New Chem/Phys Flashcards
1
Q
proton
A
positive charge and a mass of 1 amu
2
Q
neutron
A
no charge and mass of 1 amu
3
Q
electron
A
negative charge and mass of 0
4
Q
nuculeus
A
contains protons and neutrons, electrons flot around nucleus
5
Q
atomic number
A
number of protons in a given element
6
Q
mass number
A
sum of elements protons and neutrons
7
Q
atomic mass
A
equal to the mass number, sum of protons and neutrons
8
Q
isotopes
A
atoms of a given element (same atomic number) but have different mass numbers. differ in number of neiutrons
9
Q
atomic weight
A
weighted average of the naturally occurring issues
periodic table lists these
10
Q
rutherford
A
first said that the atom had a dense, + charged nucleus made up only a small fraction of the volume of the atom
11
Q
Bohr model of an atom
A
dense positive charge nucleus surrounded by electrons that revolve around a nucuekus in orbit with distinct energy levels
12
Q
the energy difference between levels is called a
A
quantum
13
Q
quanitization
A
not an infinite range of energy levels available to an electron, electrons can exist only at certain E levels. E of an electron increases as it is farther from the nuc
14
Q
atomic absorption spectrum
A
electron to jump from a lower E level to a higher E level, must absorb the amount of energy equal to the energy difference between the 2 levels
15
Q
when electrons return from the excited state to ground state, they
A
emit the amount of energy that is exactly equal to the energy difference between the 2 levels. atomic emission spectrum
16
Q
quantum mechanical model
A
electrons do not travel in defined orbits, but localized in orbitals
17
Q
orbitals
A
a region of space around the nucleus defined by the probability of finding an electron in that region
18
Q
Heisenberg uncertainty pricniple
A
impossible to know an electrons position and its momentum exactly at the same time
19
Q
quantum numbers
A
n, l, m1, ms
20
Q
principal quantum number
A
n. describes the energy of a shell.
21
Q
azimuthal quantum number
A
l. describes the sub shells within a given principal energy level (s, p, d f, )
22
Q
magnetic quantum number
A
m1 specifics the particular orbital within a subshell where an electron is likely to be found at a given moment in time
23
Q
the spin quantum number
A
ms. indicates th spin ordinations with +1/2 and -1/2 of an electron in an orbital
24
Q
electron confiuhration
A
ex. magnesium. 12 electrons
1s2, 2s2, 2p6, 3s2…
3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6
f has 14
25
electrons fill the principal energy level and subshells by this role
n +1 rule
26
HUnds rule
sub shells with many orbitals (p, d , f ) fill electrons so that every orbital in a subshell gets one electron before any of them get a second
27
paramagnetic
have unpaired electrons that align with magnetic fields, attracting material to a magnet
28
dimagnetic
materials have all paired electrons which cannot easily be realigned, repelled by magentis
29
valence electrons
electrons in outermost shell available for interaction (bonding) with other atoms
30
the difference in energy between 2 shells decreases as teh distance from the nucleus inceases
energy difference between 3 adn 4 s less than energy difference between 1 and 2
31
maximum number of electrons in a shell
2n^2
32
m1 can only be between -L and +L
yes
33
subshell s, p, d, f
L:
0,1,2,3
electrons:
2,6,10,14
34
unpaired electrons in the ground state example
helium . repelled by magments
35
E=hc/wavelength
| energy
h= 6.6 x 10 ^-34
c=3.00 x 10^8-speed o light
wavelength of light
36
if you do something with "1 mole"
multiple by avogadros number
| 6.02 x 10^23
37
spins in the same direction
parallel
38
periodic table organizes elements according to their
atomic number
39
rows are
called periods and based on teh same pricipal energy level (n)
40
columns are called
groups, have the same valence shell configuration
41
metals
shiny, lustrous, conduct electricity, malleable and ductile. metals on L side and middle
42
nonmetals
dull, poor conductors, brittle, R side
43
metalloids
metals and nonmetals and with boron stepping pattern
44
Zeff
effective nuclear charge- net positive charge experienced by electrons in their valence shell
increases from L to R,
valence electrons become increasingly separated from the nucleus as the principal energy level (n) increases from top to bottom in a group
45
atomic radius
decreases from L to R acrosss a period and increases from top to bottom in a group
46
ionic radius
size of charged species. increases as you move from top to bottom on the periodic table. Ionic radius decreases as you move across the periodic table, from left to right. cations smaller than neutral, anions greater than neutral
47
ionization energy
amount f energy needed t o remove an electron from the valence shellof a gas - increases from L to R and decreases from top to bottom
48
electroneg
F!!! measure of attractive force of the nucleus for electron within a bond. increases from L to R across a period and decreases from top to bottom
49
alkali metals
oxidation state of +1 and prefer to lose an electron to achieve Nobel gas state- very rective
all the way on the outside of L
50
alkaline earth metals
oxidation state of +2 and can lose 2 electrons to achieve noble gas configurations. reactive
form divalent cations (ions with +2 charge)
51
halogens
oxidation state of -1 and can gain electrons too have noble gas configurations.
highest electroneg
52
noble gases
fully filled valence shell and prefer to not give up
| high ionization E (for HE, NE and AR) with no electroneg and electron affinities
53
metalloid
between metals and and transition metals
54
what forms with water well
transition metals. because have many oxidation states
55
within the same period, an additional valence electron is
added with each step to teh R of the table
56
ionic bonds
unequal sharing of electrons. formed via the transfer of one or more electrons from an element with low ionization E to an element with high electron affinity.
occur between elements with large differences in electroneg (metals and nonmetals)
57
elements with an incomplete octet that are stable with fewer than 8 electrons
H, He, Li, Be, B
58
ionic ions dissociate in
polar solvents and water
59
ionic ions have high
melting point
60
covalent bond
formed via the sharing of electrons between two elements of similar electroneg.
61
bond order
refers to whether a covalent bond is a single bond, double or triple. as bond order increases, bond strength increases, bond energy increases and bond length decreses
62
non polar bonds covalent bonds
atoms that have the same electroneg
63
polar bonds covalent
significant difference in electro but not enough to transfer electrons to form an ionic bond. the more electroneg atom takes on the partial negative charge
64
coordinate covalent bond
single atom provides both bonding electrons while the other atom does not contribute any.
Lewis acid base chem
65
formal charges
when an atom is surrounded by more or fewer valence electrons than it has in neutral state (assuming equal sharing of electrons in a bond)
V-Nonbonding-1/2bonding
valence electrons
66
lewis acid
any compound that will accept a lone pair of electrons
67
lewis base
any compound that will donate a pair of electrons to form a covalent bonds
68
resonance
any molecule with pi system or duke bonds
represent all possible confirmations of electrons
small or no formal charge > formal charge
less separation between opposite charges > large seperation
negative formal charge is on electroneg atom > negative formal charge on less electroneg atom
69
all elements beyond third period can take more
than 8 electrons in their valence shell and form more than 4 bonds
70
VSEPR theory
molecular geometry
determined by repulsions between bonding and nonbonding electrons
linear (BeCL2), trigonal planar (BH3), tetrahedral (CH4),trigonal byprimada (PCL5), octahedral (SF6)
180,120,109.5, 90/120/180, 90/180
71
electronic geometru
describes the spatial arrangement of all pairs of electrons around the central atom, including bonding and lone pairs
H20, CH4 and NH3 all have 4 pairs of electrons around the central atom- tetrahedral
72
molecular geometry
describe teh spatial arrangement of only the bonding pairs
H20, CH4 and NH3- bent, tetrahedral; and trigonal pyramidal because different coordination number
coordination number- the number of atoms that surround and are bonded to a central atom
73
CCL4- non polar or polar?
non polar. although CL is electroneg and pulls, cancels each other out in all 4 directions
74
when orbitals overlap head to head
sigma bond
| free rotation of axes because electron density is single linear
75
two parallel electron cloud
pi bonds. no free rotation because parallel and cannot twist
76
intermolecular forces
electrostatic interactions between molecules
| weaker than covalent bonds (which are weaker than ionic)
77
LDF
weakest, present in all
| as the size of the atom increases, so does LDF
78
dipole dipole
occur between oppositely charged ends of polar molecules.
| not in gas phase due to distance in particles
79
hydrogen bonds
specialized idols
| H is bonded to F, O N- electroneg atoms
80
large intermolecular forces
higher BP
Kr= octet
ionic bonds are strongest
81
bond lengths decrease with larger differences in
electronegative
82
ammonium
NH4+
83
ammniia
NH3
84
ion dipole interaction
the result of an electrostatic interaction between a charged ion and a molecule that has a dipole.the interaction between a Na+ ion and water (H2O) where the sodium ion and oxygen atom are attracted to each other, while the sodium and hydrogen are repelled by each other.
