MCAT Chemistry Flashcards

Terms Definitions
Ionization Constant
Heptane chain length
Heptadecane chain length
Dodecane chain length
Conversion from
insoluble aqueous soln
metal oxides
Rate law
An emperimentally determined mathmatical expression showing the rate of a reaction as a function of the concentration of its reactants
good leaving groups
weak bases, halides
What type of equiillibrium is encountered in chemistry?
Noble gases
Contains nonmetals that are non-reactive. Full outermost energy level except helium which has 2.
Endothermic rxn
∆H is positive
Enthalpy is increased
Hess' Law
it's possible to algebraically add reactions to determine the overall enthalpy change for a complex series.
a dynamic condition in which two opposing changes occur at equal rates in a closed system
London forces
the weak attractive forces between molecules resulting from the small, instantaneous dipoles that occur because of the varying positions of the electrons during their motion about nuclei
having characteristics of both an acid and a base and capable of reacting as either
Changes in Gibbs Free Energy symbol
delta G
A base reacts with water as follows:
A- + H2O « HA + OH-
Kb = [HA] [OH-] / [A-]
(Note: the concentration of H2O is usually ignored)
pKb = -log10Kb
Note, Ka Kb = ([H+] [A-] / [HA]) x ([HA] [OH-] / [A-])
= [H+] [OH-]
= Kw
= 10-14
And since Ka Kb = 10-14
pKa + pKb = 14 = pKw
deactivating ortho/para directing (weak e- withdrawing)
F, Cl, Br, I
Bronsted-Lowry definition
Common definition of acids as proton (H+) donors and bases as proton acceptors
mixture of 2 or more substances that distills at a constant temperature and with constant composition, even though seperately the components have different boiling points
tells you how much solute is present compared to the amount of solvent
a pair of equal and opposite electric charges or magnetic poles separated by a small distance
Free radical
an uncharged molecule with a single unpaitred electron in its outer ring, very unstable, exists for only about 10 seconds
hydrogen bonding
the intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule
Group I; Valence Shell Config?
Alkali earth metals; ns1
LeChatelier's principle
When a system at equilibrium is subjected to a stress, it will shift in a direction that minimizes the effect of this stress.
Metallic Bonds
sea of e- surrounding positive metal ions
for any molec with n chiral centers there are 2 to the n sterioisomers
Electronegativity Difference
# found by taking th edifference between the e-negativities of 2 atoms. Value determines bond type.
Values differ by 1.7 or more... ionic
atoms which contain the same # of protons but diff. #'s of neutrons
List the strong acids
Hcl, HBr, HI, HNO3, HClO4, H2SO4
Solution equilibrium
When a solute is dissolved in a solvent, it will dissociate until reaching an equilibrium point at which the rate of dissociation equals the rate of precipitation of the solute, regardless of any additional solute introduced into the mixture
Formal Charge
Charge assigned to an atom in a molecule or polyatmic ion, calculated by (# valence electrons) - (# 1/2 bonding electrons) - (# nonbonding electrons). Molecules containing atoms with lower formal charges tend to be more stable than those with higher formal charges.
State Functions
A state function is a quantity whose value depends only on the state of the system and not on its history. X is a state function only if DX does not depend on the path used to go from the initial state to the final state of the system. V, G, P, H, E, S, and T are state functions.
Hess's law is a consequence of enthalpy being a state function. It states that DH of a reaction is the same regardless of whether the reaction occurs in a single step or in several steps.
First ionization energy
The energy required to remove the outermost electron from a neutral atom in the gas phase.
determining absolute configuration at a single chiral center
1)assign priority by atomic number
2)orient the molecule with the lowest priority substituent at the back
3)lowest priority to highest priority

Excited State
e- absorbs energy + moves to higher energy level above ground state
Activation Energy
Energy necessary for a rxn to begin. Obtained from the kinetic energy released durign a collsion.
What are the two completely seperate "sets of rules/considerations" for chemical reactions?
