:in the early 1900's, an important proposal was made concerning bonding. the uniqueness in the electro configuratin of noble gas atoms accounts for their inertness, and atoms of other elements combine with one another to acquire the electron cofiguration of a noble gas. :the lewis theory developed from this idea: : 1.electrons, especially those of the outermost shell (valence shell) play a role in bonding. : 2.in some cases, electrons are transformed from one atom to another. positive and negative ions form and become attracted to each other forming ionic bonds. : 3.in other cases, 2 or more electrons are shared between atoms. the sharing of electrons forms a covalent bond. (metals and nonmetals) : 4.electrons are transferred or shared in a way that each atom acquired a complete electron cofiguration (usually a noble gas configuratin with 8 valence electrons-called an octet). :a lewis symbol consists of a chemical symbol to represent the nucleus and core electrons. dots are placed to represent the valelce electron. :lewis did not know anything abour spin so the dots are placed around an element before pairing up begins. : lewis structures for: :N, P, As, Sb, Al, I, Se, Ar, Ca, O, Ne, B, Ga, Po :lewis structures are a combination of lewis symbols that reprsent the transfer or sharing of electrons. :a covalent bond is the result of electron sharing :the sharing of a single pair of electros between atoms produces a single covalent bond. the term bonding pair is used to represent the electron in the bond. the term lone pair represents the electrons not involved in bonding. :the 3 pair of electrons in n2 is called a triple bond and o2 has a double bon bond order describes the type of covalent bond: single, double, or triple :bond length is the distance between the centers of 2 atoms joined by a covalent bond a triple bond is shorter than a double bond is shorter than a single bond. :a covalent bond in which electrons are not shared equally is called a polar covalent bond. in such a bond, electrons are displaced toward the more nonmetallic element. symbols + and - indicate where the partial positive and partial negative charge are. :subtract electronegativities: : 0-.3=nonpolar covalent : .3-1.7=polar covalent : 1.7-4.0=ionic :writing lewis structures : 1. all the valence electrons of the atoms in a lewis structure must appear in the structure : 2. usually all the electrons are paired. : 3.usually each atom acquires an outer-shell octet of electron. hydrogen is limited to 2 outer shell electrons. : 4.sometimes multiple covalent bonds are needed. :rules: : 1.determine the total number of valence electrons. : 2.write the skeleton structure and joi the atoms in thsi structure by single covalent bonds. : 3.for each single bond, subracct 2 fro the total number of valence electrons. : 4.with the valence electrons remaining, first complete the octets of the terminal atoms. : 5. at this point, if the central atom lacks an octet, form miltiple covalentbods by converting lone pairs from terminal atoms into bonds. :draw structure for c,h,n together :polyatomic ions cosist of 2 or more atoms and the forces holding atoms together within such ions are covalent. :ionic compounds formed from polyatomic ions have both covalent and ionic bonds. :a coordinate covalent bond is a covalent bond in which one atom contributes both electrons (one pair). write the lewis structure for NH4. :resonance is the ability of a molecule to have more than 1 plausible lewis structure. (where 1 is not more correct than the other.) :write the lewis sturctures for NO3^1 :exceptions to the octet rule: : 1. mrs. croslin wants to give extra credit to me. : 2. she loves ap chem2 students : 3. incomplete octets: because there are only 3 valence electrons in the 4th bonding position, it hasno electrons to form a bond unless it forms a coordinate covalent bond. : 4. expanded octets: nonmetals in the 3rd period and beyond can use some of the d orbitals for their valence electrons. :write lewis structures for pcl3 pcl5 and sf6 : :SHAPES OF MOLECULES :molecules do not have a flat structure but a 3-d one based on the repulsion of electrons. :the vsepr( valence shell electron pair repulsion) model focuses on pairs of electrons in the valence shell. electron pairs repel one another, whether they are in chemical bonds or unshared pairs. electron pairs assume ashape around the atom to minimize repulsion. :the electro pair geometry refers to the shape of the pair of electrons (geometrical distribution of electrons). molecular geometry refers to the atomic shape. (where the atoms are actually located.) :possibilities of shapes: :based on central atoms :2 ep linear F2 180 :3 ep trigonal planar BF3 120 :4 ep tetrahedral CH4 109.5 :5 ep trigonal bipyramidal PCl5 90+120 :6 ep octahedral SCl6 90 : :because lone pairs spread out more than bonding pairs, molecular geometry can look different from the geometrical distribution of electrons. : :dipoles: :recall polar covalent molecules are molecules in which 1 atom has a greater pull for the electrons :the extent of the pull of the electron is given by a dipole. :pic: -+---> :dipoles can cancel each oher out and result in a nonpolar covalent molecule : :polarity depends on eelectronegativityand symmetry :CH4 - nonpolar :H2O - polar : :bond energy: :bond energy is defined, as the quantity of energy required to break 1 mole of covalent bonds in a gaseous species :bond energies can help predict if a reactin is endothermic or exothermic. : : REACTANTS PRODUCTS : weak bonds---->strong bonds exo : strong-------->weak bonds endo : :bondig theory: :we will be using the valence bond methond for understanding bonding at the electron level. it describes how covalent bonds form with the charge density of bonding electrons concentraated in the regin of orbital overlap. :let,s use carbon as an example for "hybridization" of orbitals for bonding. :according to this electron configuration, c can only form 2 bonds. in order for the carbon atom to form a filled octet, the electrons and orbitals must rearrange. the orbitals reach an intermidiate energy and hybridize to form 4 orbitals called the sp3 hybrid orbitals. :look at NH3, BF3, CH4 :we can have expanded octets that use the d orbitals. their hybridization is sp3d and sp3d2 : :multiple bonds and the VSPER theory :in order to form a multiple bond; unhybridized p orbitals must be used. the bond formed with the hybridized orbital is a bond, and the bond formed with the unhybridized p orbital is a bond. : :metallic bonding: :the challenge of a bonding theory for metals is to explain how so much bondin can occur for a few electrons. it should account for properties of metals (conduct electricity, luster, deformation, etc...) :one theory explains a solid metal of ions immersed into a "sea of electrons" electrons in a sea are free, not attached, and mobile. (can explain conductivity. if electrson are passed throgh one end from an outside source, they will leave the other end at the same rate). :free electrons are not limited in their ability to absorb photons of light so metals absorb visible light and are opaque. electrons at the surface can reradiate light that strikes the surface making it look lustrous. :deformatin occurs because the electronsare able to slide past one another. :