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Exam 2 Review.docx

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Department
Chemistry
Course
CAS CH 109
Professor
Sean Elliott
Semester
Fall

Description
Chapter 3: Chemical Bonds • Valence electrons are based on the group number of an element (ex: C is in group 4A, 4 valence electrons) • Ionic Bonding – 2 ions held together by electrostatic attraction • Only metals have low enough ionization energy and only non-metals have high enough electron affinity for the formation on a crystal  Lattice Energy • IE + (-EA) = ∆E • Potential Energy = (q q )1 2πεr)  q = C1arge on cation, q = Charg2 on anion, r = sum of ionic radii + - • Ionization Energy = x(g)  x (g) + e(g) • ElectronAffinity = x(g)  x(g) + e(g) or Negative of x(g) + e  x(g) - - • Energy from Transfer of Electrons  IE + (-EA) 2 z1 2e • ***Potential Energy of Ions (Ions to gas)  4πεd • ***Energy of Formation (Gas to gas)  Energy for transfer of electrons + Potential Energy of Ion Interactions • ***Energy from Formation (Gas to solid)  Energy for transfer of electrons – Lattice • Lattice Energy/Potential Energy of Ions = Madelung Constant • Lattice Energy = Ions to solid • Energy < 0  Stable, Energy > 0  Unstable • Trends  Metals are easier to ionized, non-metals have strong electron affinity • In solids, attractions/repulsions of all neighboring particles must be taken into account • Lattice Energy  Difference in energy between ions of a compound widely separated as gas and pack together in solid • Lattice energy of ionic solid is large when ions are small and highly charged −AN Z ZAe 1 2 2 • Lattice Energy =  Z 1nd Z ar2 absolute value of charges of atoms 4πεd • Greater charge on ion = Greater lattice energy oLarge ionic radii = Lower lattice energy • Formal Charge = Number of Valence Electrons – (bonding electrons + lone pair electrons) • Individual formal charges that are equal to zero, or close to zero indicate most likely arrangement of atoms • Lowest formal charge = Most stable structure  Predominant structure of molecule oNegative charge should be on atom with highest electronegativity oAdjacent atoms should not have same sign oC favors 4 bonds, O favors 2 bonds, N favors 3 bonds, H only forms 1 bond • Lattice Energy is proportional to Madelung Constant (A) • Atoms in Period 3 or Higher can expand their valance shell capacity (hypervalency) • CentralAtom =Atom with lowest ionization energy = Highest Electronegativity • Double-headed arrow indicates non-equivalent resonance structures • High Electron Affinity + High Ionization Energy = High Electronegativity • Low-energy structures contribute more to resonance mixture than high-energy structures +3 3+ • B andAl can have incomplete octets • Radicals = Species having electrons with unpaired spins • Equal sharing  Nonpolar covalent • Non-equal sharing  Polar covalent • Ionic Character  Higher if transfer of electrons is complete oGreater EN difference = Greater ionic character • Dipole Moments Arrow points from positive to negative • Polarizability =Ability of ions to distort electron clouds of other atoms oLarge anion  Highly polarizable (soft + large) 3+ oHigh polarizing power = Cation (small and strongly charged likeAl ) that has ability to distort electron cloud of anion oPolarizing power increases from left to right of table oPolarizing power decreases from top to bottom of table • Bond length increases as atoms become bigger • Bond length decreases as bond order increases • Strong bonds  Shorter + Stiffer • Multiple bonds cause bond dissociation energy to increase • Dissociation Energy = Energy required to separate bonded atoms  High dissociation energy indicated deep potential energy well and a strong bond • High dissociation energy = More energy necessary to break a certain bond (Strong bond) • Multiple bonds are stronger than single bonds  Need more energy to break them • Lone pairs can decrease strength of a bond  Repulsion • Bond strength decreases as atomic radi
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