Definitions Flashcards
Absolute zero
The temperature at which all substances have no thermal energy; 0 K or -273.15°C.
Absorption spectrum
The series of discrete lines at characteristic frequencies representing the energy required to excite an electron from the ground state.
Acid
A species that donates hydrogen ions or accepts electrons.
Acid dissociation constant (Ka)
The equilibrium constant that measures
the degree of dissociation of an acid under specific conditions.
Acidic solution
An aqueous solution that contains more H+ ions than OH- ions; pH <7 under standard conditions.
Actinide series
The series of chemical elements atomic numbered 89-103 and falling between the s and d blocks on theperiodic table.
Activation energy (Ea)
The minimum amount of energy required for a reaction to reach the transition state; also called energy barrier.
Actual yield
The experimental quantity of a substance obtained at the end of a reaction.
Adiabatic process
A process that occurs without the transfer of heat into or out of the system.
Alkali metals
Elements found in Group IA of the periodic table; highly reactive, readily losing one valence electron to form ionic cornpourtds with nonmetals.
Alkaline earth metals
Elements found in Group IIA ofthe periodic table; chemistry is similar to that of the alkali metals, except that they have two valence electrons and, thus, form +2 cations.
Amphiprotic species
A species that may either gain or lose a proton.
Amphoteric species
A species capable of reacting as either an acid or base, depending on the nature ofthe reactants.
Angualar momentum
The rotational analog oflinear momentum.
Anion
An ionic species witha negative charge.
Anode
The electrode at which oxidation occurs.
Antibonding orbital
A molecular orbital formed by the overlap of two or more atomic orbitals; energy is greater than the energy of the combining atomic orbitals.
Aqueous solution
A solution in which water is the solvent.
Arrhenius acid
A species that donates protons (H+) in aqueous solution.
Arrhenius base
A species that donates hydroxide ions (OH-) in aqueous solution.
Arrhenius equation
A chemical kinetics equation that relates the rate constant (k) of a reaction with the frequency factor (A), the activation energy (Ea), the ideal gas constant (R), and temperature (T) in kelvin.
Atom
The smallest unit of an element that retains the properties of the element; it cannot be further broken down by chemical means.
Atomic mass
The mass of a given isotope of an element; closely related to the mass number.
Atomic mass unit (amu)
A unit ofmass defined as 1/12 the mass of a carbon-12 atom; approximately equal to the mass of one proton or one neutron.
Atomic number
The number of protons in a given element.
Atomic orbital
Describes the region of space where there is a high probability of finding an electron.
Atomic radius
The average distance between a nucleus and its outer- most electron; usually measured as one-half the distance between two nuclei of an element in its elemental form.
Atomic weight
The weighted average mass of the atoms of an element, taking into account the relative abundance of all naturally occurring isotopes.
Aufbau principle
The concept that electrons fill energy levels in order of increasing energy, completely filling one sublevel before beginning to fill the next.
Autoionization
The process by which a molecule (usually water) spontaneously dissociates into cations and anions.
Avogadro’s number
The number of atoms or molecules in one mole of a substance: 6.02 x 10^23 mol^-1..
Avogadro’s principle
The law stating that under the same conditions of temperature and pressure, equal volumes of different gases will have the same number of molecules.
Azimuthal quantum number (I)
The quantum number denoting the sub-level or subshell in which an electrons can be found; reveals the shape of the orbital.
Balanced equation
An equation for a chemical reaction in which the number of atoms for each element in the reaction and the total charge are
the same for the reactants and the products.
Balmer series
Part of the emission spectrum for hydrogen, representing transitions of an electron from energy levels n > 2 to n= 2.
Barometer
A tool for measuring pressure.
Base
A species that donates hydroxide ions or electron pairs or that accepts
protons.
Base dissociation constant (Kb)
The equilibrium constant that measures the degree of dissociation for a base under specific conditions.
Basic solution
An aqueous solution that contains more OH- ions than tt+ ion; pH > 7 under standard conditions.
Bohr model
The model of the hydrogen atom in which electrons assume certain circular orbits around a positive nucleus.
Boiling point
The temperature at which the vapor pressure of a liquid is equal to the incident pressure; the normal boiling point of any liquid is defined as its boiling point at a pres- sure of 1 atmosphere.
Boling point elevation
The amount by which a given quantity of solute raises the boiling point of a liquid; a colligative property.
Bond energy
The energy (enthalpy change) required to break a particular bond under given conditions.
Bond enthalpy
The average energy that is required to break a particular type of bond between atoms in the gas phase.