85
mole
quantity of any substance (atoms, molecules, etc)equal to the number of particles that are found in 12 g of C
defined as avoados number - 6.022 x 10^23
86
molecular weight vs. molar mass
amu/molecule and g/mol
87
moles =
mass of sample (g)/molar mass (g/mol)
88
gram equivalent weight
molar mass/n
n= number of particles of interest produced or consumed per molecule of the compound
ex. one would need 31 g of H2Co3 (molar mass 62 g/mol) ti produce an equivalent of hydrogen ions because each mol of H2Co3 can donate 2 hydrogens (n=2)
89
equivalents
mass compound (g)/gram equivalent weight (g)
90
normality
mesure of concentration
equivlants/L
hydrogen ion concentration
1N solution = hydrogen ions equal to 1 mole per liter. 2N acid solution = 2 moles of hydrogen ions per liter
91
in a 1N HCL solution the molarity of HCL is
1 M because HCL is monoprotic
92
in a 1 N H2CO3 solution the molarity is
0.5 M because diuretic acid
93
normality
molarity x n
| n= numbest of protons, ions, electrons consumed by the solute , H+ ions or OH IONS
94
molarity=
mol/L
95
empirical formua
simplest whole number ratio of the elements in the compound
| CH- benzene
96
molecular formula
exact number of atoms of each element in the compound and is a multiple of the empirical formula
C6H6- benzene
97
percent composition
mass of element in formula/molar mass x 100
ex. what is the percent composition of Cr in K2Cr2O7
molar mass = 292 g/mol
Cr: 2 x 52
98
combination reaction
has two or more reactants forming 1 products
99
decomposition reaction
opposite of combination
| single reactant breaks down into 2 or more products
100
combustion rection
hydrocarbon (fuel) and oxidant (oxygen) to form CO2 and water
Ch4 + 2 O2 --> CO2 + h20
101
single displacement
when an atom or ion in a compound is replaced by an atom or ion of another element
Cu + AgNO3 -- Ag + CuNO3
102
double displacement
metathesis reactions
| each swap
103
neutralization reactions
acid with base to produce a salt.
| Hcl and NAOH will produce NaCL and H20
104
limiting reagent
limits the amount of product that can be formed in the reaction because it is used up first
105
theoretical yield
max amount of product that can be generated as predicted from the balanced equation, AAUMIGN ALL LIMITED REATANT IS consumed.
106
actual yield
the amount of product one actually does obtain during rection
107
percent yield
actual/theoretical x 100
108
cations
usually metas
109
anions
usually nonmetals
110
oxidation states
different charges
111
solutes that enable solutions to carry currents
electrolytes
112
-ous
lesser charge
113
-ic
greater charge
114
The solute concentration will affect the size of spots on a paper chromatogram, .
not the average migration rate
115
Rf=
= (distance traveled by aspatic acid)/(distance to the solvent front)
116
during a dehydration reaction
a carbocation forms, which leads to non stereospecific (trans and cis) results
117
What type of functional group is formed when aspartic acid reacts with another amino acid to form a peptide bond?
AMIDE GROUP
The functional group that represents a peptide bond is called an amide group.
An amine is one of the functional groups that reacts to form a peptide bond, not the peptide bond itself.
118
1,2-hydride shift
a carbocation rearrangement in which a hydrogen atom in a carbocation migrates to the carbon atom from an adjacent carbon (carbon 1).
119
In gas-liquid chromatography, the first peak to emerge will be from the least polar, most volatile compound.
hydrogen bonds= non volatile (like in alcohols)
dipole dipole (like with BR and CL)- less volatile
LDF= volatile (bonde to just a methyl group
120
1 cm=
.01 M
121
1 m
100 cm
122
40 cm=
0.4 M
123
100 M
1000 cm
124
The negative sign of the lens strength is indicative of a
diverging lens
125
enzymes affect chemical reactions by Stabilization of:
the transition state changes the activation energy of the reaction.
126
ionization is
ionization is the loss or gain of charge through H atoms.
127
ohms law
V=IR
128
index of refractio
=speed of light in vacuum/ the speed of light in the medium,
n= c/v
129
E (photon E)=
hf
h= planks constant = 6.626 * 10-34
f=frequnecy
130
speed of light =
frequency x wavelength
131
intensity of the radiation
is rested to number of photons emitted
132
THz=
1 x 10^12 Hz
133
standard atmospheric pressure is
760 mmHg or 1 atm
134
radioactive decay law
N(t) = N (1/2)t/T
135
1 mL
0.001 L
136
work and power equation
P = W/t
| time is in secondos
137
The high number of capillaries in the body means that the total cross-sectional area of these vessels is larger than any other vessel type in the circulatory system.
This causes the velocity of the blood to decrease.
138
BP is higher in the
arteries than capillaries
139
Which of the following types of orbitals of the central atom are involved in bonding in octahedral compounds?
Octahedral compounds have six σ bonds and no stereochemically active lone pairs. According to valence bond theory, the central atom requires the hybridization of six atomic orbitals, d2sp3
when there are more than 4 bonds, the sp3 hybridized orbitals can't accommodate the number of bonds, so it need to recruit the d orbital?
octahedral molecules have six atoms bonded to the central atom and no lone electron pairs,
140
The molar volume of an ideal gas at 25°C is 24.4
not 22.4
141
Retinal is composed of mainly carbon and hydrogen
making it largely hydrophobic.
142
gamma decay
emission of a proton
143
negative delta G means
positive delta S for sponteous reaction
144
Gibbs free E
determines whether or not a reaction will occur by itself without outside assistance
not necessarily mean quickly -no help with catalyst
145
colliosn theory of chemical kinetic
the rate of the reaction is proportional to the number of collisions per second between reacting molecules
146
activation E
the minimum energy of collision necessary for a reaction to take place
147
Arrhenius Equation
```
k- rate constant of reaction
A= frequency factor
Ea= activation E
R- ideal gas connate
T- temp
```
k= Ae^ (-Ea/RT)
148
rate of reaction
increases with temperature
| more molecules present
149
transition state
has greater E than both the reactants and the products
| energy required to reach here is the activation E. can either go to products or reactants
150
free energy change of the reaction
difference between the free E of products and of reactants
-= exergonic
+= endogeronic
151
exergonic
energy is given off
152
endergonic
energy is absorbed
153
reaction rates are measured in terms of
the rate disappearance of a reactant or appearance of a product
154
rate laws general form
k[A]^x[B]^y
x and y are determined EXPERIMENTALLY not coefficients
155
rate is measured in
molarity (M)/seconds (s)
156
zero order
formation of product C is independent of changes in concentration of any of the reactants
constant reaction rate equal to rate constant K
A and B raised to 0.
