General Chemistry Flashcards
1
Q
Activation Energy
A
The energy barrier that must be overcome for a reaction to proceed; often denoted by Ea. Activation energy is measured in joules (J).
2
Q
Alkali Metals
A
The highly reactive elements found in Group IA (Group 1) of the Periodic Table, except hydrogen.
3
Q
Alkaline Earth Metals
A
Elements found in Group IIA (Group 2) of the Periodic Table.
4
Q
Amphoteric
A
A species capable of reacting as either an acid or a base. In the Bronsted-Lowry sense, a species that can pick up or give off a proton. In the Lewis sense, a species that can denote or accept a lone pair of electrons.
5
Q
Anode
A
The electrode where oxidation occurs during a cell’s oxidation-reduction reaction. Electrons always flow from the anode in an electrochemical cell.
6
Q
Aqueous Solution
A
A solution containing water as its solvent.
7
Q
Atomic Radius
A
The 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 from bottom to top on the Periodic Table.
8
Q
Avogadro’s Principle
A
States that the number of moles of a gas present is proportional to its volume, assuming constant pressure and temperature.
9
Q
Boyle’s Law
A
States that at a constant temperature, the volume of an ideal gas is inversely proportional to its pressure.
10
Q
Bronsted-Lowry Definition
A
Common definition of acids as proton (H+) donors and bases as proton acceptors.
11
Q
Buffer
A
A solution containing a weak acid or base coupled with its conjugate salt, acting to prevent changes to the solution’s pH upon the addition of acidic or basic substances.
12
Q
Cathode
A
The electrode at which reduction occurs during a cell’s oxidation-reduction reaction. Electrons always flow toward the cathode in an electrochemical cell.
13
Q
Charles’s Law
A
States that at a constant pressure, the volume of an ideal gas is directly proportional to its temperature.
14
Q
Chemical Kinetics
A
The study of reaction rates and the factors that affect them.
15
Q
Colligative Properties
A
The properties of solutions- such as vapor pressure lowering, freezing point of depression, boiling point elevation, and osmotic pressure- that are affected only by the number of solute particles dissolved and not by their chemical identities.
16
Q
Collision Theory of Chemical Kinetics
A
Theory stating that the rate of a reaction is directly proportional to the number of collisions that take place between reactants per second.
17
Q
Concentration
A
The ratio of the amount of solute to the amount of solution; quantified by mole fraction, molarity, molality, or normality, among other measures
18
Q
Conjugate Acids and Bases
A
A systematic pairing of a protonated species (acid) with its deprotonated species (base) with its protonated form (conjugate acid). Conjugates appear on opposite sides of a chemical equation.
19
Q
Decomposition Reaction
A
A chemical reaction in which one substance breaks down into two substances (such as A => B + C)
20
Q
Diamagnetic
A
An atom or a substance that contains no unpaired electrons and is consequently repelled by a magnet.
21
Q
Diffusion
A
Passive dispersion of a gas or solute throughout a medium by means of random motion.
22
Q
Dipole Moment
A
The product of the magnitude of either partial charge in a dipole multiplied by the distance between the charges, given by: p = qd where p is the dipole moment, q is the partial charge, and d is the displacement vector separating the charges.
23
Q
Double-Displacement Reaction
A
A chemical reaction in which two different compounds exchange an atom or ion to form two new compounds
(such as AB + CD => AD + CB); also called a metathesis reaction.
24
Q
Electrochemical Reaction
A
A chemical reaction that either is driven by or produces electricity.
25
Electrolyte
An ion in water that is capable of conducting electricity in that solution.
26
Electrolytic Cell
An electrochemical cell that uses an external electric source to drive a nonspontaneous (unfavorable) oxidation-reduction reaction.
27
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.
28
Electronegativity
A measure of an atom's ability to pull electron density toward itself when involved in a chemical bond. Increases from left to right and from bottom to top on the Periodic Table.
29
Empirical Formula
Chemical formula showing the smallest whole-number ratio of atoms in a compound.
30
Endothermic Reaction
A reaction that proceeds with the net absorption of heat from the surroundings.
31
Enthalpy
The total heat content of a system at a constant pressure, commonly denoted by H.
32
Entropy
The chaos or randomness of a system often denoted by S. "Change in S" represents the change in entropy following a reaction.
33
Equilibrium
A dynamic point reached by a reversible reaction in which the rate of the forward reaction is equal to the rate of the reverse reaction. There is no net charge in the concentrations of the products and reactants over time.
34
Equivalence Point
The point in a titration at which an equimolar amount of titrant has been added to the unknown solution.
35
Exothermic
A reaction that proceeds with the net release of heat into the surroundings.
36
Formal Charge
The charge assigned to an atom in a molecule or polyatomic ion, assuming even division of the electrons in a bond. Molecules containing atoms with lower formal charges tend to be more stable than those with higher formal charges.
