General Chemistry Flashcards
(120 cards)
1
Q
protons
A
found in nucleus of an atom, has a fundamental unit of charge e (1.6x10^-19C) often denoted as +1
2
Q
Mass of a proton
A
1 atomic mass unit
3
Q
atomic number
A
number of protons found in an atom of that element, all elemends are defined by number of protons they contain, but they do not all necessarily have the same mass
4
Q
neutrons
A
found in nucleus of atom and has no charge
5
Q
mass number
A
sum of protons and neutrons in an antom’s nucleus, thus not always the same as atomic number
6
Q
isotopes
A
atoms that share an atomic number but different mass numbers
7
Q
electron
A
move in space surrounding nucleus of an atom associated with varing levels of energy, each electron has a charge equal in magnitude to that of a proton but with a negative sign (-e), mass of these is negligable
8
Q
electron shells
A
levels of distance from the nucleus that correspond to varying levels of electrical potential energy, higher shells are higher energy thus those farthest from the nucleus have strongest interactions with the surrounding environment and weakest interactions with the nucleus
9
Q
valence electrons
A
electrons of a high shell most likely to form bonds with other atoms because of the little electrostatic pull they experience from their own nucleus
10
Q
cation vs anion
A
positively charged atom vs negatively charged
11
Q
atomic mass vs mass number
A
nearly equal
12
Q
atomic weight
A
takes the weighted average of different isotopes to report a number on the periodic table in amu, also represents the mass of one mole of the element in grams
13
Q
planck relation
A
E=hf (h = planck constant = 6.626x10^-34 Jxs) (f=frequency of radiation)
14
Q
quanta
A
idea that energy emitted as electromagnetic radiation from matter comes in descrete bundles represented by this
15
Q
Bohr’s angular momentum of an elecron
A
L=(nh)/2pi (n is principal quantum number, h is planck constant)
16
Q
energy of an electron equation
A
E = -(RH/n^2) (RH is rydberg unit of energy equal to 2.18x10^-18 J/electron)
17
Q
The energy of the electron changes in descrete amounts with respect to the quantum number. Therefore, the electron in any of its quantized states in the atom will…
A
…have an attractive force toward the proton, the energy of an electron increases by becoming less negative the further out from the nucleus it is located
18
Q
ground state vs excited state
A
ground state is lowest energy where all electrons are in lowest possible orbitals, excited is when at least 1 electron has moved to a subshell of higher energy
19
Q
line spectrum
A
a concept in atomic emission spectra where each line corresponds to a specific electron transition, each eleement has its own unique one that is like a fingerprint for the element
20
Q
orbitals
A
regions of space around the nucleus that electrons move rapidly and are localized within based on probability of finding an electron in a given region of space
21
Q
heisenberg uncertainty principle
A
it is impossible to siultaneously determine with perfect accuracy the momentum and position of an electron, if we want to assess the position the electron has to stop removing its momentum and if we want to assess momentum the elctron has to be moving
22
Q
quantum numbers and the pauli exclusion principle
A
modern atmoic theory that postulates an electron in an atom can be completely described by 4 quantum numbers, and the pauli exclusion principle means no two electrons in a given atom can posses the same set of 4 quantum numbers. (spin up and spin down)
23
Q
Principal quantum number
A
the first quantum number denoted by n, the larger the value the higher the energy level and radius of the electron’s shell
24
Q
azimuthal quantum number
A
2nd quantum number denoted by l, refers to the number and shape of subshells in a given principle energy level shell, the range of l is 0 to n-1
25
magnetic quantum number
designated ml, specifieds the particular orbital within a subshell an electron is most likely to be found at a given moment in time. Each orbital can hold max of 2 electrons, possible values are integers between -l and +l, including 0.
26
spin quantum number
4th quantum number denoted by ms, and denoted +1/2 and -1/2, whenever 2 electrons are in the same oribtal they must have opposite spins and are referred to as paired, electrons in different orbitals with same ms values are parallel spins
27
electron confiuration subshell flow diagram
1s
2s
3s 2p
4s 3p
5s 4p 3d
6s 5p 4d
7s 6p 5d 4f
7p 6d 5f
28
hund's rule
idea that orbitals will fill so maximum number of half filled oribtals with parallel spins will fill first before doubling up with another electron
29
s subshellholds how many electrons? and p? and d? and f?
