VIEWS: 14 PAGES: 12 POSTED ON: 9/18/2011
Chapter 8 & 9: Bonding 1 Chapter 8 and 9 Chemical Bond: Bond Energy: Valence Electrons: Electronegativity: ability of an atom in a bond to __________________ Ionic Bond: Attraction of oppositely charged ions Ionic Compound: Coulomb’s Law: E = (2.31 x 10-19 J nm) (Q1Q2/r) Monatomic ions: Polyatomic ions: Active metals tend to lose electrons to obtain previous noble gas configuration. Nonmetals tend to gain electrons to obtain next noble gas configuration. Isoelectronic: Electronegativity Difference (EN): determines how likely two elements are to form an ionic compound. : large EN, ionic compound is likely (approximately 1.7) : Cs 0.8 F 4.0 EN =3.2 ionic Lattice Energy: the change in energy associated with the formation of an ionic solid from gaseous ions. Chapter 8 & 9: Bonding 2 Example: Use the following data to estimate Hfo for magnesium fluoride. (-2088 kJ/mol) Lattice Energy: -3916 kJ/mol First Ionization Energy of Mg 735 kJ/mol Second Ionization Energy of Mg 1445 kJ/mol Electron affinity of F -328 kJ/mol Bond energy of F2 154 kJ/mol Enthalpy of sublimation of Mg 150. kJ/mol Covalent Bonding Covalent Bond: two atoms sharing one or more pairs of electrons (small EN values) Lewis Formula: shows all valence electrons in an atom, ion or molecule Atoms: Atoms in molecules tend to have noble gas electron configurations (share in 8 electrons)-octet rule or rule of eight. Molecules: H2 O2 N2 Single covalent bond: sharing of Double covalent bond: sharing of Triple covalent bond: sharing of Bond Energies (B.E.) The bond energy is the energy needed to break one mole of bonds in a gas substance to form 1 mole of products in the gas state at constant T and P. H2(g) 2H(g) Ho = 432 kJ/mol O2(g) 2O(g) Ho = 495 kJ/mol When bonds are broken, H is always When bonds are formed, H is always Calculation of Ho using bond energies: Ho = broken - formed Chapter 8 & 9: Bonding 3 Example: Calculate Ho for the reaction, 2H2(g) + O2(g) 2H2O(g) using B.E. values from Table 8.4 (page 351). Ho = 2BE(H-H) + BE(O=O) – 4BE(HO) Ho = Example: C2H4(g) + F2(g) C2H4F2(g) Diatomic Elements: BrINClHOF Unshared electrons: Writing Lewis Structures (Electron Dot Structures) 1. Count number of valence electrons. (NH3) 2. Draw a skeleton structure joining adjacent atoms with a single bond. 3. Arrange remaining electrons to satisfy octet for each atom (except H). *Central atom is usually written first in formula. *Terminal atoms are usually the most electronegative. NH4+ ClO3- Cl2CO HCN Resonance: Sharing of a bonding pair of electrons between more than 2 atoms. NO3- Chapter 8 & 9: Bonding 4 SO2 Exceptions to the Octet Rule 1) Be compounds (BeCl2) 2) B compounds (BBr3) 3) Odd # of valence electrons (NO2) 4) Expanded Octets (SF6, atoms with available d orbitals) Formal Charge: difference between the number of valence electrons in the atom and the number of electrons assigned to the atom -used to determine most accurate Lewis Structure in cases where more than one possible structure exists. Formal Charge = Group Number – (number bonding pairs + nonbonding electrons) Polar and Nonpolar Covalent Bonds Nonpolar covalent bonds: Polar covalent bonds: H-F H-Cl H-Br H-I EN: 2.1-4.0 2.1-3.0 2.1-2.8 2.1-2.5 EN: 1.9 0.9 0.7 0.4 Most Polar Least Polar Electric dipole: separation of charge Unequal sharing of electrons leads to + and - Chapter 8 & 9: Bonding 5 Properties of Ionic versus Covalent Compounds Ionic Compounds Covalent Compounds 1. Solids (High melting points >400oC) 1. Gases, Liquids, and Solids 2. Generally soluble in polar solvents (water) 2. Generally insoluble in polar solvents (water). 3. Most are insoluble in nonpolar solvents. 3. Most are soluble in nonpolar solvents. 4. Molten compounds conduct electricity. 4. Liquids and molten compounds do not conduct electricity. 5. Aqueous solutions conduct electricity. 5. Aqueous solutions are usually poor conductors of electricity. Naming Compounds Naming Ionic Compounds Binary- Simple Binary Ionic Compounds 1. 2. 