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CHEM 1331- Chemistry 2 Exam 2 Study Guide/Review

by: Alexis Clowtis

CHEM 1331- Chemistry 2 Exam 2 Study Guide/Review CHEM 1331

Marketplace > University of Houston > Chemistry > CHEM 1331 > CHEM 1331 Chemistry 2 Exam 2 Study Guide Review
Alexis Clowtis
GPA 4.0

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About this Document

This study guide is a mixture between the lectures, the review Professor Teets posted, and homework review videos. The material covered is chapters 4-6.
Fundamentals of chemistry
Thomas Teets
Study Guide
Chemistry, review, molarity, stoichiometry
50 ?




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This 51 page Study Guide was uploaded by Alexis Clowtis on Wednesday March 30, 2016. The Study Guide belongs to CHEM 1331 at University of Houston taught by Thomas Teets in Spring 2016. Since its upload, it has received 253 views. For similar materials see Fundamentals of chemistry in Chemistry at University of Houston.


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Date Created: 03/30/16
CHEM 1331 Exam 2 Review Chapter 4-6 Chapter 4 VSEPR • VSEPR: Valence Shell Electron Repulsion: Place the electrons around the central atom as far away from each other as possible. • Electron- group arrangement: 3-D arrangement of all electrons • A 3-D geometric shape that the bonding and nonbonding electrons occupy • Molecular Shape: the 3-D shape that only the bonded atoms occupy i.e. what the molecule “looks like” • AX E Notation: • A= central atom • X= any atom bonded to central atom • E= lone pair on central atom • m&n= tell you how many of each you have • Steric Number: m+n- the number of electron groups, Note: Bond order between A&X doesn’t really matter X counts as 1 electron group. • angleangle: The angle between two bonded atoms, in which the central atom is the vertex of that Two Electron Groups AX2– Linear – Bond angle= 180 AXE – Linear – No bond angle (because only one bonded atom) Three Electron Groups AX3 – Trigonal Planar – Bond angles=120 AX2E – bent – Bond angles<120 Four Electron Groups AX2E2 – bend- Bond angle<109.5 Five Electron Groups AX5 – trigonal bipyramidal – Bond angles=90 and 120 AX4E – see saw AX3E – T-shaped – Bond angles<90 AX2E3 – linear – Bond angles=180 Six Electron Groups AX6 – octahedral – bond angles=90 AX5E- square pyramidal – bond angles<90 AX4E2 – square planar – Bond angles=90 Molecular Polarity • Uneven distribution of charge over an entire molecule not just a single bond • Must have polar bond • Individual bond dipoles must be asymmetrically distributed Found on studyblue A=A B=X L=E compounds&bvm=bv.118353311,d.amc&psig=AFQjCNGrCtFrWwK6z_mU00nS0Jx8vgFd1g&ust=1459441673508014 Hybridization • Sp- AX2 • Sp2- AX3 • Sp3- AX3E • Sp3d- AX5 Predicting hybridization • Steric number Types of bonds • Sigma- bonds formed by end to end overlap of atomic or hybrid orbitals • Electron density is highest between nuclei along the bond axis • Hybrid orbitals only form sigma bonds • Pi-side to side overlap of orbitals • Results in regions of electron density above and below • Unhybridized p orbitals only Valence Bond Theory • All single bonds are sigma bonds • Multiple bonds consist of one sigma bond and the rest are pi bonds Molecular Orbit Theory • Atomic orbitals are mathematically combined to form delocalized molecular orbits • Taking everything and mixing it, not just gluing together • Only orbitals that are close • Valence orbitals Types of Orbitals (MO) Bonding Orbitals Antibonding Orbitals ∗ • Wave function: ???? = ???????? +1???????? 2 • Wave function: ???? = ???????? − 1 • Additive combination ???????? 2 • Electron density: greater along • Subtractive/out of phase combination bond/between nuclei • Energy: lower in energy than • Electron density: greater outside atomic orbitals internuclear region • Energy: higher than the AO’s Molecular Orbitals Parameters determined from molecular orbital diagrams 1 • Bond order2 # ???????? ???????????????????????????? ???????????????????????????????????? − Filling orbitals with electrons Fill from bottom to top (s to p) Use electron configuration to find out how many electrons to put on Still follow Hund’s Rule Chapter 5 Atomic Mass • 1 amu=mass of carbon 12  12 • All atomic masses are measured relative to the standard, carbon 12 • Weighted average of all isotope masses • FractionAXMassA + FractionBXMassB… • Not exactly equal to mass number because of ???? = ???????? 2 • A lot of energy is released when nuclei are formed which changes the mass Percent Abundance • Relative amount of each isotope in a natural sample of that element • Occurrence in nature of each isotope basically • use to find average atomic mass (what’s on periodic table) The Mole • There are 6.022 × 1023 “things” in a mole • Things= atoms, particles, etc. • Counting number- like dozen, there are 12 “things” in a dozen no matter what the things are • Atomic mass of element is how many grams are in 1 mole of that element Molar Mass • Mass of 1 mole of a compound (like atomic mass for an element on the periodic table but it’s a compound so you add all of the individual atoms/element’s atomic masses together to find molar mass of compound) Differences between Molecular and Empirical formulas (don’t have to worry about difference for ionic compounds because already in lowest whole number ratio unlike covalent compounds) Molecular Formula Empirical Formula • Total number of each atom in • Lowest whole number ratio of that compound (can’t take atoms in a compound whole number ratio for covalent • What you can find given mass % compounds) • Gives structure • Molecular mass=empirical massXsmall whole # Reactions • Combustion- C + H + other