Thermo dynamics and Kinetics
representative elements
an element in an "A" group in the periodic table; as a group these elements display a wide range of physical and chemical properties. In their atoms, the s and p sublevels in the highest occupied energy level are partially filled
Percent composition
The Percent by mass of each element in a compound.
What does a negative delta H indicate
reaction is exothermic (exergonic)
Freezing Point Depression
DTFP = -kf x molality of solute particles
where kf is the molal freezing point depression constant
Law of Mass Action
Rate of chemical rxn is proportional to the product of the concentrations of the reactants
Hess's Law
∆H for a rxn is the same regardless of the path travelled from reactants to products
What is indicated by a negative delta H for a reaction
the process is exothermic
What is the only way to change the value of equilibrium constant, K
Change the temperature
Le chateliers Principle
States that if a stress is applied to a system at equilibrium, the system shifts in the direction that relieves the stress.
Ground State
e- is a t it's lowest energy levels as close to nucleus as possible
What are the two things that happen when energy is added to a substance?
Increase temperature (increase KE) or change phase (increase PE) assuming no chemical reaction occurs
constant temperature
Avagadro's Number
Kilo (k)
Qsp < Ksp
Heterogenous Catalyst
distinct phase
standard free energy
Average kinetic energy
Group II
alkali earth metals
Units of entropy (S)
Small discrete increments of energy.
electromagnetic energy of photons emmited from electrons at ground state
Unit cells
repeating units of cells
Combined Gas Law
P1V1/T1 = P2V2/T2
smaller than their atoms counterpart
Atomic Weight
Average mass of all naturally-occurring isotopes of a given element, measured in AMU.
Concentration of a solution calculated by (mole solute)/(1 kg solvent)
Lewis Definition
Acids defined as electron-pair acceptors and bases as electron-pair donors.
Network covalent
large molecular structures, strong covalent bonding, share qualities of IONIC AND COVALENT
The electrode at which reduction occurs during a cell's redox reaction electrons always flow to the cathode in an electrochemical cell
Does the periodic table arrange elements in increasing or decreasing atomic number?
Mole Fraction (x)
moles of solute/total moles
Standard Temperature and Pressure. 273 Kelvin (0 Celsius), 1 atmosphere (760 torr, 760 kPA).
theoretical yield =
actual /theoretical x 100
Ideal Gas
A hypothetical gas whose particles would occupy zero volume and have no attractive intermolecular forces.
A medium, usually liquid, which a solute dissolves into to create a solution.
Atomic Radius
Distance measured either between the nucleus and outermost electron of an atom or by the separation of the two nuclei in a diatomic element. Decreases from left to right, and bottom to top of the periodic table.
An atom or substance containing no unpaired electrons and is consequently repelled by a magnet.
Strong Acid
An acid that will completely dissociate in aqueous solution (e.g., HCL; HI;HCLO4)
Reversible Reaction
A process that will proceed bidirectionally to form both product and reactant
Ion Product
The product of the molar concentrations of dissociated ions in solution at any point in the reaction other than equilibrium or saturation, where each ion is raised to the power of its stoichiometric coefficient. Denoted IP
redox reaction
A chemical reaction involving the transfer of one or more electrons from one reactant to another; also called oxidation-reduction reaction.
Group 7A
halogens; ns2np5,, 2nd most reactive group, The Halogens; very active because of need to fill; form -1 ions; 7 electrons in valence shell; tend to form salts with elements from groups 1A and 2A
a subatomic particle that has a positive charge and that is found in the nucleus of an atom Charge of +1 and mass of 1.0073 amu
a substance that, when dissolved in water, results in a solution that can conduct electricity
Valence Electrons
The electrons occupying the outermost electron shell of an atom that participate in chemical bonds. Atoms with the same number of valence electrons usually have similar properties
The acronym for Valence Shell Electron Pair Repulsion theory, which states that the three-dimensional molecular geometry about some central atom is determined by the electronic repulsions between its bonding and nonbonding electron pairs
netowrk where all bonds are covalent (no imf bonds)
the average kinetic energy of the particles of the substance
Standard emf of a galvanic cell
is positive
same electrons so atom with largest nuclear charge or most protons will be smaller
Boyles Law
P1V1 = P2V2 T = constant
Catalysts and Equilbrium
catalysts move faster towards equilibrium but cant change eq.
