Chem Final

Front 1

pKb

Back 1

pKb=pKw-pKa
~ 14-pKa

Front 2

pKa

Back 2

pKa= -logKa

weak acid pKa: -2-12
strong acid pKa: -2<pKa

Front 3

Kb Value

Back 3

high: product favored and strong base
low:weak base and reactant favore rxn

Front 4

Ka value

Back 4

high:strong acids have Ka>1 and product favored
low:weak acids have Ka<1 and are reactant favored

Front 5

how carboxylic acids react as acids?

Back 5

only the hydrogen atom bond to the oxygen atom of the carboxylic acid group is sufficiently positive to be donated as an H+ ion in (aq)sol'n. The 2 highly electronegative oxygen atoms of the carboxylic acid group pull e- density away from the H atom. The C-H bond makes those H atoms non-acidic.

Front 6

Buffer/How to prepare

Back 6

->a chemical system that resists change in pH
->must contain a weak acid that can react with added base and a weak base that can react with added acid. Also, they can't react with each other.[conj acid/base pair]

Front 7

pOH

Back 7

=-log[OH-]

Front 8

pH

Back 8

=-log[H3O+]

Front 9

Henderson-Hasselbach Equation

Back 9

pH=pKa+log([conj base]/[conj acid])
pH=pKa+log([A-]/[HA])

Front 10

Water autoionization in Acid/Base Chemistry

Back 10

when water molecules react to produce H3O+ and OH-. One water molecule serves and an H+ acceptor (base) while the other is an H+ donor (acid).

Front 11

Bronsted-Lowry Definition of Acids/bases

Back 11

Acids: hydrogen ion donors
Bases: hydrogen ion acceptors

Front 12

Lewis Acid/Base Definition

Back 12

Acid: substance that can accept a pair of e-'s to form a new bond
Base:Substance that can donate a pair of e-'s to form a new bond

Front 13

How amines react as bases

Back 13

Drugs, high-molecular weight amines that are not soluble in water,blood plasma. By rxn with HCl, aminos are conveted to soluble hydrochloride salts. Resulting salts have the general formula BH+Cl- (B=basic amine)

Front 14

How salts and their ions influence the pH of (aq) sol'ns

Back 14

*salt of st base/st acid=neutral
*salt of st base/wk acid=basic
*salt of wk base/stacid=acidic

Front 15

Strength of acid from molecular shape

Back 15

H-A bond must be broken in order for the acid to transfer H+ to water. Will only occur if H-A bond is polar. H-C bond is basically nonacidic b/c of little polarity. H-Br is much more polar, thus, HBr is a strong acid. (Ka=10^8)

Front 16

Anode

Back 16

the electrode of an electochemical cell at which oxidation occurs

Front 17

Cathode

Back 17

the electrode of an electrochemical cell at which reduction occurs

Front 18

Electrochemical Cell

Back 18

A combination of anode, cathode, and other materials arranged so that a product-favored redox rxn can cause a current to flow or an electric current can cause a reactant-favored redox rxn to occur

Front 19

Nernst Equation

Back 19

Relates the potential of an e-cell to the concentrations of the chemical species involved in the redox rxns occuring in the cell.
Ecell=Eocell-(RT/nF)lnQ

Front 20

salt bridge

Back 20

a sol'n of salt arranged so that the bulk of that sol'n cannot flow into the cell sol'ns, but the ions can pass freely.

Front 21

How a battery works:

Back 21

*redox rxn that occurs must favor the formation of products
*must be an external circuit through which e-'s flow
*must be a salt bridge, porous barrier,ect. that allows ions in the salt bridge to flow into the electrode compartments to offset charge build-up.