2, 6, 10, 14
30
exceptions to hund's rule
while moving an electron to a higher level is energetically unfavorable, the extra stability from filling up a subshell outweighs that cost ***think chromium and copper***
31
presentce of paired or unpaired electrons affects the chemcial and magnetic properties of an atom or molecule. Materials composed of atoms with unpaired electrons with orient their spins in alignment to a magnetic field, thus weakly magnetic. These are considered ___, opposed to ___
paramagnetic, diagmagnetic (only paired electrons)
32
Representative elements or A elements
Elements that have their valence electrons in the orbitals of the s or p subshells
33
B elements include these two
transition and nonrepresentative elements, which have valence electrons in s and d subshells and the lanthanide and actinide series which have valence electrons in the s and f subshells
34
effective nuclear charge (zeff)
electrostatic attraction between valence shell electrons and the nucleus, for elements of the same period this increases left to right, and atomic radius decreases from left to right across a period
35
smallest atomic radius of the periodic table
helium
36
ionization energy
energy to remove an electron from a gaseous species, requires heat making it endothermic processt
37
ionization energy trend
more difficult to remove when the closer the valence elecrons are to the nucleus and the greater the atom's Zeff, left to right across a period and from bottom to top of a group
38
If losing certain number of electrons givesn an element a noble gas like configuration, then removing subsequent electrons will....
...cost much more energy, take Mg2+ for example which has the same electrons as Ne, the energy to get Mg3+ is massive
39
electron affinity
energy dissipated by a gaseous species when it gains an electron
40
electronegativity
measure of attractive force an atom will exert on electron in a chemical bond, related directly to ionization energy
41
Exception to electronegativity trend
the first 3 noble gasses despite high ionization energies have negligible electronegativity because they do not form bon
42
octet rule
atoms tend to bond with other atoms so there are 8 electrons in the outermost shell, a configuration similar to that of the noble gases
43
nonpolar vs polar covalent bonds
if the electrons are shared equally between the two atoms vs if unequal
44
5 charactereistics of ionic compounds
-high melting point
-high boiling point
-dissolve in water and other polar solvents readily
-conduct electricity
-form crystal lattice
45
pauling scale
measure of difference in electronegativity for electron transfer to occur for ion formation vs covalent bonds, >1.7 to ionize
46
dipole moment
occurs in polar covalently bonded molecules where there is a partial positive and partial negative side to the molecule
47
dipole moment equation
p=qd where p is dipole moment, q is magnitude of charge and d is displacement vector separating 2 partial charges
48
coordinate covalent bond
both of the shared electrons originated on teh same atom, generally meaning a lone pair of one atom attacked another atom with unhybridized p orbital to form a bond, typically seen in lewis acid-base reactions where a lewis acid will accept a pair of electrons and a lewis base will donate a pair to form a covalent bond
49
formal charge
to determine if a lewis structure is representative of the actual arrangement of atoms in a compound, one must calcuate this of each atom, found by the difference between number of electrons assigned to an atom in a lews structure and the number of electrons normally found in that atom's valence shell
50
formal charge equation
formal charge = V - N (nonbonding) - 1/2N(bonding) where v is normal number of electrons in atoms's valence shell, nonbonding is number of nonbonding electrons and n bonding is number of bound electrons (will be double because 2 electrons per bond)
51
resonance
when it is possible to draw 2 or ore lewis structures that demonstate the same arrangement of atoms but differ in specific placement of electrons, the actual distribution in the compound isa composite of all possible structures, formal charge can be used to assess stability of resonance structures
52
exceptions to the octeet rule
because electrons in elements in or beyond the third period can place them in the d subshell, they can form more than 4 bonds, don't discount lewis structures with central atoms with more than 4 bonds
53
valence shell electron pair repulsion theory (VSEPR)
use of lewis dot structures ot predict molecular geometry of covalently bonded molecules,
54
electronic geometry vs molecular geometry
electronic describes the spacial arrangement of all pairs of electrons around the central atom, including both the bonding and the lone pairs, in contrast, molecular geometry describes spatial arrangement of only bonding pairs of electrons, for example Nh3 and methane and h2o all have the same electronic geometry with 4 pairs of electrons around the central atom, but because of different numbers of lone pairs they have different molecular geometry
55
is any molecule with a polar bond automatically a polar molecule?