Unit Cell Shapes
simple cubic, body centered, face-centered
Aufbau principle
an electron occupies the lowest-energy orbital that can receive it
Basic building block of all matter in the universe. Made up of 3 main components: protons, neutrons and electrons.
Effective Nuclear Charge
Resulting positive nuclear charge an outer electron senses after accounting for the shielding effect of inner core electrons. Abbreviated as Z(eff). Increases from left to right, and bottom to top on the Periodic Table.
A reaction in which a species gains electrons.
Double Displacement Reaction
Chemical reaction in which two different compounds exchange an atom or ion to form two new compounds, like switching components. Also called a metathesis reaction.
Reaction Quotient
Ratio of the concentrations of the products to the concentrations of the reactants at any point during the reaction aside from equilibrium, where each reactant and product in the expression is raised to the power of its stoichiometric coefficient. Commonly denoted by Q.
Dipole Moment
Product of the amount of partial charge at either end of a molecule's dipole multiplied by the distance between them, given equation "p=qd". P is the dipole moment, q is the partial charge, and d is the distance separating the dipole
atomic emission spectrum
a set of frequencies of electromagnetic waves given off by atoms of an element; consists of a series of fine lines of individual colors
quantum numbers
numbers that specify the properties of atomic orbitals and of their electrons
bond energy
the energy required to break a chemical bond and form neutral isolated atoms
Principal Quantum Number
The first quantum number. Designated by the letter n, it takes on any positive integer value and describes an electron's energy level. An electron with a higher n value is at a higher energy state
Molecular orbital
Region in a molecule where atomic orbitals overlap, resulting in either a stable low-energy bonding orbital or an unstable high-energy antibonding orbital.
Electron Configuration
The patterned order by which electrons fill subshells and energy levels in an atom. The first number designates the principal quantum number (n); the letter-s, p, d, f, or g-specifies the subshell (L); and the superscript indicates the number of electrons in that subshell
Ionization energy
the energy required to remove an electron from a gaseous atom; generally increases in moving from left-to-right across a period and decreases in moving down a group
Boyle's Law
At a constant temperature, the volume of an ideal gas is inversely proportional to its pressure: V Alpha 1/P
Electrolytic Cell
An electrochemical cell that uses an external electric source to drive a non spontaneous (unfavorable) redox reaction to proceed.
A redox reaction in which the same species is both oxidized and reduced
Net ionic equation
A representation of a displacement reaction showing only the reactive species and omitting the spectator ions.
an atom, radical, or molecule that has gained or lost one or more electrons and has a negative or positive charge
Phase diagram
a graph of pressure versus temperature that shows the conditions under which the phases of a substance exist
Which would have a higher ionization energy? group 1 or group 7?
group 7
Entropy when more moles on product side
entropy increases
Hess Law
the enthalpy change for any overall process is equal to the sum of enthalpy changes for any set of steps that leads from the reactants to the products
Cell Potential
E(cell) = E cath - E anode
a phase change in which solid converts directly a vapor without passing through the liquid phase
ΔEevap = ΔH vap - RT vap
Group 2A
Alkaline earth metals are harder, more dense, and melt at high temperatures than alkali metals. The form 2+ cations. They are less reactive than alkali earth metals
pH of a molecule at which it contains no net electric charge, isoelectric point.
Solubility Product Constant
Product of the molar concentrations of dissociated ions in solution at saturation, where each ion is raised to the power of its stoichiometric coefficient. Denoted Ksp.
Unit of energy in the form of light equal to hf, where h is Planck's constant and f is the frequency of radiation.
