing of elections  Percent ionic character of a bone X-Y  ((Measured dipole moment of X-Y)/(Calculated dipole moment for X+ Y-))x 100%  Electronegativity- relative ability of an atom to attract shared electrons  Polarity of a bond depends on electronegativity difference of bonded atoms  Spatial order of polar bonds determines whether a molecule has a dipole moment  Dipole moment- the mathematical product of the separation of the ends of a dipole and the magnitude of the charges - Ionic Bonding  Ion has a different size than the normal atom  Anion- larger, negative  Cation- smaller, positive  Lattice energy: change in energy when ions group together to create an ionic solid (ex. salt) - Bond Energy  Energy needed to break a covalent bond  Increases as # of shared pairs goes up  Used to estimate energy change in chemical reactions - Lewis Structures  Illustrate how valence electron pairs are arranged among atoms in molecules or polyatomic ions  Most stable molecules have atoms with filled valence orbitals  Dual rule for hydrogen  Octet rule for second row elements  Atoms in the third row and beyond can go past the octet rule  Resonance- multiple ways the atoms can bond in the same molecule  Formal charge is used to choose which resonance structure is best - Compounds named depending type and other rules  Binary Compounds  Type 1- has a metal that always makes the same cation  Type 2- has a metal that can make more than one cation  Type 3- has two nonmetals  Compounds containing a polyatomic ion Ch. 4 - VSEPR Model  Based on the fact that electron pairs can be arranged around a central atom so that electron repulsions are minimized  Used to predict geometric structure of most molecules - Dipole Moment  Spatial arrangement of polar bonds determines whether the molecule has a dipole moment - Two Widely Used Bonding Models  Localized electron model  Molecular orbital model - Localized Electron Model  Molecule is pictured as a group of atoms sharing electron pairs between atomic orbitals  Hybrid orbitals (combos of native atomic orbitals) required to explain the molecular structure  d2sp3- octahedral, 6 e- pairs  dsp3- trigonal bipyramidal, 5 e- pairs  sp3- tetrahedral, 4 e- pairs  sp2- trigonal planar, 3 e- pairs  sp- linear, 2 e- pairs - Two Types of Bonds  Sigma- electrons shared via area shared by two electron clouds of two atoms, centered on a line that joins the atoms  Pi- shared electron pair from two parallel electron clouds of two atoms, space above and below a line that joins the atoms - Molecular Orbital Model  Molecule assumed to be a new thing containing positively charged nuclei and electrons  Electrons in molecule contained in molecular orbitals (constructed from atomic orbitals of the atoms in the molecule)  Correctly predicts relative bond strength, bond polarity, and magnetism  Correctly shows electrons as delocalized in polyatomic molecules  Main disadvantage: difficult to apply to polyatomic molecules qualitatively - Molecular Orbitals Classified by Energy and Shape  Energy  Bonding MO- lower in energy than its original atomic orbital, lower in molecule than in separated atoms, favor molecule formation  Antibonding MO- higher in energy than its original atomic orbital, higher in molecule than in separated atoms, do not favor molecule formation  Shape/ Symmetry  Sigma MO- centered electron probability on line passing through nuclei  Pi MO- electron probability concentrated above and below line passing through nuclei - Bond Order is Index of Bond Strength  Bond order= ((# of bonding electrons) – (# of antibonding electrons))/2 - Resonance Molecules Best Described by Combining LEM and MOM  Sigma bonds- localized  Pi bonds- delocalized


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