change by adding a catlyst or temperature
linear decreasing graph
157
first order reactions
rate is directly proportionate to only one reactant
A and B raised to 1
slightly curved line
158
second order reaction
propinate to the concentration of 2 reactants or square concentration of a singlee reactant
A and B raised both to 1 or A or B raised to 2
159
mixed order reactants
catalyst is invovled
| reactions with rate orders that vary overt time
160
if they tell u that the rate is "third order"
the sum of the exponent in the rate law must be equal to 3. s
161
changing the concentration of the reactants will not change the rate in a
0 order reaction rate
162
solvents affect the rate of reaction
along with temperature, partial pressure of gas (how molecules collide) but products have nothing to do with reaction rate so changing that is fine
163
saturated solutions with a catalysts have a max turnover rate and cannot increase the rate constant or
the reaction rate any higher by adding more reactant
164
the overall order of a reaction is the
sum of the individual orders of the reaction
165
lowest activation eneegru means
fastest rate
166
the law of mass action
for a reversible reaction, if the system is at equilibrium at a constant temperature than
Kc=Keq= concentration of products raised to their coefficient/reactants raised to their coefficient
167
forward and reverse reaction rate constants
Kf and Kr
| Keq= Kf/Kr
168
reaction quotient
at any point in time during the reaction, can measure concentrations of all the recctants and products using the same Keq formula
concentration of products raised to their coefficient/reactants raised to their coefficient
169
Q< Keq
forward reaction has not yet reached equilibrium
reactants are greater in concentration
forward rate increased to restore equilibrium
170
Q=Keq
reaction is in dynamic equilibrium
| forward and reverse are equal
171
Q>Keq
forward reaction exceeded equkiriu
greater concentration of products
reverse rate is increased
172
if teh equkibirum for a reaction is written in one direction to be keq
reverse reaction. is 1/keq
173
Keq is
temperature dependent
174
Keq>1
products are present in greater concentration at equlibrium
175
Keq=1
products and reactants are both present at equilibrium at simiar level
176
Keq<1
reactants are present in greater concenrtation
177
Keq<<1
amount of reactants converted to products is very small
178
Keq=
x^2/1-x
x on the denominator is negligible
when Keq is largely negative
179
Le Chateliers
when a chemical system experiences a stress, it will rect so as to store equilibrium
stress can beL changes in concentration, pressure, volume and temperature
180
increasing gate concentration of reactants or decreasing products will shift
to the right
181
increasing products
shift to left
182
increasing pressure on gas system (decreasing volume)
shift to fewer moles of gas side
known by coefficient
183
decreasing pressure (increase volume)
shift reaction to side with more moles of gas
| known by coefficient
184
increase temp of endothermic reaction or decreasing temp of exothermic reaction
reaction shift to R
185
decreasing temp of endothermic or increasing temp of exothermic
shift to the left
186
kinetic products
higher in free E than thermodynamic products adn can form at lower temp. f
fast products because they can form more quickly under such conditions
free E is lower than thermodynamic products
187
thermodynamic products
lower in free E and more stable than kinetic
| more sponetous tho
188
solids are NOT IN EQUILIBIRUM EXPRESSSIONS
ONLY GASES
189
Kc
concentration equilibrium
190
Kp
pressure equilibrium
191
exothermic reaction produces heat so decreasing the temperature
favors product formation, increase of forward reaction
192
adding or removing a catalyst would change the reaction rates but not
where equilibrium is
193
at very high temp
reactants and products may decomponse
194
negative delta H
exothermic
195
mRNA is moncistronic
meaning that each mRNA molecules translates Ito one protein product
196
in prok mRNA is polycistronic
mRNA can result in many different proteins depending on the lcoation
197
tRNA
convertingg NA to amino acid and peptides RRNA
198
rRNA
ribosomal machinery used during protein assembly in teh cytoplasm
199
the genetic code is degenerate becuaas e
more than 1 codon can specify a single amino acid
200
wobble position
third codon base
201
X^0=
1
202
X^-A
= 1/X^A
203
log 1 =
0
204
10^12=
tera
205
10^9
giga
206
10^6
mega
207
10^3
kilo
208
10^2
hecto
209
10^1
decta
210
10^-1
deci
211
10^-2
centi
212
10^-3
mili
213
10^-6
micro
214
10^-9
nano
215
10^-12
pico
216
isolated ysstem
system cannot exchange energy (heat and work) or matter with teh surroundings
217
closed system
the system can exchange energy (heat and work) but not matter with teh surroundings
218
open system
both energy and matter can exchange with surroundings
219
first law of thermodynamics
delta U = Q-W
u= change in internal energy
Q= heat added
W= work done by the system
change in internal energy is equal tot eh amount of heat transferred minus energy trsnadfered from work
if intneral E is positive, then increasing temperature, heat flows int eh system and work is done by teh system (expansion)
if negative internal E, decreasing temp, heat flow out of system and work is done on the system (compression)
220
isothermal processes
constant temperature so u is constant
(temperature and interval energy are directly proportional)
so Q=W
(in the delta U = Q-W equation)
Hyperbolic curve of pressure adn volume with work underneath that cuve
221
adiabatic processes
when no heat exchange between the system and the environment so thermal energy is constant
Q=0
delta U= -W (work done on the system)
also hyperbolic on the graph
222
isobaric processes
when the pressure of the system is constant.
| flat lien on the P and V graph
223
isovolumeric process
isochoric. experience no change in volume
gas neither expands nor compresses so no work is preformed so
U=Q
vertical line on PV graph
224
standard conditions
25 degrees C (298k), 1 atm pressure, 1 M concentrations
| BU STTANDARD TEMP AND PRESSURE IS 0 DEGREES C (IDEAL GAS)
225
phase changes
melting and freezing (crystallization and solidification) occur at boundary between solid and liquid
226
vaporization (evaporation or building) and condensation
occur at liquid and gas phase
vaporization: liquid to GAS
condensation: gas to liquid
227
sublimation and deposition occur at
solid and gas phase
solid to gas: sublimination
gas to solid: deposition
228
between liquid and gas phases is the
critical point- where liquid and gases are indigustushable
229
at the triple point
all three phases exist at equilibrium (the center of the graph)
230
phase diagram
qshows each phases according to pressure and temperature
231
temperature
is a scaled measured of the average kinetic energy of a substance
232
heat
transfer of energy that results from differences of temperature between 2 substances
233
the heat content of a system undergoing heating, cooling or phase changes is the
sum of all energy changes
234
process of a systematic absorbing heat
endothermic
| Q>0
235
process of system release heat
exothermic
| Q<0
236
enthaloy
is equivalent to heat (Q) oder constant pressure
237
q=
```
mc delta T
MCAT
m- mass
c- specific heat of substance
q=heat
delta T= temp
```
238
specific heat
amount of energy required to raise the temp one gram of suabyance by 1 degrees C
239
specific heat of water
1 or 4.184
240
during phase changes (solid to liquid or liquid to gas), we cannot use mCAT cuz delta T is 0
q=ml
q=heat
m=mass
l- latent heat-enthalpy (heat of fusion or heat of vaporization)
241
when going from solid to liquid, the changes of enthalpy will be
negative because heat is removed
242
enthalpy
measure of the PE of a system in intermolecular attractions and chemical bonds
=U + PV
=Q-W + PV
heat, work done by the system,
243
Hess's law
the total change of PE of a system is = to the changes of PE of individual steps int eh processes
244
second law of thermodynamics
energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing o
245
entropy
measure of sponetous dispersal of energy at specific temperatures
DELTA S = Q rev (heat that is gained or lost in a reversible process)/T (temp)
J/mol K
always greater than 0
246
when energy is distributed in a system at a given temp
entropy increases
247
delta G= delta H-T delta S
Gibbs free energy equation
248
delta H positive and delta S potivie
sponteous at high T
249
delta H positive and Delta S neg
nonspotous at all T
250
delt H neg and delta s pos
sponteous at all T
251
both delta H and delta S neg
sponetous at low T
252
boiling is an
endothermic process and delta H is motive
| abd delta S is positive
253
delta G reaction
-RTlnKeq
R- ideal gas constant
T-temp in Kelvin
Keq- equilibrium connate `
254
gases have highest entropy
solids have the lowest
255
gases are
least dense, easily compressible
256
1 tm =
760 mmHg-760 torr=101 kPA = 1.013 Pa
257
mercury barometer
measures atm pressure
| as pressure increases more mercury forced into column, increasing its height
258
STP
```
273 K (0 C) and 1 atm
1 mol of gas occupies 22.4 L
```
259
equations for ideal gases assume negligible mass and volume of
gas molecules
260
regardless of identity of gas, equimolar amounts of 2 gases will
occupy the same volume at the same temrpature and pressure
261
ideal gas law
a moleucle with no intermolecular forces and occupy no volume
PV=NRT
n- number of moles
R- ideal gas constant which is 8.21 x 10^-2
262
C to Kelvin
add 73
263
pressure of 380 mmHg
divide it by 760 mmHG to get atm
264
grams to moles
use molar mass (mols/g)
265
density
ratio of mass/volume =PM (molar mass)/RT
266
combined gas law
P1V1/T1=PV/T
267
avogadros principle
all gases are at a constant temperature and pressure occupy volumes that are directly proportional to the number of moles present
n1/V1=n2/V2
mols of gas /volume `
268
Boyles law
PV=PV
| inverse relationship I
269
Charles Law
V/T= V/T
direct relationship
270
Gay Lussac
P/T=P/T
| direct relationship
271
law of partial pressure
individual gas components of. MIXTURE OF GASES WILL EXERT INDIVIDUAL pressures in proportion to their mole frctions
total pressure of mixture of gases is equal tot eh sum of their partial pressure of many gases
272
partial pressure of a gas
Pa (pressure of A)=XPt (total pressure)
X- moles of gas A/total moles of gas
273
KE=
1/2mv^2
274
kinetic molecular theory
gases are made up of particles with low volume
gases have no intermolecular attractions or repulsions
gas are random, collisions
collisions are elastic (conversation of momentum and KE)
avg KE of gases are proportional to temp
275
