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121 Cards in this Set

  • Front
  • Back
chemistry of carbon compounds
organic chemistry
the equilibrium constant for the reaction of the acid with water to generate H3O+
HA + H2O --> H3O+ + A-
[H3O+][A-]
----------------
[HA]
acid-dissociation constant (Ka)
-log (Ka)
pKa
[H3O+][A-]
----------------
[HA]
Ka
dissociates in water to give H3O+
Arrhenius acid
dissociates in water to give -OH
Arrhenius base
proton donor
Bronsted-Lowry acid
proton acceptor
Bronsted-Lowry base
electron-pair acceptor
Lewis acid
electrophile
electron-pair donor
Lewis base
nucleophile
acid that results from protonation of a base
conjugate acid
base that results from loss of a proton from an acid
conjugate base
bonding that occurs by the sharing of electrons in the region between to nuclei
covalent bonding
a covalent bond that involves the sharing of one pair of electrons
single bond
a covalent bond that involves the sharing of two pairs of electrons
double bond
a covalent bond that involves the sharing of three pairs of electrons
triple bond
method of drawing curved arrows to keep track of electron movement from nucleophile to electrophile (or within a molecule) during the course of a reaction
curved-arrow formalism
orbitals with identical energies
degenerate orbitals
a measure of the polarity of a bond (or a molecule), proportional to the product of the charge separation times the bond length
dipole moment
the relative probability of finding an electron in a certain region of space
electron density
a measure of an element's ability to attract electrons. elements with higher electronegativities attract electrons more strongly
electronegativity
an electron-pair acceptor (Lewis acid)
electrophile
a computer-calculated molecular representation that uses colors to show the charge distribution in a molecule. Uses red to show electron-rich regions (most negative) and blue to shoe electron-poor regions (most positive). The intermediate colors orange, yellow, and green.
electrostatic potential map (EPM)
the ratios of atoms in a compound
empirical formula
a method for keeping track of charges, showing what charge would be on an atom in a particular Lewis structure
formal charges
when there are two or more unfilled orbitals of the same energy (degenerate orbitals), the lowest-energy configuration places the electrons in different orbitals (with parallel spins) rather than paired in the same orbital
Hund's rule
bonding that occurs by the attraction of oppositely charged ions. usually results in the formation of a large, 3D crystal lattice
ionic bonding
atoms with the same number of protons but different numbers of neutrons; atoms of the same element but with different atomic masses
isotopes
a structural formula that shows all valence electrons, with the bonds symbolized by dashes (--) or by pairs of dots, and nonbonding electrons symbolized by dots.
Lewis structure
a shorthand structural formula with bonds represented by lines. Carbon atoms are implied wherever two lines meet or a line begins or bends. Atoms other that C and H are drawn in, but hydrogen atoms are not shown unless they are on an atom that is drawn. Each carbon atom is assumed to have enough hydrogens to give it four bonds
line-angle formula
skeletal structure
stick figure
a pair of nonbonding electrons
lone pair
the number of atoms of each element in one molecule of a compound
molecular formula
simply gives the ratios of atoms of the different elements. does not give structural information
empirical formula
a region in an orbital with zero electron density
node
a flat (planar) region of space with zero electron density
nodal plane
valence electrons that are not used for bonding. A pair is often called a lone pair.
nonbonding electrons
an electron-pair donor (Lewis base)
nucleophile
atoms generally form bonding arrangements that give them filled shells of electrons (noble gas configurations). for the second row elements, this configuration has eight valence electrons.
octet rule
an allowed energy state for an electron bound to a nucleus; the probability function that defines the distribution of electron density in space.
orbital
states that up to two electrons can occupy each orbital if their spins are paired.
the Pauli exclusion principle
the chemistry of carbon compounds.
the study of compounds derived from living organisms and their natural products.
organic chemistry
a measure of the acidity of a solution, defined as the negative logarithm of the H3O+ concentration.
-log(H3O+)
pH
a covalent bond in which electrons are shared unequally.
polar covalent bone
a bond with equal sharing of electrons
nonpolar covalent bond
a molecule or ion for which two or more valid Lewis structures can be drawn, differing only in the placement of the valence electrons.
resonance hybrid
Lewis structures in the resonance hybrids
resonance forms
resonance structures
more important (lower energy) structures
major contributors
less important (higher energy) structures
minor contributors
when a charge is spread over two or more atoms by resonance

the molecule is said to be?
delocalized
resonance stabilized
shows all the atoms and bonds in the molecule
structural formulas
shows each central atom along with the atoms bonded to it
condensed structural formula
assumes that there is a carbon atom wherever two lines meet or a line begins or ends
line-angle formula
skeletal structure
stick figure
the number of bonds an atom usually forms
valence
those electrons that are in the outermost shell
valence electrons
the belief that syntheses of organic compounds require the presence of a "vital force"
vitalism
provides the basis for life on Earth
Carbon Compounds
the study of gases, rock, and minerals, and the compounds that could be made from them
inorganic chemistry
distinctive feature of organic compounds is?
they all contain one or more carbon atoms
four examples of organic compounds in living organisms
tobacco - nicotine
rose hips - vitamin C
carmine
opium poppies - morphine
the time it takes for half of the nuclei to decay
half life
states that we can never determine exactly where the electron is
Heisenberg uncertainty
as n (shells) increases, the shells are farther from the nucleus, ----, and can hold more ------.
higher in energy
electrons
electron density is highest at ----- and falls off exponentially.
nucleus
first electron shell?
geometric shape?