no, you have to think about the dipoles which can cancel out, such asi n the case of CCL4 which has no net dipole because the tetrahedral shape cancels out
56
sigma vs pi bonds
patterns of overlap observed in molecular bond formation, sigma allows for free roation about the axes because the electron density of the bonding orbital is linear, while a pi bond does not because the electron densities of the orbitals are parallel and cannot be twisted
57
london dispersion force
(a type of van der waals force), the intermolecular force the transient polarization that an electron cloud has from a short lived dipole moment where an electron because it is in orbit will briefly polarize the molecule and on a large scale can lead to close tistance adhesion, large molecules possess greater dispersion forces
58
dipole dipole interactions
intermolecular force where polar molecules orient where positive ends are near negative ends and form a slight electrostatic force, negligible in the gas phase because of increased distance
59
hydrogen bonds
intermolecular force unusually strong form of dipole dipole interacction that only occurs when hydrogen is bonded to fluorine, oxygen, or nitrogen, causing an unusually high boiling point
60
hydrogen bonds occur when bonded to the atoms
FON
61
molecular weight
sum of atomic weights of all atoms in a molecule, measured in amu/molecule
62
mole
quantitiy of any substance, equal to the number of particles found in 12 grams of carbon-12, defined by avogadro's number 6.022x10^23 mole^-1
63
molar mass
mass of 1 mole of a compound in grams/mol, the same value as molecular weight but the two have different units
64
equivalents
the idea that because some compounds are more potentant and can donate more of some atom to a reaction (the porotons in HCL vs H2SO4) therefore you would need 1/2 as much h2so4 to get equivalent reaction
65
normality
most commonly used for hydrogen ion concentration measured in equivalents/L,, assuming that the reaction will proceed to completion
66
percent compoosition
percent by mass of specific compound that is made up of a given element, determined by mass of element in formula/molar mass
67
combination vs decomposition reaction
combination has more reactants forming fewer products, decomposition is a single reactant into multiple products
68
limiting reagent
component of a reactant that is consumed first and once used up limits the amount of product that can be generated, ust be measured in moles, NOT grams
69
theoretical yield
maixmum product that can be generated as predicted from he balanced equaiton assuming all limiting reactant is consumed and no side reactions have occured and entire product has been collected
70
strong vs weak electrolyte
strong completely dissociates in water into its constituent ions, while weak do not completely dissolve, and some do not dissolve at all (nonelectrolytes)
71
Gibbs free energy
determines whether a reaction will occur by itself without outisde assistance or not,
72
intermediate
molecule that is in a step of a reaction that does not appear in the overall reaction, they are often difficult to detect because they are consumed almost immediately after being formed
73
rate determining step
slowest step in any probposed mechanism that acts like a kinetic bottleneck preventing the ovrall reaction from proceeding any faster than that step, the rate of the reaction is equal to this
74
activation energy
minimum energy of a collision (molecules colliding with each other in the correct orientation with sufficient energy to break their existing bonds and form new ones), also known as the energy barrier, only a fraction of colliding particles have enoguh kinetic energy to exceed activation energy meaning only a fraction of all collisions are effective
75
Arrhenius equation
measure of collision theory of chemical kinetis k=Ae^(-Ea/RT) where k is rate constant, A is frequency factor, Ea is activation energy of reaction, R is ideal gas constant, and T is temp in kelvin
76
If delta G is positive, a reaction is ____ and ____ energy, and vise versa if G is negative
endogonic, absorbs, exergonic, releases
77
factors affecting reaction rate (4)
reaction concentrations - the higher the concentration the greater number of collisions
temp - higher temp more collisions
medium - some are more likely to react in aqueous vs nonaqueous solvents
catalysts - substances that increase rxn rate without themselvs being consumed in the reaction, lowers energies of activation, but have no impact on equilibrium position
78
for all forward, irreversible reactions, rate is proportional to the concentration of the reactants, with the concentration raised to some experimentally determined exponent, for the reaction aA +bB -->cC + dD, what is the rate proportional to?
k (constant) x A^x x B^, the values for k x and y have to be determined experimentally for a given reaction at a given temp**** test will provide experimental data
79
zero order reaction
a reaction in which the product formation is independent to chagnes of concentrations of any of the reactants, they have a constant reaction rate equal to the rate constant (k) rate constant itself is dependent ontemp so the only wayt o change the rate is to chang ethe temp or add a catalyst that lowers the activation energy
80
first order reaction
a reaction that has a rtae directly proportional toonly one reactant, such that doubling the concentration of that reactant results in doubling of rate formation of the product,
81
second order reaction
a rate proportional to either the concentrations of two reactants or the square of the concentration of a single reactant,
82
dynamic vs static equilibrium
in dynamic the forward and reverse reactions are still occurring they ave not stopped, but because they are at the same rate there is no change in concentrations of products or reactants, static equilibrium is when reaction has stopped
83
at equilibirum the entropy is at ___ and gibbs free energy is at ____
maximum, minimum
84
law of mass action
for generic reversible reaction aA + bB <=> cC + dD, then Keq = C^c x D^d / A^a x B^b
85
when a rxn occurs in more than 1 step, the equilibrium constant for the overall rx iis found by....
