 chemical bond - force that holds groups of atoms together  group function as a unit  bond NRG – NRG required to break bond  indicates strength of bond  bond length – distance between two atoms bonded together  indicates most stable(least amount of NRG) state between two atoms

 ionic bond – electromagnetic force that holds two oppositely charged ions together  formed between cations and anions  cations formed when metals lose e - ‘s ▪Na  Na + + e - (oxidation)  anions formed when nonmetals gain e - ‘s ▪Cl + e -  Cl - (reduction)  opposite charges attract(electromagnetic force) ▪Na + + Cl -  Na + Cl -

 why do substances form ionic bonds???  lowest possible NRG for the system(substances involved with bonding)  see figure 8.9 on page change Li(s) to Li(g)  endothermic Li(s) kJ/mol  Li(g) 2. Li(g) oxidizes  endothermic Li(g) kJ/mol  Li + + e - 3. fluorine molecules separate and form fluorine atoms  endothermic ½ F 2 (g) + 77 kJ/mol  F(g)

4. fluorine reduces  exothermic F(g) + e -  F - (g) kJ/mol 5. ionic bond formed  extremely exothermic Li + (g) + F - (g)  Li + F kJ/mol total endothermic processes = 758 kJ/mol total exothermic processes = 1375 kJ/mol NET NRG = 617 kJ/mol  overall process is exothermic so an ionic bond forms

 metals lose e - (oxidation)  nonmetals gain e - (reduction)  ionic compound formula  total oxidation = total reduction Mg 2+ F -  MgF 2  utilize criss-cross method Mg 2+ F - MgF 2

 dissolving ionic crystals

 covalent bond – force of attraction between 2 atoms when e - ‘s are shared  each nucleus attracts the other atoms e - ‘s  balance between attraction and repulsion

 bonds result from system trying to attain lowest possible NRG state  two driving forces in nature 1)lowest NRG 2)highest entropy(disorder/chaos)  ionic and covalent bonds generally form to attain the lowest NRG state for atoms involved  single covalent bond – 1 pair e - shared  H-H, F-F, Cl-Cl, Br-Br, I-I  double covalent bond – 2 pair e - shared  O=O  triple covalent bond – 3 pair e - shared  N≡N

 molecular orbital  most probable location of e - ‘s when covalently bonded

 sigma (  ) bond centers along the internuclear axis.  single covalent bond  pi (  ) bond occupies space above and below internuclear axis.  2 nd or 3 covalent bond

 electronegativity  attraction an atom has for another atom’s e - ‘s  arbitrary scale – based on F  varies periodically (click here) (click here) ▪generally increases across and decreases down  electronegativity difference can generally predict type of bond  if diff. > 1.7 = ionic bond  if 1.7 > diff. > 0.3 = polar covalent bondpolar covalent bond  if diff. < 0.3 = nonpolar covalent bondnonpolar covalent bond

 nonpolar covalent bond  pure covalent bond  sharing of e - ‘s is equal  no poles/charges created

Lewis structures  localized e - model for diagramming bonds and molecular shapes 1) total valence e - of all atoms in molecule HCl = 8 valence e - 2) write symbols for each element least number of = interior atom hydrogen is always exterior atom H Cl

3) add a pair of e - between atoms bonding together H : Clvalence e - remaining = 6 4) add remaining e - in pairs to exterior atoms to form octets if e - remain add them to interior atoms to form octets not H, H forms duets if necessary, move e - pairs to create octets H : Cl 5) all shared pair of e - become dashes H - Cl : : : : :

6) follow VSEPR for shape VSEPR – valence shell electron pair repulsion theory model used to predict geometry/shape of a molecule based on the repulsion of e - pairs e - pairs repel each other to maximum distance in 2-dimension = 90 o in 3-dimension = o

Molecular Geometry 3-dimensional shapes determine the physical and chemical properties of molecules example – sucrose- its molecular shape fits the nerve receptors of the tongue for sweetness sugar substitutes(Splenda, Nutrasweet, …) have similar shapes as sucrose 1) linear all atoms lie in a straight line HCl, CO 2

2) bent atoms not in straight line e - pairs point to 4 corners of tetrahedron

3) trigonal pyramidal 3 atoms bonded to central atom and a pair of nonbonding e -

4) trigonal planar 3 groupings of e - around central atom all atoms lie in same plane

5) tetrahedral 4 groupings of e - around central atom

 resonance  ability to draw more than one acceptable shape for a molecule  originally believed molecule resonated between different shapes ▪benzene, nitrate ion  actual structure is an average of all resonant images

 exceptions to octet rule  C, N, O, F always follow octet rule  Be and B often have less than 8  3 rd period and heavier usually follow ▪ some may exceed by putting e - in unoccupied d- orbitals ▪when writing Lewis structures follow octet rule ▪if e - remain add them to elements with d orbitals

polar molecules molecule with oppositely, partially charged atoms on opposite sides aka – dipoles, dipole moments

molecule that has an asymmetrical distribution of charge partial charges not evenly distributed around central atom polar molecules must have: polar bonds(partial charges) unevenly distributed partial charges

HCl H = 2.1, Cl = 3.0 (electroneg diff = 0.9) = bond is polar H = δ+, Cl = δ- linear shape polar molecule N 2 N = 3.0 (electroneg diff = 0) = bond is nonpolar nonpolar molecule

H 2 O H = 2.1, O = 3.5 (electroneg diff = 1.4) = bonds are polar H = δ+, O = δ- bent shape polar molecule NH 3 N = 3.0, H = 2.1 (electroneg diff = 0.9) = bonds are polar H = δ+, N = δ- trigonal pyramidal shape polar molecule

CCl 4 C = 2.5, Cl = 3.0 (electroneg diff = 0.5) = bonds are polar C = δ+, Cl = δ- tetrahedral shape nonpolar molecule

 hybridization  formation of hybrid orbitals from atomic orbitals of similar NRG  sp 3 hybridization

 sp 2 hybridization

 sp hybridization

 naming binary molecules 1) determine # of 1 st element ▪use prefix if more than one ▪1=mono-2=di- ▪3=tri-4=tetra- ▪5=penta-6=hexa- ▪7=hepta-8=octa- ▪9=nona-10=deca- 2) name element 3) determine # of 2 nd element ▪use prefix except if bonded to H 4) use root of element name 5) end with -ide

 polar covalent bond  bond in which e - ‘s are shared unequally ▪electronegativity of one atom is higher than other