3. Binary Ionic Compounds with Variable Charge Cations The charge on the cation must be specified in the name of the compound The charge on the cation is determined from the charge on the anion Chapter 8 & 9: Bonding 6 Ionic Compounds with Polyatomic Ions 1. Name cation 2. Name anion potassium fluoride lithium phosphate Cd3(PO4)2 Al2(SO4)3 Naming Binary Covalent (Molecular) Compounds CO CO2 *Name first element, name 2nd element with “ide” ending. Use prefixes. Prefixes mono di tri tetra penta hexa hepta octa nona deca NO2 N2O4 Recognizing and Naming Acids (page 66) Acid: substances that produce H+ when dissolved in water Formulas: Acids without oxygen: Acids with oxygen: Chapter 8 & 9: Bonding 7 Lewis Structures: Major features of molecular geometry can be predicted based upon electron pair repulsion. VSEPR represents Valence Shell Electron Pair Repulsion Bonding and lone pairs of electrons are regions of High Electron Density (HED) Ideal Geometries Two electron pairs around a central atom. Example HED Geometry Bond Angles BeCl2 Three electron pairs around a central atom. Example HED Geometry Bond Angles BCl3 Four electron pairs around a central atom. Example HED Geometry Bond Angles CH4 Five electron pairs around a central atom. Example HED Geometry Bond Angles PF5 Six electron pairs around a central atom. Example HED Geometry Bond Angles SF6 Chapter 8 & 9: Bonding 8 Effect of Lone Pairs on Geometry 1. Electron pair geometry is relatively unaffected by lone pairs 2. Molecular geometry is different than EG when lone pairs are present. Example HED EG MG Bond Angles CH4 Example HED EG MG Bond Angles NH3 Example HED EG MG Bond Angles H2 O Polarity of Molecules Nonpolar covalent bond: equal sharing of electrons Polar Covalent bond: unequal sharing of electrons Polar molecules contain positive and negative poles (dipole) Two Factors Determine Polarity of a Molecule 1. bond polarity 2. molecular geometry Example: BCl3 Example: BCl2F Example: CH2Cl2 Example: CH2O VSEPR: electrons repel each other for maximum separation Valence Bond (VB): ½ filled orbitals overlap to form covalent bonds Chapter 8 & 9: Bonding 9 Why does O typically form two bonds? Why does N typically form three bonds? Example HED EG MG Bond Angles Polarity Hybridization BeH2 Lewis structure Hybridization Example HED EG MG Bond Angles Polarity Hybridization BH3 Lewis structure Hybridization Example HED EG MG Bond Angles Polarity Hybridization CH4 Lewis structure Hybridization Example HED EG MG Bond Angles Polarity Hybridization PCl5 Lewis structure Hybridization Example HED EG MG Bond Angles Polarity Hybridization SF6 Lewis structure Hybridization Example HED EG MG Bond Angles Polarity Hybridization IF7 Chapter 8 & 9: Bonding 10 Lewis structure Hybridization Example HED EG MG Bond Angles Polarity Hybridization NH3 Lewis structure Hybridization Example HED EG MG Bond Angles Polarity Hybridization AsF4- Lewis structure Hybridization Example HED EG MG Bond Angles Polarity Hybridization SeF5- Lewis structure Hybridization Example HED EG MG Bond Angles Polarity Hybridization BrF3 Lewis structure Hybridization Example HED EG MG Bond Angles Polarity Hybridization KrF4 Lewis structure Hybridization Multiple Centers Examples and Models Chapter 8 & 9: Bonding 11 Multiple Bonds Examples, Models, Overheads C2H4 C2H2 CH2O Valence Bond Theory (Localized Electron Model): Overlap of ½ filled hybridized or pure atomic orbitals leads to bonding. (hybridization is the mathematical combination of atomic orbitals on a specific atom) Molecular Orbitals (MO): Atomic orbitals on separate atoms combine to form MO. (molecular orbitals are mathematical combinations of orbitals on separate atoms, no atomic orbitals remain once molecule has formed) VB theory is easier to visualize (VSEPR) MO theory gives better description of electron cloud, magnetism and bond energies. Electrons can be treated mathematically as waves -waves added in phase interact constructively + = -waves added out of phase (subtracted) interact destructively + = Bonding Orbital: Results when two orbitals interact in phase (electron density between atoms, lower energy than atomic orbitals). Antibonding Orbital: Results when two orbitals interact out-of-phase (electron density shifts to unfavorable regions, higher in energy than atomic orbitals). bonding e' s antibonding e' s bond order 2 Chapter 8 & 9: Bonding 12 Increasing bond order correlates with increasing bond stability and decreasing bond length. Bond Order Bond Length (Å) Bond Energy (kJ/mol) N2 3 1.09 946 O2 2 1.21 498 F2 1 1.43 159 Magnetic Properties Paramagnetic substances-contain unpaired electrons. -Weakly attracted by magnetic fields. Diamagnetic substances-contain no unpaired electrons. -Weakly repelled by magnetic fields. Ferromagnetism- (Fe, Co, Ni) permanent magnets -Electrons in adjacent atoms align and generate a magnetic field.
Pages to are hidden for
"Chapter_8___9"Please download to view full document