Front 22

oxidizing agent

Back 22

the species causing the electron loss (is reduced)

Front 23

reducing agent

Back 23

species causing the the electron gain (is oxidized)

Front 24

OIL RIG

Back 24

Oxidation Is Loss; Reduction Is Gain

Front 25

relationship b/w Eo and Gibbs free energy

Back 25

ΔG=-nFEorxn=-RTlnK
Eo>0:product favored ΔG<0

Front 26

Indicator/How to pick one

Back 26

Use indicator to monitor pH changes and find the equivalence point(s)

Front 27

Arrhenius def of Acid/Bases

Back 27

Acid: increases the [H+] when dissolved in water
Base: increases the [OH-] when dissolved in water

Front 28

Oxidation States

Back 28

*The oxidation number of an atom of a free element equals 0.
*The oxidation number of a monatomic ion equals the charge on the ion.
*The oxidation number of fluorine in all compounds equals -1.
*The oxidation number of hydrogen in combination with nonmetals equals +1.
The oxidation number of hydrogen in combination with metals equals -1.
*The oxidation number of group 1 metals in compounds equals +1.
*The oxidation number of group 2 metals in compounds equals +2.
*The algebraic sum of the oxidation numbers of all the atoms in a chemical formula equals the net charge on the species.
*In ionic compounds, the oxidation numbers in the cation and the anion are independent and can be assigned separately.
*The oxidation number of oxygen in compounds equals -2.
*In binary combinations of nonmetals, the element closer to fluorine on the periodic table is given a negative oxidation number, equal to the charge on its common monatomic ion.

Front 29

Redox rxns w/changes in oxidation states

Back 29

??

Front 30

1/2 Equivilance point

Back 30

[HA]=[A-]
pH=pKa

Front 31

Relationship B/W Energy, Wavelength and Frequency

Back 31

E=h*frequency
h=plancks constant(6.626x10^-34 J*s)

Front 32

4 Quantum Numbers

Back 32

n=orbital size and energy
l= orbital shape
l=0:s orbital
l=1:p orbital
l=2:d orbital
ml:spatical orientation
-l<ml<l
ms: spin; either -1/2 or 1/2

Front 33

s,p,d orbital shapes

Back 33

s: round
p: node at the middle, two loops at either side
d:like 2 p orbitals at 90deg to each other.

Front 34

Configurations of Cations and Anions

Back 34

cation<neutral<anion

Front 35

Ionization Energy

Back 35

minimum energy required to remove one electron from each atom in a mole of atoms in the gaseous state. The first ionization energy is the energy required to remove one, the nth ionization energy is the energy required to remove the atom's nth electron, not including the n-1 electrons before it. Trend-wise, the ionization potentials tend to increase while one progresses across a period because the greater number of protons (higher nuclear charge) attract the orbiting electrons more strongly, thereby increasing the energy required to remove one of the electrons.

Front 36

Electron Affinity

Back 36

energy required to detach an electron from a singly-charged anion. More commonly, the electron affinity measures the energy released when an electron is added to a stable atom, thereby creating an anion.

Front 37

Periodic trend of Atomic radius

Back 37

r increases downward
r decreases left to right
-> due to effective nuclear charge

Front 38

periodic trend of ion radius

Back 38

cations get smaller than the original atom and anions get larger than the original atom.

Front 39

periodic trend of electronegatvity

Back 39

moves horizontally across a period in the periodic table, the electronegativity increases due to the stronger attraction that the atoms obtain as the nuclear charge increases. Moving down a group, the electronegativity decreases due to the larger distance between the nucleus and the valence electron shell, thereby decreasing the attraction, making the atom have less of an attraction for electrons or protons.

Front 40

periodic trend of ionization energy

Back 40

There will be an increase of ionization energy from left to right of a given period and a decrease from top to bottom. As a rule, it requires far less energy to remove an outer-shell electron than an inner-shell electron. As a result the ionization energies for a given element will increase steadily within a given shell, and when starting on the next shell down will show a drastic jump in ionization energy. Simply put, the lower the principal quantum number, the higher the ionization energy for the electrons within that shell.

Front 41

covalent bond

Back 41

bonds formed by sharing e-'s to attain NOble gas configurations

Front 42

ionic bond

Back 42

bond formed by the attraction between two oppositely charged ions.The metal donates one or more electrons, forming a positively charged ion or cation with a stable electron configuration. These electrons then enter the non metal, causing it to form an anion.