Bonding electrons
Electrons located in the valence shell of an atom and involved in a covalent bond.
Bonding orbital
A molecular orbital formed by the overlap of two or more atomic orbitals; energy is less than that of the combining orbitals.
Bond length
The average distance between two nuclei in a bond; as the number of shared electron pairs increases, the bond length decreases.
Bond order
The number of shared electron pairs between two atoms; a single bond has a bond order of 1, a double bond has a bond order of 2, a triple bond has a bond order of 3.
Boyle’s law
The law stating that at constant temperature, the volume of a gaseous sample is inversely proportional to its pressure.
Broken-order reaction
A reaction with noninteger orders in its rate law.
Bronsted-Lowery acid
A proton donor.
Bronsted-Lowery base
A proton acceptor.
Buffer
A solution containing a weak acid and its salt (or a weak base and its salt) that tends to resist changes in pH.
Buffer region
The portion of a titration curve in which the concen- tration of an acid is approximately equal to that of its conjugate base; pH remains relatively constant through this region.
Buffering capacity
The degree to which a system can resist changes in pH.
Calorie (cal)
A unit of thermal energy.
Calorimeter
An apparatus used to measure the heat absorbed or released by a reaction.
Catalyst
A substance that increases the rates of the forward and reverse directions of a specific reaction by lowering activation energy, but is itself left unchanged.
Cathode
The electrode at which reduction takes place.
Cation
An ionic species with a positive charge.
Celsius (°C)
A temperature scale defined by having 0°C equal to the freezing point of water and 100°C equal to the boiling point of water; otherwise known as the centigrade temperature scale.
Chalcogens
Elements found in Group VIA of the periodic table with diverse chemistry; the group contains metals, nonmetals (like oxygen), and metalloids; typically form -2 anions.
Charging
A state of an electrochemical cell in which an external electromotive force is being used to return a cell to its original state; during this process, electrons are transferred nonspontaneously from cathode to anode.
Charle’s law
The law stating that the volume of a gas at constant pressure is directly proportional to its absolute (kelvin) temperature.
Chelation
The process of binding metal ions to the same ligand at multiple points.
Chemical bond
The interaction between two atoms resulting from the sharing or transfer of electrons.
Chemical equation
An expression used to describe the quantity and identity of the reactants and product of a reaction.
Chemical properties
Those properties of a substance related to the chemical changes that it undergoes, such as ionization energy and electronegativity.
Closed system
A system that can exchange energy but not matter with its surroundings.
Colligative properties
Those properties of solutions that depend only on the number of solute particles present but not on the nature of those particles.
Collision theory of chemical kinetics
A theory that states that the rate of a reaction is proportional to the mun- ber of collisions per second between reacting molecules that have sufficient energy to overcome the activation energy barrier; implies that only a fractwn ofcollisions are sufficient.
Comination reaction
A reaction in which two or more reactants form a single product.
Combined gas law
A gas law that combines Boyle’s law, Charles’s law, and Gay-Lussac’s law to state that pressure and volume are inversely proportional to each other, and each is directly proportional to temperature.
Combustion reaction
A reaction in which an oxidant (typically oxygen) reacts with a fuel (typically a hydro- carbon) to yield water and an oxide (such as carbon dioxide if between a hydrocarbon and oxygen).
Common ion effect
A shift in the equilibrium of a solution due to the addition of ions of a species already present in the reaction mixture.
Complexation reaction
A reaction in which a central cation is bound to one or more ligands.
Complex ion
A polyatomic mol- ecule in which a central cation is bonded to electron pair donors called ligands.
Compound
A pure substance that can be decomposed to produce elements, other compounds, or both.
Compression
Reduction in the volume of a gas.
Concentrated solution
A solution with a high concentration value; the cutoff for the term “concentrated” depends on the purpose and identity of the solution.
Concentration
The amount of solute per unit ofsolvent or the relative amount of one component in a mixture.
Concentration cell
A cell that creates an electromotive force (emf or voltage) using a single chemical species in half- cells of varying concentration.
Condensation
The process in which a gas transitions to the liquid state.
Conductor
A material in which electrons are able to transfer energy in the form of heat or electricity.
Conjugate acid-base pair
The relationship between a Bronsted-Lowry acid and its deprotonated form, or a Bronsted-Lowry base and its protonated form.
Coordinate covalent bond
A covalent bond in which both electrons of the bonding pair are donated by one of the bonded atoms.
Coordination number
The number of atoms that are bound to a central atom.
Covalent bond
A chemical bond formed by the sharing of an electron pair between two atoms; can be in the form of single bonds, double bonds or triple bonds.