Lewis Structure
A method using lines and dots to represent valence electrons, and shared pairs of electrons of atoms, ions, or molecules.
Ionic Bonds
Type of chemical bond in which there is a complete transfer of valence electrons to form positive and negative ions that are subsequently bound by electrostatic forces; strong attractions holding ions together in an ionic compound.
Equivalence Point
Point in a titration at which an equimolar amount of titrant has been added to the unknown solution.
Transition Elements
-B Groups of the Periodic Table. These contain partially filled d subshells.
Endothermic Reaction
A reaction that proceeds with the net absorption of energy, heat, from the surroundings.
Closed System
A system that allows for the exchange of energy, but not matter, across its boundaries.
Principle quantum number
The quantum number that indicates the main energy level occupied by the electron. Can theoretically take on any positive interger. Denoted by the letter n.
Heisenberg Uncertainty Principle
The quantum mechanical idea that we cannot measure the exact momentum and position of an orbiting electron simultaneously. That is, the more accurately we measure an electron's momentum, the less we know about its exact position
Standard Hydrogen Electrode (SHE)
reduction potential --> higher value ---> more likely to be reduced
Buffer Solution
contains either a weak base and its conjugate acids or a weak acid and its conjugate bse
pH = pKa + log ( [A-] intial / [HA] initial)
Delta G, in terms of EMF
Delta G = -nFE
molecular formula
the exact number of elemental atoms in each molecule
Standard Reduction Potential
Tendency of a species to be reduced, as measured at 25˚C when reacting species are of 1M concentration or 1atm partial pressure (for gases).
Formula Weight
Sum of all the masses, in AMU, present in one molecule of a molecular compound.
Gibbs Free Energy
Energy of a system available to do work.
Mass Number
Sum of the protons and neutrons in an element, often denoted by the letter A
Emperical Formula
simplest whole # ration of atoms in a compound
s orbital
have the shape of a sphere, with the center of the sphere at the nucleus; completely symmetrical along all axes; 1s orbital is spherically symmetric and has no nodes; 2s orbital is also spherical but contains a node and is higher in energy
gram equivalent weight
the weight in grams of compound that can be substituted by 1 atom of Hydrogen. GEW = MW / # of acidic Hyrdogens
Temperature in equilbrium
the ONLY variable that will cause a change in the value of Keq an increase in temperature always shift the eq position in the endothermic direction
Chemical Change of State
The amount of reactants and products change
Behavior of Gas Kinetic Molecular Theory
same behavior regardless of identity
Neutralization Reaction
A reaction in which an acid and a base are combined to form water and a salt.
Diprotic Base
A base that can accept two moles of H+ per mole of itself (ex: SO₄²-).
real gasses at moderately high pressure
gas volumne is less than predicted
Atomic Mass Unit
Unit of mass equal to 1/12 gram of a carbon-12 atom, roughly equal to the mass of one proton.
Chemical Kinetics rate
rate = f * z
f = fraction of collision that are effective
z = total collisions per second
Kinetic Molecular Theory of Gases
A series of ideas used to account for the behavior of ideal gases. The theory describes gas as volumeless particles in constant, random motion that exhibit no intermolecular attractions and undergo completely elastic collisions with each other and their container walls
IMF forces are _____ a physical change because _______
not, because bonds are broken and reformed
standard entropy
Centi (c)
Qsp > Ksp
1 cal
4.814 J
(chemistry) p(otential of) H(ydrogen)
Reactant concentrations
reaction order dependent
Qc 〉 Keq
exceeded equilbrium
electrode where oxidation occurs,, half reactions with the more negative reduction potential occurs at the node
Capacitance is proportional to what?
Reaction Mechanism
Play-by-play showing the individual steps of a reaction, including the formation and destruction of any reaction intermediates that may occur.
Atom or substance containing unpaired electrons and is consequently attracted by a magnet.