....multiplyting together equilibrium constants for each step of the reaction
86
reaction quotient
Qc, used in same equilibrium law of mass action Keq equation but if
87
Le Chatlier's principle
if a rxn is temporarily moved out of equilibrium state by a stress, the reaction responds by reacting in whichever direction will result in reestablishment of equilibrium
88
how does pressure impact equilibrium
as pressure increases the system will move to whichever side ha the lower number of moles of gas
89
isolated vs closed vs open system
cannot exchange energy (heat and work) with environment, vs can exchange but these but not matter vs can exchange everything
90
Heat vs temp, heat considers the
amount
91
processes where system absorbs heat is ___ delta Q >0, and processes where system releases heat is ___ delta Q<0
endothermic, exothermic
92
enthalpy
total heat content of a particular system, equal to Q under constant pressure
93
derived formula for bomb calorimeter
q system = - q surroundings or q cold = -q hot assuming no phase change (and remember q=mcdeltaT
94
bond breakage is ___thermic, bond formation is ___
endo, exothermic
95
sublimation
transfer from solid togas
96
change in entropy equation
delta S = Q/T
97
gibbs free energy equation
delta G = delta H - T delta S
98
gases, unlike liquids, are ___, although not infinitely so
compressible
99
atmospheric pressure and how does a barometer work
a downward force of all the air above onto an object, in the case of a barometer this creates a downward fore on the pool of mecurary at the base of the barometer while the mercury in the column exerts an opposing force based on tis density, creating a vacuum in the top of the tube causing it to rise when external air exerts greater force and lower when external air exerts less fore
100
STP conditions
standard temp and pressure conditions of 273 K (0 degrees C) and 1 atm
101
standard state conditions
298 K, 1 atm, 1 M concentrations
102
ideal gas law formula
PV = nRT where P is presure, V is volume, n is number of moles R is ideal gas constant, T is temp
103
density of a gas formula
PV = nRT where n (mass/molar mass)
104
combined gas law equation
P1V1/T1 = P2V2/T2 (1 is state of gas at STP and 2 is actual temp and pressure
105
avogadro's principle
alll gasses at a constant temp and pressure occupy volumes directly proportional to number of moels present, equal amoutns of all gasses at the same tmep and pressure will occupy equal volumes
106
boyle's law
for a given gaseous sample held at constant temp, isothermal conditions, the volume of gas is inversely proprtional to its pressure
107
boyle's law equation
PV = k or P1V1=P2V2
108
charle's law
under constant pressure the volume of a gas is proportional to its absolute temp expressed in kelvins
109
charle's law equation
V/T=K or V1/T1=V2/T2
110
Gay lussac's law
same as charles law but relates pressure to temp instead
111
Gay lussac's law equation
P/T = k or P1/T1=P2/T2
112
dalton's law of partial pressures
idea that gases that do not chemically interact that are found in one vessel each gas will behave independently as though it were the only gas in the container,
113
dalton's law of partial pressures equations
PT = Pa+ Pb+ Pc +....
114
henry's law
at various applied pressures, the concentration of a gas in a liquid increased or decreased, a characteristic of a gas's vapor pressure
115
henry's law equation
A = KH x Pa or A1/P1 = A2/P2 = KH where A is concentraiton of A in solution, KH is henry's consant, PA is partial pressure of A
116
kinetic molecular theory and assumptions (5)
explains the behavior of gases based on certain assumptions below
-gases are made up of particles with volumes negligibile to container volume
-gas atoms or molecules exhibit no intermolecular attractions or repulsions
-gas particles are in continuous, random motion, undergoing collisions with other particles and contanier walls
-collisions between any 2 gas particles or the wall are elastic meaning there is conservation of both momentum and kinetic energy
-average kinetic energy of gas particles is proportional to the absolute temp in kelvins and is the same for all gases at a given temp
117
average kinetic energy of a gas particle equation
KE=1/2mv^2 = 3/2KbT where Kb is boltzmann constant 1.38x10^-23 J/K
118
root mean square speed equation
u = square root of 3RT/M where R is ideal gas constant, T is temp and M is molar mass
119
grahams law of diffusion and effusion
idea that during diffsuion heavier gases diffuse more slowly than lighter oens because of differing average speeds, because all gas particles have same kinetic energy at same temp, larger mass molecuels move at slower speed
120
graham's law equation
r1/r2 = square root of M2/M1 (rates of diffusion and molar masses, notice that a gas that ahs molar mass 4x that of another gas will travel half as fast as the lighter gas)