Critical point
The point in a phase diagram beyond which the phase boundary between liquid and bas no longer exists.
Critical pressure
The vapor pressure at the critical temperature of a given substance.
Critical temperature
Also known as the critical point. The highest temperature at which the liquid and gas phases of a substance can coexist; above this temperature, the liquid and gas phases are indistinguishable.
Crystal
A solid in which atoms, ions, or molecules are arranged in a regular, three-dimensional lattice structure.
d subshell
Subshell corresponding to the angular momentum quantum number / = 2; contains five orbitals and is found in the third and higher principal energy levels.
Dalton’s law of partial pressures
The law stating that the sum of the par- tial pressures of the components of a gaseous mixture must equal the total pressure of the sample.
Daniell cell
An electrochemical cell in which the anode is the site of Zn metal oxidation and the cathode is the site of Cu2+ ion reduction.
Decomposition reaction
A reaction in which a single compound breaks down into two or more products.
Delocalized orbitals
Molecular orbitals in which electron density is spread over an entire molecule, or a portion thereof, rather than being localized between two atoms.
Density (p)
A physical property of a substance, defined as the mass contained in a unit of volume.
Deposition
In most chemical processes, the direct transition of a substance from the gaseous state to the solid state; in elec- trochemical reactions, the build up of a solid precipitate onto an electrode.
Diamagnetism
A condition that arises when a substance has no unpaired electrons and is slightly repelled by a magnetic field
Diffusion
The random motion of gas or solute particles across a concentration gradient, leading to uniform distribution of the gas or solute throughout the container.
Dilute solution
A solution with a low concentration of a given solute.
Dipole
A species containing bonds between elements of different electronegativities, resulting in an unequal distribution of charge.
Dipole-dipole interactions
The attractive forces between two dipoles; magnitude is dependent on both the dipole moments and the distance between the two species.
Dipole moment
A vector quantity with a magnitude that is dependent on the product of the charges and the distance between them; oriented from the positive to the negative pole.
Discharging
The state of a rechargeable electrochemical cell that is providing an electromotive force by allowing electrons to flow spontaneously from anode to cathode.
Disproportionation
An oxidation-reduction reaction in which the samr species acts as the oxidizing agent and as the reducing agent; also called dismutation.
Dissociation
The separation of a single species into two separate species; usually used in reference to salts or weak acids or bases.
Double-displacement reaction
A reaction in which ions from two different compounds swap their associated counterions; typically, one of the products of this type of reaction is insoluble in solution and will precipitate.
Ductilitiy
The property of metals that allows a material to be drawn into thinly stretched wires.
Effective nuclear charge (Zeff)
The charge perceived by an electron from the nucleus; applies most often to valence electrons and influences periodic trends such as atomic radius and ionization energy.
Effusion
The movement of gas from one compartment to another under pressure through a small opening; follows Graham’s law.
Electrocchemical cell
A cell within which an oxidation-reduction reaction takes place, containing two electrodes between which there is an electrical potential difference.
Electrode
An electrical conductor through which an electrical current enters or leaves a medium.
Electrolysis
The process in which an electrical current is used to power an otherwise nonspontaneous decomposition reaction.
Electrolyte
A compound that ionizes in water and increases the conductance of the solution.
Electrolytic cell
An electrochemical cell that uses an external voltage source to drive a nonspontaneous oxidation-reduction reaction.
Electromaagnetic radiation
An electrochemical cell that uses an external voltage source to drive a nonspontaneous oxidation-reduction reaction.
Electrolmagnetic spectrum
The range of all possible frequencies or wavelengths of electromagnetic radiation.
Electromotive force (emf)
The potential difference developed between the cathode and the anode of an electrochemical cell; also called voltage.
Electron (e-)
A subatomic particle that remains outside the nucleus and carries a single negative charge; in most cases, its mass is considered to be negligible.
Electron affinity
The energy dissipat-ed by a gaseous species when it gains an electron.
Electron configuration
The symbolic representation used to describe the electron arrangement within the energy sublevels in a given atom.
Electron spin
The intrinsic angular momentum ofan electron, represented by ms; has arbitrary values of + 1/2 and - 1/2.
Electronegativity
A measure of the ability of an atom to attract the electrons in a bond; commonly measured with the Pauling scale.
Electronic geometry
The spatial arrmgement of all pairs of electrons around a central atom, including both the bonding and lone pairs.
Electron shell
The space occupied by/path followed by an elec- tron around an atom’s nucleus. Electron shell (also called principle energy level) for a given electron is indicated by its principle quantum number.