Front 43

polar covalent bond

Back 43

dipole-dipole intermolecular forces between the slightly positively-charged end of one molecule to the negative end of another or the same molecule. Molecular polarity is dependent on the difference in electronegativity between atoms in a compound and the asymmetry of the compound's structure.

Front 44

unit cell

Back 44

smallest part of a lattice that when repeated along the directions specified by its edges reproduces the crystal.
ex: sc, bcc,fcc,hcp,ccp

Front 45

formal charge

Back 45

FC = number of valence electrons of the atom in isolation - number of lone pair electrons on this atom in the molecule - half the total number of electrons participating in covalent bonds with this atom in the molecule.

Front 46

ionic solids

Back 46

chemical compound in which ions are held together in a lattice structure by ionic bonds. High melting and boiling points, brittle.

Front 47

molecular substances

Back 47

covalent bonded molecules; weaker forces b/w molecules; stronger forces inside molecules;soft; low melting and boiling pts;poor conductors

Front 48

metals

Back 48

closely packed metal atoms; metal ions in sea of e-'s; good conductors

Front 49

network substances

Back 49

covalently bonded arrays of atoms; graphite, diamond, quartz

Front 50

hydrocarbon

Back 50

has only C and H atoms

Front 51

sigma bond

Back 51

strongest type of covalent chemical bond. symmetrical with respect to rotation about the bond axis.They are obtained by an overlap of atomic orbitals. According to the sigma bond rule, the number of sigma bonds in a molecule is equivalent to the number of atoms plus the number of rings minus one.

Front 52

pi bond

Back 52

covalent chemical bonds where two lobes of one involved electron orbital overlap two lobes of the other involved electron orbital.this bond's weakness is explained by significantly less overlap between the component p-orbitals due to their parallel orientation.

Front 53

alcohols

Back 53

any organic compound in which a hydroxyl group (-OH) is bound to a carbon atom of an alkyl or substituted alkyl group.

Front 54

aldehydes

Back 54

organic compound containing a terminal carbonyl group. This functional group, which consists of a carbon atom bonded to a hydrogen atom and double-bonded to an oxygen atom (chemical formula O=CH-), is called the aldehyde group.

Front 55

ketones

Back 55

chemical compound that contains a carbonyl group. A ketone can be generally represented by the chemical formula:

R1(CO)R2.

Front 56

carboxylic acids

Back 56

characterized by the presence of a carboxyl group, which has the formula -C(=O)OH, usually written -COOH or -CO2H. [1] Carboxylic acids are Brønsted-Lowry acids — they are proton donors. Salts and anions of carboxylic acids are called carboxylates.

Front 57

esters

Back 57

consist of an inorganic or organic acid in which at least one -OH (hydroxyl) group is replaced by an -O-alkyl (alkoxy) group

Front 58

amines

Back 58

contain a basic nitrogen atom with a lone pair. Amines are derivatives of ammonia, wherein one or more hydrogen atoms are replaced by organic substituents such as alkyl and aryl groups.

Front 59

amides

Back 59

an amine where one of the nitrogen substituents is an acyl group; it is generally represented by the formula: R1(CO)NR2R3 , where either or both R2 and R3 may be hydrogen.

Front 60

primary alcohol

Back 60

an alcohol which has the hydroxyl radical connected to a primary carbon. It can also be defined as a molecule containing a “–CH2OH” group

Front 61

esterification rxn

Back 61

chemical reaction in which two reactants (typically an alcohol and an acid) form an ester as the reaction product. used in the fragrance and flavor industries.

Front 62

condensation rxn

Back 62

reaction in which two molecules or (functional groups) combine to form one single molecule, together with the loss of a small molecule.

Front 63

hydrolysis rxn

Back 63

reaction during which one or more water molecules are split into hydrogen and hydroxide ions which may go on to participate in further reactions.It is the type of reaction that is used to break down certain polymers.