photoelectric efect
when light of high Frequency (blue to UV light) is shown on a metal vacuum, metal atoms emit electrons. these electrons will produce a net charge flow (current)
276
larger intensity means
greater current and the greater the number of photons per unit time and greater number of electron per unit time
277
current is proportional to
intensity (amplitude) of the light
278
threshold freqnecy
minimum frequency of light that causes ejection of electrons
if less, no electron will be ejected
279
energy of each photon equation
E=hf
E- energy of 1 photon of light
h-plancks constant (6.6 x 1^-34
f- frequency of the light
280
wavelength adn frequency equation
c=f wavelength
| c- speed o light (3.0 x 10^8
281
waves with higher frequency have shorter wavelengths and
```
higher energy (toward blue and UV)
inverse of each other
```
282
waves with lower frequency have longer wavelengths and
lower energy (to3ward red and infrared spectrum)
283
if the photon has more than enough energy to eject a single electron (above threshold NOT at threshold ) then the excess energy will be converted to
kinetic energy in the ejected electron
KE= hf-W
w- work
284
work function
minimum energy required to eject an electron
| W=hf
285
an electron can jump from a lower to a higher orbital b y
absorbing a photon of light at the right frequency to match teh energy difference between oribital s(E=hf)
286
when an electron falls from a higher energy level to a lower one
photon of light is emitted
| emitted with lower frequency because lower energy so longer wavelengths
287
IR spec
is used to determined chemical structure because different bonds absorb different wavelengths of light
288
UV VIS spectra
absorption of light in the visible and UV range of chemical structures
289
we see the color on IR and UV spec that are
not absorbed
290
photons of UV light have
high frequency (Short wavelengths)
291
mass defect
E=mc^2
E- energy
m-mass
c-speed of light
matter is converted to energy, so nuclei is slightly smaller than the protons and neutrons added together
small amount of mass can have a lot of energy because of the c sqaured
292
isotopic notation
A X
Z
Z- atomic number- protons in the nucleus
A- mass number- protons plus neutrons
293
fusion
small nuclei combine and form a larger nuclei
294
fission
large nucleus splits into smaller nuceli
295
alpha decay
emission of an alpha particle, which is 4/2 HE (alpha)secular that has 2 proton 2 neutrons and 0 electrons
massive compared to beta and gamma `
296
beta decay
emission of a beta particle, which is an electron
| 0/-1 Beta
297
positron emission
position is released, which has the same mass and effect of an electron in beta decay but carries a positive charge
298
gamma decay
emission of gamma rays which are high energy and high frequency photons
carry no charge and Lower the energy of the nucleus without changing the mass number or atomic number
0/0 gamm
299
half life example
if the half life is 4 years, what fraction of the sample will remain after 12 years
do teh normal half life thing
3 half LIVES at 12.5% remains
(1/2)^3= 1/8 is the fraction
300
mass defect is related to binding energy
E=Mc2
301
the flow of positive charges
current
302
metallic conductivity
metals (good electrical and thermal conductors because electrons can move around)
303
electrolytic conductivty
seen in solutions conductivity
water has low conductivity (insulator)
sea water is good conductor
304
electrical current
Flow of charge between two points at different electrical potentials connected by a conductor, like a copper wire
305
Current equation (I)
I= Q (amount of charge passing through a conductor)/time
| amps
306
because electrons are negatively charged, they move from a point of lower electrical potential to a point of
higher electrical poneital (reducing their electric potential E)
the pathway that positive charges flow, or current, actually flows from higher potential to lower potential and so current is the opposite of the direction of electron flow
307
voltage
can be produced by an electrical generator, a voltaic cell (galvanic) or cells with a battery
308
Kirchfoos junction rule
at any point in a circuit, the sum of currents directed into that point equals the sum of currents directed away from that point
I (into)= I out
309
Kirchoffs loop rule
in a closed loop, the sum of voltage sources is equal to the sum of voltage drpos
310
resistanve
the opposition within` any material to teh movement and flow of charge
think of it like air resistance, friction
311
materials that have low resistance are called
conductors
312
Resistane equation
R= p (resistivity) x L (length of the resistor)/ A (cross sectional area)
313
resisitiy is in
ohms
| intrinsic resistance to current flow
314
a longer resistor mean that electrons will have a greater length to travel in its resistant material
L in resisvity equation
315
inverse between resistance and cross sectional area
an increase in area will increase the number of pathways through the resistor, and more current can flow the wider it is.
316
conductor have greater resistance at
higher temp
317
ohms law
V=IR
| voltage, I- current, R- resistance in OHMS
318
in a series
current is equal
| add the resistance up up
319
power adn resistance
P= IV= I (R)^2=V^2/R
320
parallel circiut
current will divide to pass through resistors seperately
so add them together
321
in a series circuits, the resistance are
added together
| as ell as the voltages
322
parallel ciruits
voltage is the same for every resistor
current can be added together
resistance ill be different (electrons prefer the path of least resistance)
add together resistor by 1/R + 1/R = R total (which will decrease as more R are adde)
323
capacitors
the ability to hold charge at a particular voltage
| defibillaterors!
324
captiance
the ratio of the magnitidue of the charged stored on one plate to the potential difference (voltage) across the capacitor
both positive and negative plates
C=Q/V
Q- charge
V- volrage
C- capacitance
Farad
325
uniform electric field equation
E= V/d
d- distance
E- electric field
326
potential energy of a stored capictor
PE = 1/2CV^2
327
when capacitors are connected in series, the total capacitance decreases
as more are added
must share the voltage drop and therefore cannot store as much charge
1/C total = 1/c + 1/c....
328
capitanors in parallel
C total = C1 + C2 + C3
| C Toal increases as more is added
329
ammeters
used to measure the current at some point within a circuit
330
voltmeter
measure voltages
331
ohmmeters
measure resistance
332
amps
C/s
333
sinusoidal waves
oscillate back and forth
334
traverse vs longuidundal waves
Transverse waves cause the medium to move perpendicular to the direction of the wave (up and down). Longitudinal waves cause the medium to move parallel to the direction of the wave (sound)
335
the distance from one crest (maximum) to the next
wavelength
336
frequency of a wave
number of wavelengths passing a fixed point per second
| Hz
337
propagation speed of wave equation
v= f wavelength
338
period (T)
inverse of freuency
| 1/f
339
angular frequqncy
radians per sec
used ins simple harmonic motion in springs
angular frequency = 2 pi F or 2 pi/T
340
maximum magnitude of displacement in a wave is
amplitude
| from equilibrium position (middle) to crest (top) or trough (bottom)
341
in phase
2 waves that have the same frequency wavelength adn amplitude through teh same space at same Tim because crest adn trough line up
phase difference is 0
342
phase outout of phase
when crest of 1 wave goes with teh trough of another
phase difference is 1/2 of a wave
wavelength/2 or 180 degrees
343
principle of superposition
when waves interact with each other, the displacement of the resultant wave at any point is the sum of the displacement of 2 interaction waves
344
constructuve interence
when the waves are perfeclty in phase and the displacement add together and the amplitude is equal tot eh sum of the amplitude of the two eves
345
when waves are perfectly out of phase, displacement counteract each other and the amplitude is the difference between teh 2 waves
destructive interference
346
string is fixed at one end moving up and down
traveling wave and will reflect and invert when rect the end
| two waves shown
347
when both ends of her string are fixed, the traveling wave will look
stationary, looking like standing waves
348
points in teh wave that remain at rest
nodes (amplitude is 0)
349
antinodes
midway between the nodes flucttuate with maximum amplitude
350
audible frequency
20 Hz-20,000 Hz
351
sound
longuidindal wave trnmitted by the oscillations of particles in a deformed medium
produced by the mechanical disturbance of particles in a material
vibrate
352
pitch
frequency
| determined by the length of the air column, which can be changed ether by covering holes or changing its lenght
353
doppler effet
the differnec between teh actual frequency of a sound and its perceived regency of the sound and the sounds detector are moving relative to one another
towards each other- perceived is greater than actual
away from each other (source and detector)- perceived is less than actual
perceived = actual (v -/+ vD/v -/+ vs)
v- speed of sound
vD- speed of detector
vS- speed of source
first sign os used when going towards (NEG), second is when it away (+)
354
km/hr to m/s
1 hour has 3600 seconds
| 1000 m in 1 km
355
loudness
or volume of sound is the way in which we perceive its intensity
356
intensity
average rate of E transfer per unit area across a surface that is perpendicular to then wave
I = P/A
p- power
A-area
357
intensity is proportional to the square of ampltidue
so doubling amplitduing means quadrupling intensity
358
sound level
B= 10 log I/Io
359
change ins ound level
Bf= B9 + 10 log If/Ii
360
damping
attenutation
| sound will be lower than expected from calculations
361
standing wave
L= wavelength/2, 2 wavelength /2, 3 wavelength over 2
or wavelength = 2L/n
n=1,2,3,etc (harmonics)- corrresponds to the number of half wavelengths (HUMPS)
IN ESSSENCE
wavelength = 2 L/ count how many waves (OR HUMPS)
362
open pipes
```
has antinodes on both ends
(nodes are the only thing there for closed pipes)
L= 1wavelength/2
L= 2wavelength/2
L= 3 wavelength/2
```
SAME AS STRING
wavelength = 2L/n (COUNT THE ANTINODE)
363
closed pipes
the harmonics are now equals to eh number of quarter wavelengths supported by teh pipe
odd harmonics only
wavelength = 4L/n or
L= wavelevnth/4
L= 3 wavelength /4
L=5 wavelength/4
364
ultrasound
uses high frequency sound waves to compare relative densities of tissue sin body
based on travel time of reflece sound
365
speed of sound is fastest in solid
then liquids than gases
366
