CHEM 2211 1st Editiom Lecture 1Outline of Last Lecture I. Introduction to organic chemistry Outline of Current Lecture II. Rules and principles III. Bonding IV. Structures Current Lecture II. Rules and Principle Quantum Numbers Principle- n= row (electron shells) - l= electron orbital o n-1 o 0=So 1=Po 2=Do 3=F- Ml= magnetic quantum number o Range -1 to +1 - Ms= electron spin quantum number o + ½ or – ½ o Refers to up or down arrow in electron configurations Electron configuration 1s 2s 2p 3s 3p 4s 3d 4p Aafbau’s principle - Electrons go in the lowest energy orbital first Pauli principle These notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.- only two electrons in a pair Hunds rule - every orbital in a subshell must have one electron before doubling up III. Bonding Ionic bonding - always between a metal and a non-metal - transfer of electrons between metal to non-metal leaving a charge - Li. + :Cl: Li+ + Cl- covalent bonding - bond that involves the sharing of electrons between atoms- can be polar or non-polar non-polar- bond between 2 nonmetals - the electron or electrons they share are equally distributed o because they have the same electronegativity - Ex: H-H polar - bond between 2 non-metals - different electronegativities o results in uneven sharing of the electron or electrons - the separation of charges results in a bond dipole - Ex: H---Cl- S+ S-IV. StructureLewis dot Structures - Lines are drawn between atoms to show chemical bonds - Dots are put around the atoms to show lone pairs that are not bondedo H-:C=C:-HKekule structures - Same basic structure as lewis structure - Lone pairs are emitted though o H-C=C-H Condensed structure - Does not show connections - H-C-H instead CH2- Butane condensed o
View Full Document
Unlocking...