Element
A substance that cannot be further broken down by chemical means; defined by its number of pro- tons (atomic number).
Emission spectrum
A series of discrete lines at characteristic frequen- cies, each representing the energy emitted when electrons in an atom return from an excited state to their ground state.
Empirical formula
The simplest whole-number ratio of the different elements in a compound.
Endothermic reaction
A reaction that absorbs heat from the surroundings as the reaction proceeds (positive delta H).
Endpoint
The point in a titration at which the indicator changes to its final color.
Energy density
An equivalence unit regarding the amount of electrochemical energy capable of being stored per unit weight; a battery with a large energy density can produce a large amount of energy with a small amount of material.
Enthalpy (H)
The heat content of a system at constant pressure; the change in enthalpy (delta H) in the course of a reaction is the difference between the enthalpies of the products and the reactants.
Ethropy (S)
A property related to dispersion of energy through a system or the degree of disorder in that system; the change in en- tropy (6.S) in the course of a reac- tion is the difference between the entropies of the products and the reactants.
Equilibrium
The state of balance in which the forward and reverse reaction rates of a reversible reaction are equal; the concentrations of all species will remain constant over time unless there is a change in the reaction conditions.
Equilibrum constant (Keq)
The ratio of the concentrations of the products to the concentrations of the reactants for a certain reaction at equilibrium, all raised to their stoichiometric coefficients.
Equivalence point
The point in a titration at which the moles of acid present equal the moles of base added, or vice-versa.
Equivalent
A mole of charge in the form of electrons, protons, ions, or other measurable quantities that are produced by a substance.
Evaporation
The transition from a liquid to a gaseous state.
Excess reagent
In a chemical reaction, any reagent that does not limit the amount of product that can be formed.
Exitation
The promotion of an electron to a higher energy level by absorption of an energy quantum.
Excited state
An electronic state having a higher energy than the ground state; typically attained by the absorption of a photon of a certain energy.
Exothermic reaction
A reaction that gives off heat to the surroundings (negative delta H) as the reaction proceeds.
f subshell
The subshell corresponding to the angular momentum quantum number I = 3; contains seven orbitals and is found in the fourth and higher principal energy levels.
Faraday constant (F)
The total charge on I mole of electrons
[ F= 96, 485 C/ mole-]; not to be confused
with the farad (also denoted F), a unit of capacitance.
First law of thermodynamics
The law stating that the total energy of a system and its surroundings remains constant.
First-order reaction
A reaction in which the rate is directly proportional to the concentration of only one reactant.
Fluid
A substance that flows due to weak intermolecular attractions between molecules and that takes the shape of its container; liquids and gases are considered fluids.
Formal charge
The conventional assignment of charges to individual atoms of a Lewis structure for a molecule; the total number of valence electrons in the free atom minus the total number of electrons when the atom is bonded (assuming equal splitting of the electrons in bonds).
Formula weight
The sum of the atomic weights of constituent ions according an ionic compound’s empirical formula.
Freezing
The process in which a liquid transitions to the solid state; also known as solidification or crystallization.
Freezing point
At a given pressure, the temperature at which the solid and liquid phases of a substance coexist in equilibrium; identical to the melting point.
Freezing point depression
Amount by which given quantity of solute lowers the freezing point of a liquid; a colligative property.
Galvanic cell
An electrochemical cell that uses a spontaneous oxidation- reduction reaction to generate an electromotive force; also called a voltaic cell.
Galvanization
In electrochemical cells, the precipitationn process onto the cathode itself; also called plating.
Gas
The physical state of matter possessing the most disorder, in which molecules interact through very weak attractions; found at relatively low pressure and high temperatures.
Gas constant (R)
A proportionality constant that appears in the ideal gas law equation, PV = nRT. Its value depends on the units of pressure, temperature, and volume used in a given situation.
Gay-Lussac’s law
The law stating that the pressure of a gaseous sample at constant volume is directly proportional to its absolute temperature.
Gibbs free energy
The energy of a system available to do work. The charge in Gibbs free energy, delta G can be determined for a given reaction equation from the enthalpy change, temperature, and entropy change; a negative delta G denotes a spontaneous reaction, while a positive delta G denotes a nonspontaneous reaction.
Graham’s law
The law stating that the rate of effusion or diffusion for a gas is inversely proportional to the square root of the gas’s molar mass.
Gram equivalent weight (GEW)
The amount of a compound that contains 1 mole of reacting capacity when fully dissociated; one GEW equals the molar mass divided by the reactive capacity (how many of the species of interest is obtained) per formula unit.
Ground state
The unexcited state of an electron.