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121 Cards in this Set
- Front
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chemistry of carbon compounds
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organic chemistry
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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)
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-log (Ka)
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pKa
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[H3O+][A-]
---------------- [HA] |
Ka
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dissociates in water to give H3O+
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Arrhenius acid
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dissociates in water to give -OH
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Arrhenius base
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proton donor
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Bronsted-Lowry acid
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proton acceptor
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Bronsted-Lowry base
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electron-pair acceptor
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Lewis acid
electrophile |
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electron-pair donor
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Lewis base
nucleophile |
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acid that results from protonation of a base
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conjugate acid
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base that results from loss of a proton from an acid
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conjugate base
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bonding that occurs by the sharing of electrons in the region between to nuclei
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covalent bonding
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a covalent bond that involves the sharing of one pair of electrons
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single bond
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a covalent bond that involves the sharing of two pairs of electrons
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double bond
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a covalent bond that involves the sharing of three pairs of electrons
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triple bond
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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
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curved-arrow formalism
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orbitals with identical energies
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degenerate orbitals
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a measure of the polarity of a bond (or a molecule), proportional to the product of the charge separation times the bond length
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dipole moment
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the relative probability of finding an electron in a certain region of space
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electron density
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a measure of an element's ability to attract electrons. elements with higher electronegativities attract electrons more strongly
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electronegativity
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an electron-pair acceptor (Lewis acid)
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electrophile
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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.
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electrostatic potential map (EPM)
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the ratios of atoms in a compound
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empirical formula
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a method for keeping track of charges, showing what charge would be on an atom in a particular Lewis structure
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formal charges
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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
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Hund's rule
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bonding that occurs by the attraction of oppositely charged ions. usually results in the formation of a large, 3D crystal lattice
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ionic bonding
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atoms with the same number of protons but different numbers of neutrons; atoms of the same element but with different atomic masses
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isotopes
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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.
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Lewis structure
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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
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line-angle formula
skeletal structure stick figure |
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a pair of nonbonding electrons
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lone pair
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the number of atoms of each element in one molecule of a compound
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molecular formula
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simply gives the ratios of atoms of the different elements. does not give structural information
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empirical formula
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a region in an orbital with zero electron density
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node
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a flat (planar) region of space with zero electron density
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nodal plane
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valence electrons that are not used for bonding. A pair is often called a lone pair.
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nonbonding electrons
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an electron-pair donor (Lewis base)
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nucleophile
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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.
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octet rule
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an allowed energy state for an electron bound to a nucleus; the probability function that defines the distribution of electron density in space.
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orbital
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states that up to two electrons can occupy each orbital if their spins are paired.
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the Pauli exclusion principle
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the chemistry of carbon compounds.
the study of compounds derived from living organisms and their natural products. |
organic chemistry
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a measure of the acidity of a solution, defined as the negative logarithm of the H3O+ concentration.
-log(H3O+) |
pH
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a covalent bond in which electrons are shared unequally.
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polar covalent bone
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a bond with equal sharing of electrons
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nonpolar covalent bond
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a molecule or ion for which two or more valid Lewis structures can be drawn, differing only in the placement of the valence electrons.
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resonance hybrid
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Lewis structures in the resonance hybrids
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resonance forms
resonance structures |
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more important (lower energy) structures
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major contributors
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less important (higher energy) structures
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minor contributors
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when a charge is spread over two or more atoms by resonance
the molecule is said to be? |
delocalized
resonance stabilized |
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shows all the atoms and bonds in the molecule
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structural formulas
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shows each central atom along with the atoms bonded to it
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condensed structural formula
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assumes that there is a carbon atom wherever two lines meet or a line begins or ends
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line-angle formula
skeletal structure stick figure |
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the number of bonds an atom usually forms
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valence
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those electrons that are in the outermost shell
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valence electrons
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the belief that syntheses of organic compounds require the presence of a "vital force"
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vitalism
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provides the basis for life on Earth
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Carbon Compounds
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the study of gases, rock, and minerals, and the compounds that could be made from them
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inorganic chemistry
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distinctive feature of organic compounds is?
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they all contain one or more carbon atoms
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four examples of organic compounds in living organisms
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tobacco - nicotine
rose hips - vitamin C carmine opium poppies - morphine |
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the time it takes for half of the nuclei to decay
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half life
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states that we can never determine exactly where the electron is
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Heisenberg uncertainty
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as n (shells) increases, the shells are farther from the nucleus, ----, and can hold more ------.