intensity is related to
sound level (which is related to amplitude of vibration)
367
mass unts
lb s^2/ft OR lbs s^2/in
368
force units
g cm/ s^2 OR kg m/s^2 (N)
369
work and energy unit
ft lb OR g cm^2/s^2 OR kg m^2 / s^2 (JOULES)
370
POWER unit
ft lb/sec OR g cm^2/s OR kg m^2/s^3 (WATT)
371
1 angstrom =
1 x 10^-10
372
1 eV =
1.6 x 10^-19
373
v=
d/t
374
friction =
mu N
mu- kinetic friction coeff
N- normal force- mg
375
a=
v/t
376
v=
v0 + at
377
d=
v0t + 1/2at^2
378
v^2=
v0^2 +2ad
379
d=
vt
380
Fc=
mv^2/r
381
kn to N
1 kn = 1 x 10^3
382
TORQUE
r f sin theta
383
zeroth law of thermodynamics
objects are in thermal equilibrium when they are at the same temperature adn experience no net change of heat energy
384
temperature
qualitative measures how hot or cold something is
| quantitively, it is related to KE of the particles that make up a substance
385
thermal expansion
how a substate changes in length or volume as a function of the change in temperature
386
heat moves from materials that have higher tempature
to materials that have lower temperature
| heat is the transfer of thermal E
387
kelvin scale
starts from 0, adds 273 + C
| no negatives
388
Fahrenheit to Celsius equation
F= 9/5C +32
389
rising temperature causes a decrease or increase of length
increase of length for solids, and increase in volume for liquids
390
1 Cal=
1000 cal = 4184 J
391
conduction
transfer of E from molecule to molecule through molecular collisions
must be physical direct contact between objects
transfer KE from hot to cooler matter through collisions
392
metals are the best conductors because
metallic bonds contain a density of atoms embedded in sea of electrons which facilitate rapid energy transfer
gases re the poorest because too much space between molecules that collisions are infrequency
393
convection
transfer of heat by the physical motion of a fluid over a material
ovens
394
radiation
transfer of E by electromagnetic waves
through a cavvumn
sun able to warm earth
395
heat of fusion
liquids to solids, solids to liquid (which have a melting point)
396
heat of vaporizaition
liquid yo gas or gas to liwuid
397
tofigure out how much heat is required, use mc delta T to figure out heat to melting point,
then use q=mL for heat of fusion
| add both numbers together
398
entrophy of the universe is increasing
delta S universe= delta S system + delta S surroudsning >0
399
E (heat/Q)
P x T
400
500 gq
0.5 kg
401
density =
mass/volume
kg/m^3 or g/ml or g/cm^3
402
density of water
1 g/cm^3=1000 kg/m^3
403
weight of substance with density=
Fg= density x Volume x 9.8
404
specific gravity
SG= density/ 1 g/cm^3
405
pressure
Pressure= Force/Area
Pa (pascal)= 1 N/m^2
406
scalar
magnitude but no direction
| like pressure
407
absolute hydrostatic pressure
total pressure that is exerted on an object the is submerged in a fluid
P= P0 (pressure of the surface/ambient pressure) + density x 10 x depth of the object (z)
408
guage pressure
difference between the absolute pressure inside the tire and the atmosphericpressure outside the tire
Pguage= P- Patm = (P0 (pressure of the surface/ambient pressure) + density x 10 x depth of the object (z))- Patm
P0 and Patm can most likely equal each other
409
absolute pressure
P= Patm + Pguage
1.013 x 10^5 Pa if 1 atm
410
pascals principle
pressure applied to an incompressible force will be distorted undiminished throughout the entire volume of the fluid
411
archimdes principle
buoyancy
always teh opposite direction of gravity
if higher than force of gravity, will float (less dense objects)-- can be equal to mg if floating
Fb= density of the fluid x volume of the fluid displace x 10 = density of the fluid x volume subermedg x 10
412
cohesive forces
give rise to surface tension because cohesion together fluid
| attractive force a molecule feels toward other molecules of the same liquid
413
adhesion
attractive force that a molecule of liquid feels toward the molecules of other substance
water and glass
414
viscosity
reistance of fluid
415
Poiseuielles Law
Q= pi r ^4 delta Pressure/ 8 n L
Q- flow rate
r-radius of the tube
n-viscosity
L- length o pipe
look at relationship mroe than equation
416
contuinty equation of fluids
fluids flow more quickly in narrow spaces than larg enes
417
Bernoullis equation
pressure + 1/2 density (volume)^2= pressure + 1/2 density volume ^2 + density(g)(h)
418
couloumb
1.60 x 10^-19 C in 1 electron
negative in electron, positive in proton (with much greater mass)
419
insultors
do not distbrute charge over a surface
420
couloumbs law
Fe= kq1q2/r^2
electrostatic force= Fe
k- 8.99 x 10^9
q- charge (coulomb of electron)
421
gravitational force
Gm1m2/r^2
422
magnitude of electric field (E)=
Fe/q= kQ/r^2
423
field lines
represent how a positive charge would move in teh presence of a positive or negative charge
424
electric field is vector or scalar
vector- has direction
425
electric potential energy
= kQq/r
if both charges are positive or both negative, then potential will be positive
if both charges are different from each other, then both charges will be negative
426
electic potential
measured in volts= kQ/r
different points in space of an electric field surronding a source charge will have different electric potential values
the charge of electric potential that accompanies the mvoement of a test charge from one position to another
move in the direction that shows. decrease in electric PE (from high to low potential for positive test charges, and low to high for negative test charges)
427
equipotential line
every point is the same potential , so 0 is the difference
designate teh set of points around a source charge or multiple source charge that have the same electric potential
work is done when one charge is moved from 1 equipotential line to anther (but no work if move to same line)
428
two charges of opposite sign separated by d generates an
electric dipole
429
dimagentic
no unpaired electrons
| not MAGNETIC
430
paramagentic
unpaired electrons
| MAGNETIF
431
dipole potential
V= kqd/r^2 x cos theta
q- dipol moment
432
magentic force
```
Fb= qvBsin theta
q- charge
v- velocity
B- magenta field magnitude
theta- smallest angle between velocity and melanic field vector
```
433
sn1
1. loss of LG
forms carbocayion
+1 charge for carbocation functions as nucleophile
-1 LG charge acts as electrophile
RDS- rate determining step. increase by 2, rate increases by 2. first order
``
2. nucleophilic attack
434
SN2
there's a dipole, alkyl halide is a bit more negative (nucleophile)and C that is it attached to is more positive
the thing it reacts with (oh for ex) is the electrophile (the oxygen)
negative oxygen in OH attacks Carbon positive, at the same time, 2 electrons from the bond come off of kick bromine/alkyl group off
no carbocation
OH adn carbon that was attached to alkyl halide forms a bond (to form an alcohol)
435
nucloephile
onating an electron pair. ... Because nucleophiles donate electrons, they are Lewis bases.
436
electrophile
lectrophiles react by accepting an electron pair in order to form a bond to a nucleophile including the interactions of a proton and a base.
437
amide
Amides are functional groups in which a carbonyl carbon atom is linked by a single bond to a nitrogen atom and either a hydrogen or a carbon atom.
n secondary structure of proteins
438
amine
basic nitrogen atom with a lone pair.
| H-N-H
439
hydride shift
arbocation rearrangement in which a hydrogen atom in a carbocation migrates to the carbon atom bearing the formal charge of +1 (carbon 2) from an adjacent carbon (carbon 1).
440
bronsted acid
an acid is a proton (H⁺) donor, a
441
bronsted base
base is a proton accepto
442
intensity of radiation and photons emitted
, intensity is inversely proportional to wavelength.
| higher intensity more photons emitted
443
solvation
electrostatic interaction between solute adn solvent molecules (aka dissolution)
when water is the solvent, can be called hydration
involves breaking intermolecular interactions between mc and forming solute and solvent interactions together (if new interaction is stronger than separately-exothermic, favored at low temp)-- gas into liquid
444
when the interaction between solute and solvent is weaker than if they were seperate
then it is endothermic at high temperatures
| NACL and water
445
solubility
the maximum amount of the substance that can be dissolved in a particular solvent at a given temperature
soluable in sponetous reactions (neg delat G)
446
all salts of Group 1 metals and all nitrate salts are soluable
are soluale in water
447
coordinate covalent bonds
Lewis base (electron donating pair) and an electron acceptor (lewis acid0 form very stable Lewis base acid complexes
448
the formation of complex ions (Lewis acid and lewis bases to form one end product)
what that says yes
449
concentration
amount of solute dissolved in a solvent
450
mole fraction
moles of A/total moles of species
sum of all mole factions will equal one
451
molarity
moles of solute/liters of solution
452
molality
moles of solute/kg of solvent
453
normality=
molar mass x n (which could be number of protons or number of OH ions)
454
m1v1=
m2v2
Molarity
Volume
455
for a dissociation reaction
Ksp is the equilibrium constant (products /reactants of AQ only)
456
ion product and Ksp determines level of saturation
if less than Ksp- unsatured and more solute is added it will dissolve
if = to ksp - saturated and no change in concentrations
if more than Ksp then solution is supersaturated and precipitate will form
457
ksp vs. kf
ksp wil show dissoltion of OG solution
formation of complex ion is Kf (usually larger than Ksp)
multiply Kf and Ksp together when you combine both equations together
458
to find molar solubility
set up ICE table which Is baed on concentrationa
for C- look at coefficient
459
formation of complex ion increases the solubility of other salts with teh same ions because it uses up the products of those dissolution reactions, shifting equiilibrium tot eh R
look at Ksp adn Kf both equations
460
common ion effect
decrease the solubility of a compound in a solution that already has one of the ions in teh compound . shifts dissolution to eh L, decreasing dissociation
461
arehnnius acid
dissociate to form an excess of H+ in solution
462
arrhenius base
disociate to form an excess of OH-
463
amphortic speies
acts as an acid in basic envroent and a base in acidic environments
gain or loss a proton in bronzed
water for example
464
in pure wter
equal concentrations of H30+ and OH- if solution is equilibrium
465
pH=
-log H+= log 1/H+
466
pOH=
-log OH-= log 1/OH-
467
pH + pOH=
14
468
pH= -log (10^/3)=
3
469
ka and kb
Ka = [H3O+][A-] / [HA]. General Kb expressions take the form Kb = [BH+][OH-] / [B].