(chemistry) a substance that changes color to indicate the presence of some ion or substance
Group 6A
chalcogens,, Oxide O²⁻, Sulfide S²⁻, Selenide Se²⁻ - chalcogens - gain 2 electrons to become a noble gas - s2p4 Oxygen is the second most electronegative element.
Group 1A
Alkali metals: highly reactive, therefore always compounds., 1 valence electron +1 ion, Hydrogen H⁺, Lithium Li⁺, Sodium Na⁺, Potassium K⁺, Rubidium Rb⁺, Cesium Cs⁺
A cagelike network of solvent molecules that forms around a solute in a solution
Anode Electrolytic
positive (because attracted to positive poll of source)
Real Gases
1. have intermolecular attractions (non-elastic collisions) called Van Der Waals Force
2. contain molecules that have volume (or measurable size)
Intermolecular Forces
1. Dispersion Forces
2. Dipolar Forces
3. Hydrogen Bonding
3rd atomic number
ml (1 to -1)
moles =
grams/ atomic or molecular weight
Smallest unit of a substance, composed of two or more atoms joined in covalent bonds, retaining all the chemical properties of that substance.
Dense, positively charged center of an atom, containing protons and neutrons.
Atoms that share the same atomic number (Z) but have a different number of neutrons.
Avogadro's Principle
Principle stating that when different gases of equal volumes are at identical temperatures and pressures, they contain equal numbers of molecules.
Redox Half-Reaction
The hypothetical equation showing only the species that is oxidized or reduced in a redox reaction and the correct number of electrons transferred between the species in the complete, balanced equation
Graham's Law
temperature is constant; effusion and temperature are proportional to the square root of their masses
Colligative Properties
The properties of solutions-such as vapor pressure lowering, freezing point depression, boiling point elevation, and osmotic pressure-that are affected only by the number of solute particles dissolved and not by their chemical identities
in a solution, the substance that dissolves in the solvent
A subatomic particle that orbits the nucleus and has a charge of -1. The electron has a negligible mass and is often denoted by the symbol e-
Kinetic Energy
proportional to temperature and speed is proportional to the square root of kinetic energy
What is a prereq for hydrogen bonding?
Large Keq
indicates that the reaction goes virtually to completion
Boiling Point Elevation
ΔTb = kb m i
Molality (m)
moles of solute/ kg of solvent
measure of how strong an element attracts electrons in a bond
negative Ions are much _______ than neutral
The characteristics of metals
lustrous, ductile, malleable, thermally, and electrically conductive
Raoult's Law
Vapor pressure of one component above a solution is proportional to the mole fraction of that component in the solution. Pa = Xa - P(total).
Atomic Emission Spectra
Discontinuous line spectra of light produced when excited atoms return to their ground state and emit photons of a certain frequency.
Aqueous Solution
A solution containing water as its solvent.
Adiabatic Process
Process in which no heat is transferred to or from the system by its surroundings.
Dipole-Dipole Interaction
Type of intermolecular force in which opposite poles of neighboring dipole molecules are drawn together.
Covalent Bond
A chemical bond formed when atoms share bonding electron pairs
Spin Quantum Number
The fourth quantum number. Designated by ms, it specifies an electron's intrinsic spin value or angular momentum in an orbital. Since there can be no more that two electrons per orbital, the value of ms can only be +1/2 or -1/2
Alkali Metals
The highly reactive elements found in Group IA of the periodic table, except hydrogen
Resonance structure
structure that occurs when it is possible to draw two or more valid electron dot structures that have the same number of electron pairs for a molecule or ion
Magnetic Quantum Number
The third quantum number. Designated by ml, it describes a particular orbital within a subshell where an electron is very likely to be found. The possible values for ml are integers in the -l to l range, including 0
Galvanic Cell
An electrochemical cell powered by a spontaneous redox reaction that produces an electric current flow; voltaic cell
A ratio that measures how much solute can dissolve in a solvent at a given temperature, expressed in units of (g solute)/(100 g solvent)
An increase in atomic number means an (increase, decrease) in first ionization energy
Molecular Solid
molecules are held in place dispersion forces, dipolar internactions and/or hydrogen bonding
Entropy when energy out of the system
entropy decreases
is the transfer of energy from one substance to another as a result of the change in temperature
Standard Enthalpy of Formation
enthalpy change accompanying the formation of one mole of chemical substance from pure elements in their most stable form under standard conditions
Transition State
High energy complex in which old bonds are partially broken and new bonds are partially formed. Charges existing only prior to or after the formation of the complex are designated as partial charges.