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higher in energy
electrons |
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electron density is highest at ----- and falls off exponentially.
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nucleus
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first electron shell?
geometric shape? |
1s orbital
spherical |
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second electron shell(s)?
geometric shape? |
2s and 2p orbitals
2s spherical |
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smaller amount of electron density close to the nucleus and higher in energy, what electron shell?
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2s
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2p geometric shape
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consists of 2 lobes one on either side of nucleus
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What orbitals?
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2p orbitals
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can accommodate 2 e-
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first shell
one 1s orbital |
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can accommodate 8 e-
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second shell
one 2s orbital and three 2p orbitals |
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can accommodate 18 e-
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third shell
one 3s orbital, three 3p orbitals, and five 3d orbitals |
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1s1
1 ve- |
H
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1s2
2 ve- |
He
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"building up"
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Aufbau
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tells us how to build up the electronic configuration of an atom's ground (most stable) state
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aufbau principle
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1s2 2s1
1ve- |
Li
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1s2 2s2
2ve- |
Be
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1s2 2s2 2px1
3ve- |
B
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1s2 2s2 2px1 2py1
4ve- |
C
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1s2 2s2 2px1 2py1 2pz1
5ve- |
N
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1s2 2s2 2px2 2py1 2pz1
6ve- |
O
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1s2 2s2 2px2 2py2 2pz1
7ve- |
F
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S+
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a small amount of positive charge
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S-
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a small amount of negative charge
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useful for predicting the polarity of cavalent bonds
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Pauling electronegativity scale
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higher electronegativities generally have --- attraction for the bonding electrons
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more
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electronegativities increase from ---- to ---- across the periodic table
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left to right
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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 |
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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 |
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How to calculate formal charges
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count how many e- contribute to the charge of each atom
compare with the number of valence e- |
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the electrons that contribute to an atom's charge are:
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- all its unshared (nonbonding) e-
- half the (bonding) e- it shares with other atoms, or one e- of each bonding pair |
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formal charge calculation
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FC= [group #] - [nonbonding e-] - 1/2[shared e-]
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bonds between atoms with very large electronegativity differences (>2) are usually drawn as -----.
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ionic
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Calculation of empirical formula?
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1. # g / atomic weight of each element in molecule
2. divide each by the smallest you should get the ratios ie. CH2O |
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Calculation of molecular formula?
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known molecular weight divided by molecular weight.
ie 2x as much. 2(CH2O) = C2H4O2) |
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acids as substances that dissociate in water to give H3O+ ions
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Arrhenius theory
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acidity or basicity of an aqueous (water) solution is measured by the concentration of ----
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H3O+
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Kw
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= [H3O+][-OH]
= 1.00 * 10 (^-14) |
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[H3O+] > 10^-7
[-OH] < 10^-7 |
acidic
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[H3O+] < 10 ^ -7
[-OH] > 10 ^ -7 |
basic
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pH
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-log (H3O+)
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stronger the acid, the more it dissociates, giving a ---- of Ka
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larger value
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strong acids are almost completely ionized in water, and their dissociate constants are ----.
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greater than 1
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pKa
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= -log (Ka)
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strong acids generally have pKa value around --- .
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0, zero
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weak acids have pKa values that are greater than ---.
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4
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weaker acids have ---- pKa values
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larger
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in the reaction of an acid with a base, the equilibrium favors the --- acid and base
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weaker
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Ka value?
HCl |
1.6 * 10 ^2
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Ka value?
HF |
6.8 * 10 ^-4
strong acid, weak base |
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Ka value?
HCOOH |
1.7 * 10 ^-4
strong acid, weak base |
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Ka value?
CH3--COOH |
1.8 * 10 ^-5
strong acid, weak base |
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Ka value?
HCN: |
6.0 * 10 ^-10
weak acid |
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Ka value?
+NH4 |
5.8 * 10 ^-10
weak acid |
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Ka value?
CH3--OH |
3.2 * 10 ^-16
weak acid |
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Ka value?
H2O |
1.8 * 10 ^-16
weak acid, strong base |
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Ka value?
NH3 |
1.0 * 10 ^-33
very weak acid, very strong base |
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Ka Value?
CH4 |
<10 ^-40; pKa > 40
very strong base |
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Kb =
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[HA][-OH]
--------------- [A-] |
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pKb =
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= - log Kb
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pKa + pKb =
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= 14
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