470
strong acid and bases completley dissociate
ph=14-pOH
| ph= 14-(-log[OH])= 14 + log (cocnentration fo OH)=
471
log 1=
0
472
log 10=1
YES
473
Kw=
[H30+][OH-]
474
strong acids
HCl, HBr. HI H2SO4, HNO3, HCLO4
475
stong bases
NaOH, KOH,
476
there can be negative PH and above 14 PH values
yes
477
weak acids and bases
dissociate partially
| Ka= (products)/reactas
478
smaller Ka means
weaker the acid and less it dissociates
479
small kb mean
weaker the base and it will dissociate less
480
conjugate acid
when a base gets a proto
481
conjugate base
when an acid loses a proton
482
autionization of water the Kw is
[H30+][OH-]=10^-14
483
Ka +Kb= Kw=
10^-14
484
Kb adn Ka are inversely related so
if Ka is large, Kb is small
| strong acid will produce a weak conjugate base
485
10 ^1
10
486
10^0
1
487
neutralization reaction
when acid and bases form a salt
488
when salt ions react with water to give back an acid or base
hydrolysis
489
titrations
used to determine the concentration of a known reactna tin a solution
titrant has a knoewn concentration and is added slowly to the tetrad to reach teh equvilance point
490
acid base titrations
equivalencies point is reached when the number of acid equivalents present in the OG solution equals th enumbe of based equvlant added
491
strong acid strong base will have an equvilance point of
7
492
indicators
weak acids or bases that have different colors in their protonated an deprotonated states
when it dodes change color= end point or rlly equivlance point
493
weak acid and sting base titrations will have equilateral point of
ph>7
494
weak base and strong acid will have equilibrium point of
pH<7
495
indicatorsshould have a pKA close to the PH of the
expected equvialnce point
496
bufering solitions
mixture of weak acid and its conjugate salt (conjugate base and cation) or weak base and its conjugate salt (conjugate acid and anion)
497
can have half equvilance points
half of the species has been prootonted in polyvalent acids (H2Co3) and bases
498
buffer soltions can resist changes in PH when small amounts of acid or base are added
yes
499
bicarb buffe system
Co2 + H20--> H2C03--> H+ + HCO3-
| cellular respiration maintain pH of blood
500
henderson hasslbach equation
pH= Pka + log [A-]/[HA]
A-= conjugate vase
HA- weak acid
501
weak buffer solution:n
POH= Pkb + log ([B+][BOH]
```
B+= conjugate acid
BOH= concentation of weak base
```
pOH=PKb when conjugate acid = weak base
502
first pka in a graph
halfway between starting point and first equivalence point
503
eqvulance point
steep slopes upwards
504
N1V1=
N2V2
normaly and volume
505
reactions that involve the transfer of electrons from one chemical species to another
oxidation reduction redox reactions
506
oxidation
loss of electrons
507
reduction
gain of electrons
508
oxidizing agent
causes another atom to undergo oxidation by being reduced itself
509
reducing agent
causes the other atom to be reduced and is oxidized itself
510
oxidation numbers
assigned to atoms in order to keep track of the redistribution of electrons during chemical reactions
N2, He- free elements- oxidation # is 0
monatomic ions= charge of ion: Na+, Cu2+, Fe3+ (+1, +2, +3)
511
group 1A elements oxidation numer
+1
512
group 2 oxidation number
+2
513
oxdaition number of oxugen
-2
514
if neutal atom
no charge above it, must be 0 oxidation number
515
if oxidatio number goes from +2--> +4
loses electrons, oxidation (making it a reducing agent)
516
If oxidation number goes from +4-+2
reduced, gains electrons (oxidizing agent)
517
optical isomers
stereoisomers- compounds that have the same chemical formula but differ from one another in terms of the spatial arrangement of their atoms.
518
enantiomers
nonidentical, non superimposable mirror images
chiral compounds wit no symmetry
every carbon Is different between D and L
519
chiral comoound
4 different groups attached to it
520
D and L / R and S
R=D- OH on R side
S=L- OH on left side
for last OH group
521
two sugars that are in the same family (ketoses or alcoves or have same number of C) but are not identical or no mirror images are
diastereomers
522
epimers
have one chiral carbon in a diastereoisomer
523
alpha anomeric carbons
point down OH's
| less favored because axial adding to steric strai
524
beta anomeric C
pointed up OH's
525
monosacarrides contain
OH groups and carbonyl C (R-C-R with OH on top of C)
526
aldoses can be oxidized and so are considered
reducing agents
527
1 chiral carbon has how many stereoisomers
2
2^chiral C= stereoisomers
528
fisher projections
when drawing OH groups think downright up - left
on the R side go down, L side go up
529
aldehyde or ketone (KETO FORM) + acid or base you get
enol
double bond with alcohol, from shifting alpha H
has a carbocation
enola nd keto are tatuermers
530
glucose is an
aldohexose
2^4 (chiral groups)= 16 stereocenters
531
2 monosaccarides linked together
form disaccarides by glycosidic linkages
532
glycolsidic linkages
C1 and usually C4 (from the anomeric C- the one on the lowest R side)
533
if OR group is CIS with respect to the C6 C (the one with CH2OH) then it is
beta
534
if OR group is trans with respect to the C6 C (the one with CH2OH) then it is
alpha
535
electroagetic spectrum
radio waves (long wavelength, low frequency) and game rays (short wavelengths, high frequency, high energy)
between them from lowest to highest we have: microwaves, inferred, visible light, UV and x rays
536
electromagnetic waves are
transverse waves because oscillating vectors are perpendicular to each other
537
mm, um, nm, angstrom
10^-3, 10^-6, 10^-9, 10^-10
538
sped of light =
frequency x wavelength
539
in visible light
red is 700 nm, violet is 400 nm
540
reflection
rebounding of incident light waves at the boundary a medium
541
angle of incidence =
angle of reflection
| as measured from the normal, which is a line drawn between each angle and straight up
542
real image
if light converges at the position fo the image
543
virtual
only I light appears to be coming from the position of the image but does not actually converge there
544
center of curbature
center of the spherical shaped mirror if it were a complete sphere
545
concave
inside of the sphere to its surface
in front of mirror ris center of curvature and radius
converging mirrors aka
546
convex
outside of the sphere to inside
| diverging mirrors
547
focal length
distance between the focal point and the mirror
f=r/2 (radius of curvature is the distance from C and the mirror)
distance between object and mirror is o
image and mirrors distance- i
548
focal length equation
1/f=1/o + 1/9= 2/r
549
if an image has a positive distance (I>0)
it is a real image, implying image is in front of the mirro r
550
if image has a negative distance (I<0)
it is virtual and located behind the mirror
551
magnification
= -I/o
ratio of size of image to the size of an object
552
- magnification
= inverted image
553
+ magnification
= upright image
554
if the object is placed at the focal point
no image is formed because reflected light rays are parallel t each other
555
if the object is between F and the mirror
the image is vital, upright and magnified
556
If the object is beyond F
image is real, inverted and magnitifed
557
divergin mirro
forms only virtual, upright adn reduced image
farther away object, smalle rimage will ve
558
object is in front of mirror
o is positive (and real)
559
o is negative
objet is behind mirror (virtual)
560
I is positive
image is in front of mirror (real)
561
I is negative
image is behind mirror (virtual)
562
r or f is positive
mirror is concave (converging)
563
r or F is negative
mirror is convex (diverging)
564
m is positive
image is upright
565
m is negative
image is inverted
566
refraction
bending of light as it passes from one medium to another and it changes speed
567
index of refraction equation
n= c/v
| index ofrefraction = speed o light in a vacuums, speed of light in medium
568
snell s law
refracted light
n1sin=n2sin
1- where light is coming from (when this is bigger, bend away from normal)
2- where light is entering (when this is bgiger, bend to normal)
569
when light travels from a medium of high index of refraction (such as water) to a medium with lower index of refraction (such as air)
the refracted angle is larger than the incident angle (theta 2> theta 1), bends away from normal
570
as the incident angle increases, refracted angle also increases
making the critical angle, refracted is equal to 90 degrees (theta 2)
refracted light passes along the interface between teh two media
critical angle = sin ^-1 (n2/n1)
571
lenses refract light, mirrors
reflect it
572
lenses
there's two surfaces (the glass when light goes in adn out) so equation is:
1/f= (n-1)(1/r1 - 1/r2)
n-index of refraction
r-radius of curvature
573
the cornea acts as the
primary source of refractive power
574
lens:
```
o positive
I positive
r positive
f postiive
m positive
```
object is on same side of lens as light source
image is on opposite side of lens from light source (real)
lens is convex (converging)-- both r and f positive
image is upright
575
```
lens:
negative
o
I
r
f
m
```
object is on opposite side of lens from light sourec
image is on same side of elns as light source (virtual)
lens is concave (diverging) - for r and f
image is inverted
576
POwer and focal length
P= 1/F
577
power is positive for
converging lens
578
power is negative for
diverging lens
579
people who are near sighted (can see near objects- myopia)
need diverging lens
580
peopel who are far sighted need (hypeopia)
converging lens
581
or a thermodynamic mixture of isomeric products, the relative mole ratio of products is directly related to the
relative stability of these products.
