Atomic Bonds and Lewis Structures
Atomic Bonds and Lewis Structures CHEM - 10060 - 001
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This 4 page Class Notes was uploaded by Nick Manning on Sunday April 17, 2016. The Class Notes belongs to CHEM - 10060 - 001 at Kent State University taught by TBA in Fall 2015. Since its upload, it has received 10 views. For similar materials see GENERAL CHEMISTRY I in Chemistry at Kent State University.
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Date Created: 04/17/16
Week of 4/11 – 4/16 Bond type is influenced by metallic character of the things bonding Metal + metal = metallic bond. Metal + nonmetal = ionic bond Nonmetal + nonmetal = covalent bond COVALENT BONDS -Happen bw elements of similar metallic character - tend to be stronger than ionic bonds -involve e- sharing -Involve orbital overlap - bonds usually don’t break in phase changes of small molecules or in solution; whereas in an ionic bond they always break in phase changes, (s) (l), and in solution if they are soluble - both nuclei of bonded atoms attract shared e- - not a static bond, it’s more of a dynamic pulsing bond bw the two atoms There are specific partners in covalent bonds and not specific in ionic bonds Covalent bond length- length bw two bonded atoms’ nuclei Vanderwaals distance- always larger than bond length, distance bw the nuclei of the two unbonded atoms nearest to each other # of covalent bonds an atom forms depends on its need for valence e- For example, Hydrogen (H) needs one, so it makes one bond H --- H H follows the duet rule, meaning it needs two valence electrons and a single bond pair Period 2 elements follow the octet rule (aiming for full 2s2p shells) Period 3 elements = octet rule & often expanded octet rule BOND ORDER Bond order is just a number for each set of different bonds that represents the amount of bonds made. A single bond has a bond order of 1, double bonds have a bond order of 2, and triple bonds have a bond order of 3. BOND LENGTH Bond Length increases as the size of bonded atoms increases, so it is important to know your trends for sizes now. BOND ENERGY Bond Energy is the energy required to break the covalent bond by overcoming the attraction between the nuclei and the shared e-. A stronger bond results in a higher bond energy. Bond energy increases as bond order increases &bond length decreases. POLAR & NON POLAR COVALENT BONDS Nonpolar covalent bonds- have only one pair of hared e-, e- are shared equally, e- density is the highest bw nuclei, even distribution of e- Polar Covalent Bonds- e- aren’t shared equally, uneven distribution of e- ELECTRONEGATIVITY -the ability to attract shared e- pairs, must be measured in a covalent bond - different from e- Affinity; e- affinity is the ability to completely gain an e-; measured in gas phase. - Fluorine is the most electronegative element, everything is based off of this rating. - increases up a group and increases left to right down a period -⧍EN is the rating system to determine if a bond is polar, nonpolar, or ionic. - if ⧍ EN is above .4 then it is polar, above 1.7 is ionic EX) C has an EN of 2.5, N has an EN of 3.0; find the ⧍EN and state what type of bond this is. 3.0-2.5 = .5 making this a polar covalent bond. METALLIC BONDING -uses e- pooling and forms an e- sea where covalent just shares one pair of e- e- sea: the outer orbitals overlap and the valence e- are pooled together - Delocalized e- sea provides strength w/o rigidity - Valence e- are free to move independently, allowing for great heat and electricity conduction - Metals are strong, but can be shaped & worked w/o breaking LEWIS STRUCTURES - Used for covalent species - Some prefer the octet to make them neutral O H H That is an example of an H2O bond using Lewis Structures. Formal charge must also be shown with these structures - Compares the valence e- of unbonded atoms to “e- owned” by bonded atoms - Always 0 when bonded in preferred pattern e- owned = all the non-bonding e- (dots; the lone pairs and single e-) - Also the lines = ½ the bonding e- Formal charge is calculated by: val e- - e- owned A formal charge of 0 mean the atom is making its preferred bonding. RESONANCE Electrons in a covalent bond are not just localized between two atoms, they are delocalized between all the atoms in a bond. This can be shown in a model called the Resonance Hybrid model. It is a sort of weighted average between the two contributing atoms. There are certain rules for using the resonance model; resonance structures: have the same skeleton, the same total # of e-, the same net charge, and the Lewis structures must be valid.
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