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Chem 102 Exam 1 Study Guide

by: Kaitlin Notetaker

Chem 102 Exam 1 Study Guide CHEM 102 001

Marketplace > University of South Carolina > Chemistry > CHEM 102 001 > Chem 102 Exam 1 Study Guide
Kaitlin Notetaker
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About this Document

This study guide follows an outline of all the main points of what is going to be on our exam for chapters 1-10.
Fundamental Chemistry II
James Sodetz
Study Guide
Chemistry, chemistry 102, chem 102, Sodetz, James Sodetz, chapter 1-10, University of South Carolina
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This 9 page Study Guide was uploaded by Kaitlin Notetaker on Friday February 5, 2016. The Study Guide belongs to CHEM 102 001 at University of South Carolina taught by James Sodetz in Spring 2016. Since its upload, it has received 120 views. For similar materials see Fundamental Chemistry II in Chemistry at University of South Carolina.


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Date Created: 02/05/16
Friday, May 20, y Chemistry 102 Exam 1 Study Guide Chapters 1-5 ­ Elements, Isotopes and Atomic Weights • particles in the atom: ­     oton   (P) ­ positive (+) charge, 1 amu •  located in the nucleus ­     utron   (n) ­ neutral charge, 1 amu • located in the nucleus  electron ­ ­        (   ) ­ negative (­) charge, 0 amu • located in electron cloud surrounding nucleus • important definitions: atomic ­        number ­ number of protons of the atom (determines identity) mass ­        number ­ sum of the number of protons and neutrons in the atom     isotopes ­ atoms of the same element (same number of protons) with different  number of neutrons  ­     omic   weight ­ average mass of all naturally occurring isotopes of an element  (expressed in amu) ­ Metals, Nonmetals and Noble Gases •    metal ­ has luster (shiny, reflective) and is a good electrical conductor ­ located  on center and left side of periodic table •    nonmetal ­ typically a nonconductor  ­ located  on top right part of periodic table •    metalloid ­ element with properties of metals and nonmetals •    noble gases ­ stable, inert gases in group 8A  1 Friday, May 20, y ­ Valence Electrons, Electron Dot, Octet Rule •   valence electrons ­ electrons found in outermost shell of atoms ­ involved  in all chemical reactions ­ for groups 1A­8A:  group number = number of valence electrons  •   electron­dot symbols ­ an element’s symbol surrounded by a dot for each valence  electron they possess  ­ examples  in table 2.5 on page 65 of the textbook  •   octet rule ­ elements tend to combine by gaining or losing electrons in order to  attain 8 valence electrons (noble gas electron configuration) ­ groups  1A­3A tend to lose electrons and groups 5A­7A tend to gain electrons ­ Ionic Bonds, Compounds and Formulas • Naming Ionic Bonds  ­   cation  ­ a positively charged ion (+) formed when an atom loses one or more  electrons (typically groups 1A­3A) • named by identifying the metal followed by “ion” ­ ex : Mg 2+  magnesium ion ­   anion  ­ a negatively charged ion (­) formed when an atom gains one or more  electrons (typically groups 5A­7A) • named by replacing the end of an element name with “­ide” followed by “ion” ex ­ : F­  fluoride ion ­   Polyatomic Ions  ­ group of atoms with net charge and behave as single particle • table of common polyatomic ions on page 85 • Ionic Compounds      oni  bonding ­a “transfer of electrons” that involves the electrostatic attraction  from oppositely charged ions (cations + anions)      ionic formula ­ shows lowest possible ratio of atoms with zero net charge  • ex: K +F  —> KF 2 Friday, May 20, y + 2­ • ex: K +O  —> K O 2 • more examples 87 and practice on page 97 • Naming Ionic Compounds ­ the cation is named first with its unchanged name followed by the anion,  changing the ending of the name to ­ide • ex:  CaO (calcium + oxygen) = calcium oxide  ­ Covalent Bonds, Sharing of Electrons, and Naming Bonds •   covalent bond ­ bond formed when atoms share electrons to achieve noble gas  configuration  ­ multiple  covalent bonds: • single bond ­ shares one electron pair • double bond ­ shares two electron pairs • triple bond ­ shares three electron pairs      onpolar   covalent bond ­ equal sharing of electrons      olar  covalent bond ­ unequal sharing of electrons  ­ naming  covalent bonds: • rules for the order of element names: elements ­  farthest left in periodic table or closest to bottom in a group are first hydrogen ­  ­ first with 6A & 7A, second with 1A­5A ­ oxygen  ­ always second except with fluorine • their numerical prefixes (based on subscripts):  • one  ­ mono­ •  ­ di­ two • three ­ tri­ • four ­ tetra­  • five ­ penta­ 3 Friday, May 20, y • six ­ hexa­ ­ Lewis Structures • Drawing Lewis Structures ­ general  method : • 1. sum total number of valence electrons in molecule/ion • 2. draw skeleton structure and subtract 2 electrons for each pair  • 3. add lone pairs using remaining electrons so each atom receives an octet • 4. place any remaining electrons as a lone pairs on the central atom  • 5. if central atom still does not have octet, take lone pair from neighboring atom to form double bond  ­ polyatomic  ions are determined the same way, but charge of ion needs to be  considered  • ex: NH  4 N(5 VE x 1 atom ) + H(1 VE  x 4 atoms) ­ 1(charge) = 8  • *know the common bonding patterns on slide 60!! ­ VSEPR Model  • the shape of a molecule depends on the VSEPR ­ Valence­Shell Electron­Pair  Repulsion  ­ predicts  molecular shape based on a molecule’s:  steric ­         number ­ number of lone e  pairs + number of atoms bonded to central  atom       lectron  pair geometry ­ describes arrangement of atoms bonded to central  carbon and lone electron pairs      olecular   geometry ­ describes only arrangement of atoms bonded to central  carbon   ­ table on slide 59 and page 115 Electronegativity, Bond Polarity and Molecular Polarity ­ • Bond Polarity 4 Friday, May 20, y dipole         moment ­ the measure of the unequal sharing of a molecule (net polarity) • *not an actual moment in time, is a measurement.  • molecules with a dipole moment are polar and molecules lacking one are  nonpolar    Electronegativity  ­ stronger  towards the top and right­side of periodic table  • electronegativity > 2 is ionic  • electronegativity < 2 is covalent • Polarity of Molecules ­ depends  on shape, magnitude and direction  • equal magnitude, symmetrical shape and opposite directions tends to equal no dipole moment • different magnitudes, same direction and bent shapes tend to equal a dipole moment 5 Friday, May 20, y Chapters 6-10 ­ Chemical Equations and Balancing  • chemical equation ­ the symbolic notation that represents a chemical reaction  ­ reactant  + reactants —> products • balanced equations have same number of atoms for each element on both  sides of the equation (shows what happens during reaction) ­ can  only use coefficients to balance equations, no subscripts  ­ practice and examples on page 136 ­ Types of Reactions •    Precipitation Reaction ­ soluble reactants react in water to form a solid, insoluble  product      verall  equation shows balanced equation       omplete   ionic equation separates overall equation and shows all electrolytes as  ions net         ionic equation cancels out spectator ions (do not participate in reaction, show up on both sides) and only shows ions that undergo change during reaction     Neutralization Reactions ­ occurs  between an acid and a base  • nonmetal anion from acid + metal cation from base —> water +salt  • know common acids and bases table on slide 102 •    Oxidation­Reduction Reactions  ­ occurs  when electrons are transferred ­ key  definitions:      oxidation ­ loss of electrons, gain of O or loss of H     reduction ­ gain of electrons, loss of O or gain of H ­ assigning  oxidation numbers (chart on slide 104) 6 Friday, May 20, y • elements have charge of 0 • monatomic ion’s number is the charge of ion • when in combination: F is ­1, O is ­2, H is +1 with nonmetals and ­1 with metals and other halogens are ­1 ­ Equilibrium • caused by reversible reactions:  as product accumulates, it may convert back to reactant •    chemical equilibrium ­ when the forward rate and reverse rate are equal ­ does  not mean amounts of reactants and products are equal, just their formation rates •    Equilibrium Constant ­ K eq ­ amounts  of products and reactants at equilibrium ­  K eq , then products > reactants (and vice versa) ­ Moles •    molecular weight ­ sum of all atomic weights in the molecule •    molar mass ­ mass (in grams) of 1 mole of a substance  ­ molecular  weight (amu) = molar mass (g/mol) 1 mole = 6.022 x 10 23  ­ entities  • Converting Moles and Mass ­ moles  —> mass (using molar mass) ­ examples  on page 161­162 ­ Intermolecular Interactions    Dipole­Dipole Attraction ­ results  from electrostatic forces between molecular dipoles (polar molecules)    London Dispersion Forces ­ results  from attractions between instantaneous and induced dipoles •    Hydrogen Bonding ­ occurs between H and N,O or F and a lone pair 7 Friday, May 20, y ­ networking  between lone pairs and bonds makes it hard to break up which is why the boiling point is so high Solutions, Molarity, Dilutions ­   Solutions solutes ­  are dissolved in solvents (it is considered an aqueous solution when  water is the solvent) • strong electrolyte is a compound that completely dissociates into ions  • weak electrolyte is a compound that only partially dissociates into ions    Molarity  ­ number  of moles of solute in one Liter of solution    Dilutions  ­ adding  solvent to lower concentration of a solution  • molarity(conc) x volume (conc) = molarity (dil) x volume (dil) • examples on page 272 ­ Acids and Bases • Arrhenius vs. Bronsted­Lowry  ­   Arrhenius  ­ acids and bases must be dissolved in water  • acids produce H+ ions and bases produce OH­ ions when dissolved in water      ronsted  ­Lowry ­ not limited to aqueous solutions  • acid is a proton (H+) donor and bases are proton (H+) acceptor    Strength of Acids and Bases  ­ based  on ability to dissociate in water, the more dissociated = stronger  ­ know  table of strong acids and bases on slide 147 • Measuring Acidity pH ­     is the numerical measure of acidity or basicity  8 Friday, May 20, y • pH = ­log [H O 3 + the ­  bracketing means the “concentration of”  + ­ • a pH of 7 is neutral so [H O ]3= [OH] ­ i pH < 7 the solution is acidic so [H O ] 3 [OH] ­ ­ i pH > 7 the solution is basic so [H O ] < [OH] ­ 3 + •    Conjugate Acid­Base Pairs ­ molecules or ions that differ by a single proton (H ) ­ acid  contains the proton transferred to a molecule/ion ­ conjugate  base is the molecule/ion remaining after the loss of the proton  • acid + base = conjugate base + conjugate acid  ­ know  how to recognize what is what in an equation    Buffers mixture ­  of weak acid + conjugate base that resists pH change  • can categorize a weak acid by its K  “acidadissociation constant” ­ Ka < 1 weak  (the smaller the K athe weaker the acid) • pK = ­logK (way to express really small Ka values) a  a   ­ Ka > 1 strong ­ an  ideal buffer • pH = pK + a g( [A ]/[HA] ) ideally ­ , there will be equal amounts of A­ and HA because the log(1)=0 and  the pH will equal the pK a • provides max buffering capacity  9


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