Front 64

primary structure of proteins

Back 64

the amino acid sequence

Front 65

secondary structure of proteins

Back 65

regularly repeating local structures stabilized by hydrogen bonds. The most common examples are the alpha helix and beta sheet.Because secondary structures are local, many regions of different secondary structure can be present in the same protein molecule.

Front 66

tertiary structure of proteins

Back 66

the overall shape of a single protein molecule; the spatial relationship of the secondary structures to one another.

Front 67

quaternary structure of proteins

Back 67

the shape or structure that results from the interaction of more than one protein molecule.

Front 68

alpha helix

Back 68

a right/left-handed coiled conformation, resembling a spring, in which every backbone N-H group donates a hydrogen bond to the backbone C=O group of the aminoacid 4 residues earlier.

Front 69

beta sheet

Back 69

consisting of beta strands connected laterally by three or more hydrogen bonds, forming a generally twisted, pleated sheet

Front 70

rate law

Back 70

r=k(T)[A]^m[B]^n
where, k(T)=rate constant.

Front 71

Arrhenius Equation

Back 71

k=Ae^(-Ea/RT)
gives the dependence of the rate constant k of chemical reactions on the temperature and activation energy

Front 72

Exceptions to the octet rule

Back 72

boron compounds such as BF3 have only 6 electrons in the valence shell;free radicals, phosphorus, fluoride

Front 73

monomer

Back 73

a small molecule that may become chemically bonded to other monomers to form a polymer

Front 74

dipole moment

Back 74

the measured polarity of a polar covalent bond. It is defined as the product magnitude of charge on the atoms and the distance between the two bonded atoms.

Front 75

VSPER theory

Back 75

used for predicting the shapes of individual molecules, based upon their extent of electron-pair electrostatic repulsion, determined using steric numbers;predict the geometric shape and lone-pair behavior of a molecule one must consider : the constructed Lewis structure, expanded to show all lone pairs of electrons, alongside protruding and projecting bonds

Front 76

intermediates and catalysts

Back 76

catalysts are not consumed
intermediates are made and comsumed during the rxn

Front 77

exothermic

Back 77

releases heat/energy

Front 78

enzyme

Back 78

biomolecules that catalyze chemical reactions. Almost all enzymes are proteins.

Front 79

substrate

Back 79

binds with the enzyme's active site, and an enzyme-substrate complex is formed. substrate is broken down into a product and is released from the active site.

Front 80

inhibitor

Back 80

molecules that bind to enzymes and decrease their activity.

Front 81

enthalpy

Back 81

ΔHrxn= ΣΔH(products)-ΣΔH(reactants)
ΔHrxn= Σ(bonds broken)-Σ(bonds formed)

Front 82

enthropy

Back 82

Increases when matter and/or energy spread out.Reflects disorder in a system.
ΔS=Sfinal-Sinitial=qrev/T

Front 83

Gibbs free energy

Back 83

ΔGsys=ΔHsys-ΔSsysT
ΔG<0: product favored (exo)
ΔG>0: reactant favored (endo)

Front 84

Equilibrium Constant

Back 84

aA+bB->cC+dD
K= ([C]^c *[D]^d)/([A]^a*[B]^b)
**solids do not appear
**in dilute sol'ns solvent does not appear.

Front 85

Kc

Back 85

Kc<1: reactant favored
Kc>1: product favored
Kc=1: prod & reac comparable
Kc=Keq

Front 86

Kp

Back 86

Kp=Kc(RT)^Δn

Front 87

Hess's Law

Back 87

The enthalpy of a rxn depends on the states of the reagents. Allows for predictions. ΔHoverall= ΣΔHsteps

Front 88

Reaction Quotient

Back 88

same form as the K, except use instantaneous concent @time T.
K>Q:Forward
K<Q:Reverse
K=Q:Equilibrium

Front 89

Kinetically Stable

Back 89

If Ea is large. If Thermodynamically stable, kinetics doesn't matter.

Front 90

Thermodynamically Stable

Back 90

If ΔG>0 with respect to the product.