582
mF to F (for captors)
1000 mF is = to 1 F
583
q=CV
```
q= charge stored
C= catpiance
V= votlage
```
584
energy with captiance =
1/2 CV^2
585
doppler effect
f prime > f before passing the horn, f prime< f after passing it
586
P = Fv
power= force x velocity
587
work is
= to KE
588
kW to W
1 Kw= 1000 W
589
slowing down/ stopping or anything
v't graph is a parabola
590
hermodynamic stability of isomers can be determined based on the amount of heat produced when the compounds are combusted;
less heat, greater stability.
tability is related to energy in that more stable molecules have less energy while the less stable molecules have more energy to start with. So the ones that give more heat off when burned were less stable.
591
Boiling point is about the intermolecular forces, while heat of combustion is about the intramolecular forces.
Thermodynamic stability is about intramolecular forces
592
example problem with labelling
The researchers want to know where guanine binds/is incorporated in reactions. To do this, they decided to use radioactive hydrogen which can be traced.
If they want to put the hydrogen in guanine, they have to place it somewhere where it won’t be likely to fall off. Position 2, 3, and 4 all have nitrogens and hydrogens. If the researchers attach the radioactive hydrogen to any of these spots, they might risk the nitrogen’s reacting with the water(solvent) and lose the hydrogen. If the hydrogen is lost, the researchers will have no idea where the guanine attached.
If they put the hydrogen at position 1, it is very unlikely that the hydrogen will fall off b/c the double bond is fairly unreactive towards the solvent.
593
activation energy for a reaction represents the minimum energy barrier necessary to be overcome by the reactants on the path to products.
.the activated complex minus the energy of the reactants.
594
period
from one place to same next place on the wave
top to top
time it takes to complete one full cycle
595
The glucose meter measures the current produced during Reaction 2. If 0.67 μmol of electrons were measured, what mass of glucose was present in the sample? (Note: The molar mass of glucose is 180 g/mol = 180 μg/μmol.)
stoichiometry of the reaction is 2 mol e– per mole of glucose consumed. The device measured 0.67 (2/3) μmol of electrons, indicating that 0.33 (1/3) μmol of glucose was consumed. This weighs 60 μg, based on its molar mass of 180 g/mol.
596
Alter the ratio of monosodium/disodium phosphate added to favor the monosodium species.
will increase acid production because monosodium has extra H
597
the formation of a peptide bond is accompanied by the formation of water as a by-product, and the mass of water is 18 amu.
byproduct of peptide bond formation
598
acetylation (attachment of RC=O)
to N
599
n atom must be a Lewis base to coordinate to calcium ions, and oxygen is the only Lewis basic atom present in the side chains or backbones of the listed amino acids that has either a partially negative charge (in the peptide backbone or Ser side chain) or a negative charge (in the Asp side chain).
A coordinate covalent bond is a bond where both the electrons in it come from the same atom.
A Lewis base is an electron donor.
So yes, the Lewis base will be responsible for
600
pyrrole ring
five-sided heterocycle containing one nitrogen atom
601
resistivity is the inverse of the conductivity,
yes
602
involves transfer of a methyl group to cytosine, and the next step involves the hydroxylation of that methyl group.
refore, the two classes of enzymes needed are a transferase and an oxidoreductase.
603
hydroylxation
a chemical process that introduces a hydroxyl group (-OH) into an organic compound
oxydation reduction
oxidation reaction in which carbon–hydrogen (C-H) bond oxidizes into carbon–hydroxyl (C-OH) bond. I
604
deta G equation
ΔG′° = −RTlnKeq = −RTln([unfolded]/[native]).
605
Because DNA unfolding occurs as the pH increases, it
can be inferred that this is due to cytosine deprotonation.
606
glucose and dextran (sugar)
polysaccharide of glucose has numerous hydroxyl groups that can hydrogen bond to the polar side chains that are typically exposed on a protein surface.
607
the frequency of a wave is not affected by the medium through which it propagates
but amplitude, wave speed and wavelength are
608
sponetous reactiona and delta K
because a spontaneous reaction is one that exhibits ΔG < 0. Since ΔG = –RTln(Keq), this means that Keq must be > 1
609
mass is
density x volume
610
electrohemical cels
contained systems in which oxidation reduction reactions occur
galvanic cells (voltaic cells), electrolytic cells, concenrtation cells
611
galvanic cells
sponeteous reactions
612
electrolyytic cells
nonsponteous reactions
613
oxidation occurs at
anode
614
reduction occurs at
cthode
615
electromotive force
corresponds to teh voltage or electrical potential difference f the cell
if positive, cell releases energy (delta G less than 0)
if negative, aborb energy (delta G positive and nonspop)
616
movement og electrons in electrochemical cells is from
anode to cathode, and current runs from cathode to anode
617
galvanic cells
release E into teh environment
2 electrodes connected to each other through a conductive metal, like copper and the solutions contain a salt bridge
sponetous redox reaction - as the reaction proceeds toward equilibrium, the movement of electrons results in conversion of electrical potential to KE and this can hel pdo work
618
salt bridge
charge disperses so that when the reaction runs, not too much negative or too much positive is on one side
619
cell diagarm
short hand notation for representing reaction
Zn (s) [ Zn 2+ [ Cu2+ [ Cu
anode, anode solution cathode solution, cathode
620
electrolytic cells require the
input of E to proceed, driven by external voltage source called electrolysis
621
Cl- cells migate to the
anode
| where they are oxidized to Cl2
622
Na+ ions migrate to teh
cathode
| where they are reduce to NA
623
1 electron carries a charge of
1.6 x 10^ -19 C
624
Farday constant
equvilant to eh charge contained in 1 mol of electron, as 96,485 C= 1 F
10^5
625
concentration cells
specialized form of galvanic cells in which electrodes are made up fo the same material , concentration gradient causes movement of charge, not potential difference
626
charge of electrode dependent on type of electrochemical cell
galvanic cells- anode is neg and cathoe is positive
| electrolytic cells- anode I positive and cathode is neg
627
electrons typically move from negative (low electric PE) to
postiive (positive electrical PC)
628
isoelectric focusing
technique used to seperate amino acids or polypeptides based on isoelectric points
postiive charged aa will igrate to cathode (protonated at soltiosn pH) and nehatively charged aa will migrate to anode *deportonated )
629
reduction potential
gien in Volts
oxidized or reduced which psecies can be determined (which gains e to be reduced)
more positive potential, more likely to be reduced
630
electrode witht the more positive reduction potential
cathode (opposite for electrolytic cells)
631
lectrodue with less positive reduction potential
anode (opposite for electrolytic cells)
632
beause species wants to gain electrons, reaction I sponteous
so negative delta G
633
since reduction adn oxidation are opposite processes, to brain teh oxidation potential,
switch sign of reduction potential
634
electrolytic potential of cell =
E reductionc cathode - E reduction anode
635
for galvanic cells
difference inr eduction potentials must be pos
636
or electrolyti cells
difference in reduction potentials I negative
637
electromagnitude F (E) and change in free E have opposite signs (G)
if E is pos, G is neg and galvanic
E is neg, G is pos= electrolytic
E and G are 0= concentration cells
638
Nerst equaton
describes the relationship between the concentration fo species in a solution under nonstandard conditions adn electromotive F
E cell = E^o cell --RT/nF. ln Q
e cell- no standard condition, E^o cell standard conditions, R - ideal gas constant , temp, number of moles of electrons, Farday, Q is reaction quoeitn
Q- reactants/products of aq
639
delta G =
-nFE cell
n- number of moles of eelctron exhcnaged
F- Farday constant
E- emf of cell
640
Work= delta q x V
amount of work is need to transport charge across potential difference (voltage)
641
when kEq (ratio of products concentrations at equilibrium over reactant raised to stoic coefficient) isgreater than 1
E cell is positive
642
when Keq is less than 1
E cell is negative
643
When keq=1
E cell is 0
644
gluocose transporters
GLut 1-4
| Glut 2 and 4 are important because in specific cells and highly regulated
645
GLUT 2
low affinity transporter in hepatocytes and pancreatic cells
after a meal, blood traveling from a hepatic portal vein from teh intestine is rich in glucose. glut 2 captures excess glucose for storage.