stuff= CO2 +H2O + other stuff • All carbon is converted into CO2, all hydrogen into H2O • Precipitation reaction • Acid-Base Reactions • Recombination of hydrogen and hydroxide • Redox reactions • Titration Chemical Equations • Reactants (left)- compounds consumed • Products (right)- compounds formed • Coefficients- numbers in front of each symbol that tell you number of moles of each species involved • Balance: 1. Arrange reactants and producs 2. Add coefficients 3. Adjust coefficients 4. Check 5. States of matter Stoichiometry • Calculating the amounts of reactants and/or products involved in a chemical reaction • Allows you to predict amount of product • Mole ratios given by coefficients in balanced chemical equation • Stoichiometric ratio ( R) ONLY CAN BE USED TO DETERMINE LIMITING REACTANT Limiting Reactant • Use stoichiometric ratio ( R)= Moles of reactant/reaction coefficient • Smaller ratio is limiting reactant • Other one is in excess • After finding which is limiting, use the mass given for that compound to determine mass/moles/whatever for product it’s asking for Percent Yield ???????????????????????? ???????????????????? ????100 = % ???????????????????? ????ℎ???????????????????????????????????? ???????????????????? Actual is given in the problem (when someone does the experiment) Theoretical you have to find based off of the values given in the problem (what you would ideally/theoretically get if the reaction happened 100% perfectly) Chapter 6 Definitions Polarity of Water • Water is polar because: • O is more electronegative than H • H2O has a bent shape- bond dipoles do not cancel out • Negative and positive “poles” (regions that are negative and others that are positive) Solutions of Ionic Compounds in Water • Ionic compounds separate in solution • Conduct electricity in solution  electrolyte solution • Covalent compounds: no bond breaking Electrolytes • Strong electrolytes are soluble in ionic compounds or strong acids • Conduct electricity well • Weak electrolytes- compounds partially break down into ions and weakly conduct electricity • Weak acid • Nonelectrolytes- do not conduct electricity • Covalent compounds that are not acids or bases Molarity ???? = ???????????????????? ???????? ???????????????????????? ???????????????????????? ???????? ???????????????????????????????? -Concentration unit -Can be used as conversion factor Dilution M1V1=M2V2 =moles of solute Dilution is the solution! Moles of solute stays the same before and after dilution Higher M, smaller V Lower M, larger V Precipitation Reaction • 2 solutions combined to form 1+ insoluble products • Determine if precipitate will form: 1. Determine ions in R’s (separate reactants into constituent ions) 2. Consider all combinations- usually 2 products are possible 3. Decide whether any combination is insoluble- Solubility rules Equations for Aqueous Ionic Reactions • Molecular (formula) equation: the full set of reactants and products written as intact species • Total ionic equation: separate all SOLUBLE ionic compounds into their ions • Net ionic equation: remove “spectator”(aq) ions which appear on both sides of the equation Acide-Base Reactions • Acid- compound that releases H+ when dissolved in water • Formula: ???????????? • n=number of hydrogens=anions charge • X=anion • Base- compound that releases OH- in water • Formula: ???? ????????) ???? • n= charge on ????????+cation • Strong acids/bases- dissociate completely into ions they’re composed of (strong electrolytes) • EBases- soluble hydroxides (LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2 • Weak acids/bases- do not dissolve completely but release some ions so most of the molecule stays intact • Ex. Acids- HF, H3PO4, CH3COOH (acetate) Bases- NH3 • Acid + Base= Water + Salt (recombination of H and OH into water) Recombination of H+ and OH- • Net ionic equation: H+(aq) + OH-(aq) = H2O(l) • Spectator ions and coefficients depend on identity of acids and bases ions Redox Reaction • Always has Oxidation and reduction “happening” • Oxidation- losing electrons LEO • Reduction- gaining electrons GER • If you have a chemical by itself in a chemical reaction then its almost definitely a redox reaction • Oxidation number rules (somewhat related to EN- O.N. gives all electrons to more EN atom) • Balance Redox reactions based off of oxidation number changes Titration • using solution of known concentration to determine unknown concentration • Acid/base- a solution of base is added to acid or vise versa • Equivalence/end point of titration: point where moles of acid(or base) added equals the moles of the other initially present • Completely reacted to form water Review Question T opics • Electron-group arrangement and molecular geometry for given molecule • Hybridization of an atom • Molecular orbital theory/diatomic molecule • Bond order • Percent abundance • Moles of an ion in a compound given mass of compound • Empirical formula given percent by mass of ions • Combustion empirical formula • Mass of product in combustion • Moles of product given mL of both reactants • Percent yield • Equation that forms a precipitate • Concentration of ions in M • Volume to neutralize other reactant given mL and Molarity • Oxidation number • Balance equation with ions • Net ionic equation Extras • Dipole moment- Do dipoles cancel out or add together?Polarity • If cancel out, nonpolar • If don’t cancel out, polar (with dipole moment in direction that vector addition would give you) • C-H is essentially nonpolar • If there is ONE lone pair on central atom, then it will be polar • If there are 0 lone pairs on central atom, it will be nonpolar -Pi bonds don’t have any rotational freedom


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