State Function
A function that depends only on the initial and final states of a system, not on the path in between.
Open System
A system that allows for the exchange of energy and matter across its boundaries.
Henry's Law
Partial pressure of a gas dissolved in a solution is directly proportional to the partial pressure of this gas above the solution.
Arrhenius Definition
A definition of acids as producers of H+ and bases as producers of OH- in aqueous solution
Quantum Mechanics
Study of physics at the atomic level where energy is quantized in discrete, rather than continuous, levels
Electron Affinity
The energy released when an atom or ion in the gaseous state gains an electron. increases from left to right and from bottom to top on the periodic table
Octet Rule
A rule stating that atoms-except a few such as Be, H, and B-tend to react in order to form a complete octet of valence electrons. Hydrogen can have a maximum of 2 valence electrons, Be can have 4 valence electrons, and B can have 6 valence electrons
theoretical yield
the maximum amount of product that can be produced from a given amount of reactant
Ion dipole interactions
When dipoles are dissolved in a solution where ions are present ions wil arrange themselves with the opposite charged end of the dipole.
Second Law of Thermodynamics
any spontaneous process increases the disorder of the universe
Electromotive Force (emf)
or E°cell is the difference in potential between two half cells
Group 4A elements
the elements can form four covalent bonds with nonmetals. All but Carbon can can form two additional bonds with Lewis bases.
Triple Point
Point on a phase diagram at which a substance exists in equilibrium between all three phases.
Polar Covalent Bond
A type of covalent bond between atoms with different electronegativities that results in an unequal sharing of electron pairs, giving the bond partial positive and negative poles
Bronsted Lowry
A model of acids and bases which an acid is hydrogen ion donor and base is a hydrogen ion acceptor.
Common ion effect
The molar solubility of one salt is reduced when another salt, having a common ion is brought into the same solution
Single Displacement Reaction
A chemical reaction in which an atom or ion of one compound is replaced by another atom or ion (e.g., A + BC -> B + AC)
Nonpolar covalent bond
A covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge
The Periodic Law states...
The properties of the elements are periodic functions of their atomic numbers
Entropy equation
ΔS = Q / T
spread of heat over time
Molar Heat of Solution
measures net energy flow that occurs as substance dissolves
Buffer Capacity
the amount of added H+ or OH- the buffer solution can tolerate without exceeding a specified pH range
weak pi bonds
large atoms are unable to make strong pi bonds
Pauli Exclusion Principle
No two electrons in an atom can have the same set of four quantum number values.
Lewis acid base reaction
the formation of one or more covalent bonds between an electron pair donor and an electron pair acceptor
Balancing a Half Reaction
1. Balance the elements other than the H and O
2. balance the O atoms by adding H20
3. Balance the H atoms by adding H+
4. Balance the charge by adding e-
Atomic radius
Metallic Character
it increases going to the left and moving from top to bottom
Collision Theory of Chemical Kinetics
Theory Stating that the rate of a reaction is directly proportional to the number of collisions that take place between reactants per second
The collision theory of chemical reactions
the rate of reaction is proportional to the number of collisions per second between reacting molecules
What are the charges at the electrodes in an electrolytic cell?
Cathode (-) and anode (+)
Isoelectronic ions (ions with the same number of electrons)
tend to get smaller with increasing atomic number because more protons pull inward on the same number of electrons.
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