Chemistry 1027A/B Chapter Notes - Chapter 1-9: Methyl Acetate, Chloromethane, Acetamide

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Chemistry 2213: Textbook Notes
1.1
Ground-state electron configuration is the electron configuration of lowest energy
Rules of electron configuration:
oOrbitals in elements 1-18 fill in the order 1s,2s,2p,3p
oEach orbital contains max 2 electrons, however the 3 2p orbitals (x,y,z)
can be just grouped together and condensed to 2p6
oThe p orbitals are all equal in energy so we fill each orbital with an
electron before adding a second.
For elements in groups 4-7 and row 3, the valence electrons are in the third shell.
With 8 electrons, 3s and 3p are full but the 5 3d orbitals can accommodate an
additional 10 valence electrons. This is what makes significant differences in the
covalent bonding of second and third period elements like sulfur and oxygen.
1.2
According to the Pauling scale, fluorine is the most electronegative (4.0) and all
other values are assigned in relation to fluorine
Electronegativity increases from left to right in a period (increasing positive
charge on nucleus) and from bottom up (decreasing distance of valence electrons
from the nucleus) (from one side to fluorine)
Electronegativity also varies based on the oxidation state
1.9 or < is ionic, 1.9> is covalent
Separation of charge produces a dipole (two poles). The presence of a dipole is
shown by an arrow with the head at the negative and the tail at the positive.
Formal charge = number of valence electrons in neutral unbounded atom – (all
unshared electrons + one-half of all shared electrons)
1.3
VSEPR – Valence Shell
Electron Pair Repulsion
1.4
A molecule is polar if:
oIt has polar bonds
oAND the vector
sum of its bond
dipoles is NOT
zero
If the vector sum of a
molecule’s two bond
dipoles is zero, the
molecule is nonpolar.
1.5
To describe why no
molecules/ions Lewis structures are adequate, we turn to the theory of resonance:
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oDeveloped by Linus Pauling in 1930s
oMany molecules/ions are best described by writing two or more Lewis
structures and considering the real molecule/ion to be a composite of these
structures.
oResonance contributing structures are individual Lewis structures that
differ only in the distribution of valence electrons; all of equal energy.
oWe show that the real molecule/ion is a resonance hybrid of all the various
contributing structures by interconnecting them with double-headed
arrows.
oMolecules/ions, however, do not constantly change their electron pairs and
bonds back and forth, they have only one real structure that is best
described as a hybrid of its various contributing structures.
A curved arrow shows the redistribution of valence electrons from the electron
pair’s origin to its destination.
Rules from writing resonance contributing structures:
oMust have same number of valence electrons
oMust obey the rules of covalent bonding (3rd period elements can have up
to 12 electrons in their valence shells)
oPositions of nuclei must be the same
oSame total number of paired and unpaired electrons
1.6
According to the orbital overlap model, a covalent bond is formed when a portion
of an atomic orbital of one atom overlaps a portion of an atomic orbital of another
atom
A sigma bond ( ) is a covalent bond where the overlap of atomic orbitals is σ
concentrated along the bond axis
Bond angles in molecules:
oSingle bonds: 109.5°
oDouble bonds: 120°
oTriple bonds: 180°
Hybrid orbitals are orbitals produced from a combination of two or more atomic
orbitals
The superscript on hybrid orbitals tells you how many atomic orbitals have been
combined to form the hybrid orbitals.
sp3 hybrid orbitals: 109.5°
ohave one 2s and three 2p orbitals; occurs in sets of four (tetrahedral)
oeach orbital either forms a sigma bond with hydrogen or holds an unshared
pair of electrons
sp2 hybrid orbitals: 120°
ohave one 2s and two 2p orbitals; occurs in sets of three
othe third 2p orbital not involved in hybridization is two lobes lying
perpendicular to the plane of the hybrid orbitals.
oForm double bonds
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oSigma bond between the two central carbons and between carbon and
hydrogens are formed
oA pi ( )π bond is a covalent bond formed by the overlap of parallel p
orbitals; weaker than sigma bonds
sp hybrid orbitals: 180°
ohave one 2s and one 2p orbital; occurs in sets of two
oa carbon-carbon triple bond has 1 sigma bond and 2 pi bonds
1.7
Functional groups:
oSites of chemical reactions
oDetermine the physical properties
oHow we divide organic
compounds into
families
oBasis for naming
organic compounds
Alcohols
osp3 hybridized carbon
oprimary (1°), secondary
(2°), or tertiary (3°)
depending on # of
carbon atoms bonded to
carbon bearing the
hydroxyl
Amine
osp3 hybridized nitrogen
oprimary (1°), secondary
(2°), or tertiary (3°)
depending on # of
carbon atoms bonded to
nitrogen
Aldehydes and ketones
oCarbonyl group
oAldehyde is at the end,
ketone is in middle
2.1
Arrhenius acid: dissolves in water to produce H+ ions
Arrhenius base: dissolves in water to produce OH- ions
Problem: H+ ions don’t exist in water because they react with water to form
hydronium
2.2
Bronsted-Lowry
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