1s orbital
spherical
second electron shell(s)?
geometric shape?
2s and 2p orbitals
2s spherical
smaller amount of electron density close to the nucleus and higher in energy, what electron shell?
2s
2p geometric shape
consists of 2 lobes one on either side of nucleus
What orbitals?
2p orbitals
can accommodate 2 e-
first shell
one 1s orbital
can accommodate 8 e-
second shell
one 2s orbital and three 2p orbitals
can accommodate 18 e-
third shell
one 3s orbital, three 3p orbitals, and five 3d orbitals
1s1
1 ve-
H
1s2
2 ve-
He
"building up"
Aufbau
tells us how to build up the electronic configuration of an atom's ground (most stable) state
aufbau principle
1s2 2s1
1ve-
Li
1s2 2s2
2ve-
Be
1s2 2s2 2px1
3ve-
B
1s2 2s2 2px1 2py1
4ve-
C
1s2 2s2 2px1 2py1 2pz1
5ve-
N
1s2 2s2 2px2 2py1 2pz1
6ve-
O
1s2 2s2 2px2 2py2 2pz1
7ve-
F
S+
a small amount of positive charge
S-
a small amount of negative charge
useful for predicting the polarity of cavalent bonds
Pauling electronegativity scale
higher electronegativities generally have --- attraction for the bonding electrons
more
electronegativities increase from ---- to ---- across the periodic table
left to right
Electronegativity of:
H
Li
Be
B
C
N
O
F
H = 2.2
Li = 1.0
Be = 1.6
B = 1.8
C = 2.5
N = 3.0
O = 3.4
F = 4.0
Electronegativity of:
Na
Mg
Al
Si
P
S
Cl
K
Br
I
Na = 0.9
Mg = 1.3
Al = 1.6
Si = 1.9
P = 2.2
S = 2.6
Cl = 3.2
K = 0.8
Br = 3.0
I = 2.7
How to calculate formal charges
count how many e- contribute to the charge of each atom
compare with the number of valence e-
the electrons that contribute to an atom's charge are:
- all its unshared (nonbonding) e-
- half the (bonding) e- it shares with other atoms, or one e- of each bonding pair
formal charge calculation
FC= [group #] - [nonbonding e-] - 1/2[shared e-]
bonds between atoms with very large electronegativity differences (>2) are usually drawn as -----.
ionic
Calculation of empirical formula?
1. # g / atomic weight of each element in molecule
2. divide each by the smallest
you should get the ratios
ie. CH2O
Calculation of molecular formula?
known molecular weight divided by molecular weight.
ie 2x as much. 2(CH2O) = C2H4O2)
acids as substances that dissociate in water to give H3O+ ions
Arrhenius theory
acidity or basicity of an aqueous (water) solution is measured by the concentration of ----
H3O+
Kw
= [H3O+][-OH]
= 1.00 * 10 (^-14)
[H3O+] > 10^-7
[-OH] < 10^-7
acidic
[H3O+] < 10 ^ -7
[-OH] > 10 ^ -7
basic
pH
-log (H3O+)
stronger the acid, the more it dissociates, giving a ---- of Ka
larger value
strong acids are almost completely ionized in water, and their dissociate constants are ----.
greater than 1
pKa
= -log (Ka)
strong acids generally have pKa value around --- .
0, zero
weak acids have pKa values that are greater than ---.
4
weaker acids have ---- pKa values
larger
in the reaction of an acid with a base, the equilibrium favors the --- acid and base
weaker
Ka value?
HCl
1.6 * 10 ^2
Ka value?
HF
6.8 * 10 ^-4
strong acid, weak base
Ka value?
HCOOH
1.7 * 10 ^-4
strong acid, weak base
Ka value?
CH3--COOH
1.8 * 10 ^-5
strong acid, weak base
Ka value?
HCN:
6.0 * 10 ^-10
weak acid
Ka value?
+NH4
5.8 * 10 ^-10
weak acid
Ka value?
CH3--OH
3.2 * 10 ^-16
weak acid
Ka value?
H2O
1.8 * 10 ^-16
weak acid, strong base
Ka value?
NH3
1.0 * 10 ^-33
very weak acid, very strong base
Ka Value?
CH4
<10 ^-40; pKa > 40
very strong base
Kb =
[HA][-OH]
---------------
[A-]
pKb =
= - log Kb
pKa + pKb =
= 14