Chem-is-try Review Final ExamCh 8: Basic Concepts of Chemical BondingLewis Dot Diagrams1. Sum the valence electrons from all of the atoms.2. Connect the appropriate atoms using a single bond (a line, which represents two electrons). The central atom is usually the least electronegative atom, but never hydrogen.3. Add lone pairs to complete the octets of the outer atoms. Keep in mind that hydrogen fills its "octet" with only two electrons.4. Place any leftover electrons on the central atom. Keep in mind that elements in row 3 of the periodic table and beyond can exceed the general octet.5. If any atom lacks an octet, try using double or triple bonds. You may have to use one or more lonepairs to make the double or triple bonds so that the total number of valence electrons remains the same.Exceptions to Octet:- Be and B uses its valence elections as bonds, so B often forms 3 bonds and Be 2.- Elements in the third row of the periodic table and beyond can have more than an octet of electrons in a covalent compound due to the existence of an empty d subshell available to these elements, which allows them to expand their valence to a number greater than eight.Formal Charge:- FC= valence electrons of element-lone pairs-bonds.Non/Polar Covalent and Ionic Bonds:Nonpolar Covalent- When two bonded atoms attract electrons with equal strength.Polar covalent- When electrons are unequally shared between the atoms. (The more electronegative element is negative)Ionic- When the sharing is so unequal that fully charged ions form.- Electronegativity difference can be used to predict bond type. - If differ by more than 2 units, the bond is substantially ionic.- If they differ by less than 2 units, the bond is polar covalent. - If the values are equal, the bond is nonpolar covalent.- Can still predict the bond type using the periodic table. (Know trend)- Metals have low electronegativity compared to nonmetals so any metal–nonmetal combination will be ionic.- Any nonmetal–nonmetal combination will be covalent. - Without electronegativity you should assume that a covalent bond is polar unless it is between two atoms of the same element.Resonance Structures: Alternate versions of the same molecule.Bond length: Depends on how many bonds are with another atom. The more bonds the shorter it is.Ch 9: Molecular Geometry and Bonding TheoriesHybridization- The Lewis structure for a molecule allows you to predict its shape, bond angles, and hybridization based on the number of charge clouds on the central atom.Clouds Shapes Angle (degrees) Hybrid orbitals2 linear 180 sp3 trigonal planar, bent 120 sp24 tetrahedral, trigonal pyramidal, bent 109.5 sp35 trigonal bipyramidal, seesaw, T-shaped 90, 120, 180 sp3d6 octahedral, square pyramidal, square planar 90, 180 sp3d2- The magnitude of repulsion between two lone pairs is higher than the magnitude of repulsion between a lone pair and a bonded pair of electrons.Molecular Geometry and Hybridization:With hybridization count how many elements are attached to another (not the bonds) and that will help. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/table.htmlhttp://www.stolaf.edu/people/jackson/08-125/molshape/molshape.htmMolecular Polarity (Dipole moment):- A molecule can be polar or nonpolar depending upon the nature of the bonds and the shape of the molecule. For a molecule that has different outer atoms the molecular symmetry will decidethe polarity.- If the molecular geometry is such that the dipole moments of each polar bond cancel each otherthan the molecule is nonpolar.- However, if the molecular geometry is such that the dipole moments of each polar bond don’t cancel each other than the molecule is polar.Sigma and Pi Bonding:- Overlap occurs between s and sp2 (H and C), sp and sp2 (C andC), sp and sp (C and C), as well as p and p. - Each Pi bond results from sideways overlap of 2 p orbitals. Allother overlaps are head on resulting in sigma bonds.-Chemical bonds have 1 sigma bond, the rest are Pi.Single Bond- 1 sigmaDouble Bond- 1 sigma, 1 PiTriple Bond- 1 sigma, 2 PiPi Orbital (bond)- Side/side overlap of atomic orbitals. Electron density above and below the bond axis.Sigma Orbital (bond)- Electron density found along axis of bond.Electron Domain- The shape created from all bonds including lone pairs. Molecular Geometry- The shape of the molecule from the electron domain but without the lone pairs.Delocalization- Electrons belonging to a certain molecule but not attached to a particular atom/bond in molecule. Ergo delocalized since they do not have a particular location and exist in orbitals, clouds surrounding the molecules suggesting the space the electron could be in. General the Pi orbitals.Ex. In O3 (ozone) there are 4 delocalized electrons.Ch 11: Liquids and Intermolecular Forceshttp://www.chem.purdue.edu/gchelp/liquids/hbond2.htmlIntermolecular forces: - If increased the bonds are stronger ergo higher melting point, surface tension, viscosity, and boiling point while decreased vapor pressure. - An increase in number of bonds in a straight line increases intermolecular forces from one end to another, however if another molecule has O2, its high electro negativity means stronger intermolecular forces. - Anything with stacked methyl groups decreases the intermolecular forces. Ex. 2,2-dimethylbutane. - If adding an O to the entire molecule London Dispersion Forces- The weakest attraction that occurs between any 2 adjasent atoms polar or non-polar. Electrons constantly move so create asymmetrical distribution, ergo temporarydipoles, stronger as atom gets larger. + attracted to -. Dipole-Dipole- Between + end of one polar molecule and the - end of another polar molecule. Weak so only have a significant effect only when the molecules involved are close together.Ex. The more electronegative Clbears a partial – and I partial+, ergo attraction.Ion-Dipole- Results from an electrostatic attraction between an ion and a neutral molecule that has a dipole. Stronger if ion charge increases or magnitude of dipole increase. Overall strong.Ex. Cation to partially – end of a neutral polar mol, Anion to partially + end of a neutralpolar mol.Hydrogen Bond- Special dipole-dipole attraction, not covalent bond to hydrogen atom but very strong attractive force between very electronegative atoms N,O, or F. Molecules containing N-H, O-H or F-H bonds have a large diff in electronegativity between the H