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Exam 2 Study guide

by: Poshita Nigam

Exam 2 Study guide CHM 115-06

Poshita Nigam

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About this Document

These notes cover everything that could be on our next exam
Principles of Chemistry I
Dr. Heidi Cuticchia
Study Guide
CHM 115, Exam 2, midterm, unit 2, lewis structures, Chemical Bonds, ionic, covalent, octet, electronegative, empirical, Molecular
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This 6 page Study Guide was uploaded by Poshita Nigam on Monday February 29, 2016. The Study Guide belongs to CHM 115-06 at Grand Valley State University taught by Dr. Heidi Cuticchia in Winter 2016. Since its upload, it has received 47 views. For similar materials see Principles of Chemistry I in Chemistry at Grand Valley State University.


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Date Created: 02/29/16
Monday, February 29, 2016 Exam 2 Study Guide Experimental Evidence for Compounds—Moles - Determine Percent Composition of a compound, given its formula • % mass of element = (mass of element / total mass of compound) x 100 - Determine Empirical and Molecular Formulas from data • Empirical formula: subscripts in compound in lowest possible terms • *doesn’t actually indicate the actual number of atoms in the compound • Step 0: Given % Mass of elements • Step 1: Assume that the given is 100g of the pure compound. By doing so, your percents are now grams of elements • Step 2: Convert elements from grams to moles • Step 3: Divide all elements by the smallest moles (Mole Ratios) • Step 4: Empirical Formula • Molecular Formula: A factor larger by mass than the empirical formula • Step 5: Determine Empirical formula mass. • Step 6: Divide the molecular mass (often given) by the empirical mass to determine the factor. • Step 7: Multiply the empirical formula by the factor to get the molecular formula. If empirical formula is Ax By then molecular formula is A(nx) B(ny) • 1 Monday, February 29, 2016 Chemical Bonds and Classifications of Bonding - Combustion Analysis: combustion reaction; unknown compound undergoes combustion (burning) in the presence of pure oxygen. (see page 175 figure 5.10) - Ionic Bonding: electrostatic attraction between a positive cation and a negative anion. • electron transfer from metal (cation) to a nonmetal (anion) • total # of e-’s lost by cation(s) = total # of e-’s gained by anion(s) • Lattice Energy: energy required to form a crystalline lattice from its gaseous ions. Measurement of how strong the electrostatic attraction is in an ionic compound - depends on ionic size and ionic charge Naming Ionic compounds: the first element (metal cation) retains its name. The • second element (nonmetal anion) has an adjusted ending with “ide” added. - NaCl2: sodium chloride - CaO: calcium oxide - Polyatomic Ions: know them. See page 159 Table 5.5 - “ate” ending - “ite” ending: one less hydrogen than the “ate” - SO4: Sulfate —— SO3: Sulfite - Acids: have only H+ ion as only cation. Only 2 elements. - hydro= H+ - “ic” ending for the ending and then add “acid” to the end. - HF: hydrofluoric acid - HCl: hydrochloric acid - Know: HF, HCl, HBr, HI, H2S 2 Monday, February 29, 2016 - Oxyacids: acids from polyatomic ions (anions) - cation: H+ - if “ate” anion then the acid ending is “ic” - if “ite” anion then the acid ending is “ous” - H2CO3: carbonic acid - HNO2: nitrous acid - Covalent Bonding: sharing electrons • only nonmetals • to the right of metaloid line: covalent • to the left of metaloid line: ionic • Naming: add prefixes according to number of atoms. - N2O: dinitrogen monoxide - P4O10: tetraphosphorus decoxide - Electronegativity: ability of an atom to attract a pair of electrons in a chemical bond • Pauling Scale: 0-4 (fluorine) • Difference in Electronegativity Scale: 0 to 0.4 is polar covalent. >0.4 to 2.0 is polar covalent. >2.0 is ionic • Polar covalent: electrons not quite equally shared • Nonpolar covalent: equally shared electrons - Valence Electrons: those electrons in outer most energy level that are used for bonding - Duet Rule: give, receive, or share to have 2e-’s (helium) - Octet Rule: give, receive, or share to have 8e-’s (s and p orbitals only) 3 Monday, February 29, 2016 Reactions - Synthesis/Combination Reaction: forms only 1 product • Diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2 - 4Al + 3O2 —> 2Al2O3 - Decomposition • One reactant decomposes - chlorate —> metal chloride + oxygen gas - carbonate —> metal oxide + carbon dioxide - Single Replacement Reaction • Element + ionic compound —> new element + new ionic compound • Three types: metal; active metal reacting with and acid; nonmetal - Double Replacement Reaction Precipitation Reaction: solid forms as a product • • Solubility rules: page 312 Lewis Structures - Determining Lewis Dot Structures • Step 1: Determine # of valence electrons each atom contributes and calculate total for molecule - 1A: if you have cation, then subtract # of electrons that correspond to the charge - 1B: if you have anion, then add # of electrons that correspond to the charge Step 2 : Write down central atom (usually single atom in the formula) and single • bond all other atoms/ terminal atoms to it. • Step 3 : Give full octet (3 pairs of electrons) to all terminal atoms, except hydrogen • Step 4 : Count electrons in structure, subtract from total # of electrons (step # 1) 4 Monday, February 29, 2016 - add remaining electrons in pairs to central atoms • Step 5 : Check central atoms fro 8 electrons. If not then double/triple bond. *except boron • Step 6 : Calculate formal charge for each atom in the structure. - formal charge= # of valence electrons - # of electrons in lone pairs - 1/2( # of electrons in bonds) - Good structure has a formal charge of 0 - most EN atom carries the negative formal charge - Polyatomic: charge of ion = sum of the formal charges - Resonance: more than one plausible (based on formal charges) structure - Multicenter Molecules: all atoms have zero formal charges, except when its an ion. - Electron Deficient: can have a maximum of 6 electrons - Expanded Octet: can have up to 16e-’s but no more than 6 groups of electrons on central atom • Enthalpy of Reaction (ΔH) bond energies: energy required to break one mole of bond. (Given in a table) • • ΔH: change in energy • coefficient x # of bonds of one type x bond energy • sum of reactants - sum of products • ΔH < 0 : exothermic : bond making/releasing energy • ΔH > 0 : endothermic : bond breaking absorb energy • VSEPR Theory: Valence shell electron pair repulsion 3-D structures 5 Monday, February 29, 2016 • Strongest Repulsion: Lone pair- Lone pair • Medium Repulsion: Lone paid- bond pair • Weakest Repulsion: Bond pair- bond pair 6


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