beta ce;;s and also glut 2 help with insulin release
646
glut 4
adipose tissue adn responds to glucose concentration in peripheral blood
increased by insulin (rate of glucose transport), stimulating the movement of additional glut 4 transporters to membrane
647
glycolysis
pathway that converts glucose into 2 PYVRUATE molecules , realign energy captured in two substate level phosphorlyations adn one oxidation reaction
needs mitochondria and oxygen - aerobic through NADH
if lacking- RBC , glycolysis occurs anaerobically bu energy is lost
648
hexokinase and glucokinase
the first steps in glucose metabolism in any cell are transport across the membrane and phosphorylation by kinase enzymes inside the cell to precent glucose from leaving via the transporter
convert glucose to glucose 6 phosphate
649
PFK 1
rate limiting enzyme and main control point in glycolysis
fructose 6 phpshatr is phosporlyated to frustose 1,6 bisphoat using ATP
inhibited when made enough ATP and citrate (meaning Krebs cycle is working)
activated by AMP and fructose 2,6 bisphosphate
650
PFK 2
produces fructose 2,6 bisphosphate that activates PK1 . activated by insulin and inhibited by glucagon
651
glyceraldehyde 3 phosphate dehydrogenase
produces NADH which can feed into teh electron transport chain
652
3 posphoglucerate kinase and pyruvate kinase
put phsphate on ADP to make ATP through substrate level phosphorlation
653
irreversible reactions enzymes
glucokinase/hexokinase, PFK 1 and pyruvate kinase
654
NADH produced in glycolysis is
oxidized by the ETC in mito when o2 is present
655
if o2 and mito are absent, NADH in glucolysis is
oxidized by lactate dehydrogenase
656
lgycolysis overview
glucose + 2 NaD+ + 2 ADP + 2 Pi= 2 pyruvate + 2NADH + 2 ATP + 2 H20 + 2 H+
(use 2 ATP because use 2 in process)
657
pyrutes are
3C molecules
658
glycolysis
1. phosphorlyation of glucose by hekokinase to glucose 6 phosphate (uses ATP)
glucose 6 phosphate to fructose 6 phosphate by phosphoglucose isomerase
fructose 6 phosphate ohosprulated by PFK 2 to make fructose 2 6 bisphosphate
fructose 2 6 phosphate is then catalyzed via PFK 1 to make fructose 1 6 bis phosphate which splits into DHAP and glyceride 3-P after aldolase
Glyceradlude 3P is mixed with glyceradhyde 3p dehydrogenase to make 1 3 bisphophosoglycertae which combined with phosphoglycerate kinase to make 3 phosphoglycerate
last step is PEP phosphorylated by pyruvate kinase to make pyruvate
659
lactate to pyruvate by lactate dehydrogenase
fermentation
oxidizes NADH to NAD+
generating more NAD+ fo rglycolysis to contain
660
cellular respiration
glycolysis (in cytosol) which splits up glucose into 3 C pyruvate to make 2 ATP and NADH.
pyruvate fermentation, where pyruvate is converted to acetyl COA for entry into teh citric acid cycle if ATP is needed= carboxyl group dt4ripped off of pyrvuate adn exerted as CO2 . acetyl attaches to COA (where pyruvate dehydrogenase complex is- inhibited by acetyl COA in inhibition). Pyruvate can also be converted to lactate by lactate dehydrogenase or oxaloacetate by pyruvate carboxylase.
acetyl COA can then be used in teh citric acid cycle to make CO2 and H20 or FA synthesis to make FA
661
citric acid cycle
acetyl COA + oxaloacetic acid (4C) to make citrate/citric aci (6C). oxidized over a bunch of steps to form oxaloacetatic acid again. Cabers that are cleaved off make Co2 and leave. Makes 3 molecules of Co2 each spin, twice usually - so 6 CO2
reduced 10 NAD to 10 NADH
and 2 FAD to 2 FADH2 (oxidized)- inputs to ETC
2 ATP
662
ETC
produces 32 ATP
actually generate ATP
each NADH produces 3 ARP
FADH2 each produces 2 ATP
so 10 NADH come in (30 ATP) total, FADH2 has 2 go in so (4ish) ATP
NADH is oxidized and electrons go down complexes and losing energy
pump hydrogens into the outer membrane to move electrons down and ATP H+ synthase
663
glycogen synthesis and degradation occurs in teh
liver adn skeletal muscle
| mobile between meals to prevent low blood sugar
664
glycogenesis
synthesis of glycogen with UDP
665
gluconeogeneisis`
synthesis fo glucose from noncarbo precursors
reverse og glycolysis
666
boiling chips
The air bubbles break the surface tension of the liquid being heated and prevent superheating and bumping
667
The boiling point of a liquid is the temperature at which the vapor pressure of the liquid equals the surface pressure.
which is usually 1atm
668
color and orbitals
The color arises because nickel(II) ion has partially filled d orbitals and the electrons in the lower energy d orbitals absorb visible light to move to the higher energy d orbitals.
669
inductive effect
The effect of the sigma electron displacement towards the more electronegative atom by which one end becomes positively charged and the other end negatively charged
670
electric field
Volts/Meter= N/C
voltage / distance
671
If the first harmonic has a frequency of 100 Hz, then the second harmonic
is 200 and the third is 300, etc.
672
(NH3 + H+ → NH4+) does the oxidation state of N change?
No! tell me why??
673
One mole of gas occupies 22.4 L at STP.
yes
674
you cannot turn a tertiary alc to a
ketone
675
Z
same side
676
E
opposite sides
677
Two vectors of magnitudes |A| = 8 units and |B| = 5 units make an angle that can vary from 0° to 180°. The magnitude of the resultant vector A + B CANNOT have the value of:
add them together and then subtract them together- this is the range so cannot have 2!
678
to correct nearsightedness, the laser beam is directed onto the central part of the cornea, resulting in a flattening of the cornea.
This means that the radius of curvature of the cornea is increased.
679
time internval also measn
period
680
pulsed vs. continuous laseer
The pulsed laser radiation interacts with the cornea for a shorter time than a continuous laser radiation, thus less heat is transferred to the cornea. This allows the cornea to maintain a lower average temperature by cooling off between pulses
681
if you are given Kb - set up an equilibriu m
of concentrstions!
682
square root of 1.2 x 10^-4 is 1 x 10^-2
yes
683
A carbonyl group contains a C=O double bond. .
The first bonding interaction between atoms is always a σ bond. The second bond is formed from π symmetry orbitals
684
cyclic ester
lacton
685
ester
R-C(=O)-O-R
686
The substrate is covalently attached to:
they are asking about how the substrate, glycerol 3-phosphate, covalently attaches to the active site of enzyme G3PP. In the hydrolysis reaction, we know that the active site Asp that will covalently attach to the substrate is the one that's a nucleophile because the nucleophile is the one that is capable of bond formation, not a general acid. We are told that Asp14 is the nucleophile, so we can immediately eliminate options C and D. Next, they are asking what atom on the substrate the Asp14 attacks. We know that Asp has a carboxyl group in the side chain, which means the oxygen on the Asp is the one that attacks. We know the nucleophile in the reaction. The oxygen (from Asp14) is not going to attack another oxygen, it's going to attack a phosphorus, and hence we can deduce that the electrophilic atom on the substrate is the phosphorus, not an oxygen. Thus, the answer is A. Hope that helps!
687
v max and k cat are properotional
both increase or decrease
688
32P undergoes the β– decay according to the reaction 32/15P →
32/16S + e– + energy.
689
GC has
3 hydrogen bonds
690
One company sells a defibrillator for home use that uses a 9-volt DC battery. The battery is rated at 4.2 A•hr (amp•hour). Roughly how much charge can the battery deliver?
Current has units of Amperes (A) which is charge per unit time, or C/s. The question is asking for how much CHARGE is delivered, meaning we need the answer to be in Coulombs. They gave us 4.2 amp •hour, which has the units 4.2 C/s • hour. In order to convert that to just coulombs, we convert hour to seconds (1 hour = 3600 seconds), and if we multiply
4.2 C/s x 3600s, this leaves us with 15,120 C
691
v/t =
acceleration
692
In a neutral atom, the number of electrons is equal to the number of protons.
yes!!!
693
ester
made from acid and alcohol
694
protectign group
no reaction from molecule