37
Free Radical
An atom or molecule that has an unpaired electron in its outermost shell.
38
Galvanic Cell
An electrochemical cell powered by a spontaneous oxidation-reduction that produces an electric current; also called a voltaic cell.
39
Gay-Lussac's Law
States that at a constant volume, the pressure of an ideal gas is directly proportional to its temperature.
40
Gibbs Free Energy
The energy of a system available to do work.
41
Half-Cell
An electrode immersed in an electrolytic solution that is the site of either oxidation or reduction in an electrochemical cell.
42
Half-Equivalence Point
The point in a titration at which exactly half the molar equivalence of reactant is consumed by the titrant being added. At this point in an acid-base titration, the pH is equal to the pKa of the titrand.
43
Halogens
Elements found in Group VIIA (Group 17) of the Periodic Table.
44
Heat
A transfer of energy from a substance with a higher temperature to a substance with a lower temperature.
45
Handerson-Hasselbalch Equation
An equation commonly used in titration-based problems that relates the pH or pOH of a solution to the pKa and the ratio of the dissociated species.
pH = pKa + log (A-/HA)
46
Henry's Law
States that the partial pressure of a gas dissolved in a solution is directly proportional to the partial pressure of this gas above the solution.
47
Hess's Law
States that the enthalpy change of an overall reaction is equal to the sum of the standard heats of formation of the products minus the sum of the standard heats of formation of the reactants.
48
Hund's Rule
States that electrons will first fill orbitals within a subshell unpaired and with parallel spins before being couples with other electrons of opposite spins in the same orbital. This method of maximizing the number of half-filled orbitals allows for the most stable distribution of electrons within a subshell.
49
Hydrogen Bonding
Very strong intermolecular force where a hydrogen covalently bonded to an N, O, or F is attracted to another N, O, or F.
50
Ideal Gas Law
A unification of Boyle's law, Charles's law, Gay-Lussac's law, and Avogadro's principle into the formula that describes the behavior of ideal gases:
PV = nRT.
51
Indicator
A chemical species that changes color when undergoing dissociation. Indicators are used to signal the end point of a titration.
52
Ion
A single particle or polyatomic species with an electric charge.
53
Ionization Energy
The amount of energy required to remove an electron from orbit around a gaseous atom into free space. Increases from left to right and from bottom to top on the Periodic Table.
54
Isobaric Process
A process that occurs at a constant pressure.
55
Isoelectronic
Two different elements that share the same electronic configuration (such as K+ and Ar).
56
Isolated System
A system that can exchange neither energy nor matter with its surroundings.
57
Isothermal Process
A process in which temperature (and, therefore, internal energy) remains constant.
58
Isovolumetric (Isochoric) Process
A process in which volume remains constant and no net pressure-volume work is done.
59
Kinetic Molecular Theory of Gases
A series of assumptions used to account for the behavior of ideal gases. The theory describes gases as volumeless particles in constant, random motion that exhibit no intermolecular attractions and undergo completely elastic collisions with each other and the walls of their container.
60
Law of Conservation of Energy
Law stating that energy cannot be created or destroyed but only transferred and transformed.
61
Le Chatelier's Principle
States that when a system in equilibrium is places under one of several stressors, it will react in order to regain equilibrium.
62
Lewis Definition
A definition of acids as electron pair acceptors and bases as electron pair donors.
63
Limiting Reagent
The reactant in a chemical equation that, given nonstoichiometric amounts, determines the amount of product that can form; the reactant that runs out first.
64
Metals
Elements that are characteristically electropositive, malleable, and ductile. These elements tend to be found on the left side of the Periodic Table, are lustrous, and have relatively low ionization energies and electron affinities.
65
Molar Mass
The sum of all the masses present (in amu) in one molecule of a molecular compound.
66
Molar Solubility
The molarity of a solute such that a solution is saturated (at equilibrium).
67
Molarity
Concentration of a solution calculated by moles of solute per liters of solution, often denoted by M.
68
Molecular Orbital
The region in a molecule where atomic orbitals overlap, resulting in either a stable low-energy bonding orbital or an unstable high-energy antibonding orbital.
69
Net Ionic Equation
A representation of a displacement reaction showing only the reactive species and omitting the spectator ions.
70
Noble Gas
Inert elements naturally existing in a gaseous state that comprise Group VIII (Group 18) of the Periodic Table.
71
Nonmetals
Elements that have characteristically high electronegativity, ionization energy, and electron affinity. These elements tend to be found on the right side of the Periodic Table and are poor conductors of electricity.
72
Nonpolar Covalent Bonds
A type of covalent bond between atoms with the same electronegativities, resulting in an even distribution of electron density along the bond.
73
Octet Rule
A rule stating that atoms tend to react in order to from a complete octet of valence electrons. Exceptions include hydrogen (stable with two electrons), Be (4), B (6); any element in Period 3 and lower (can have more than 8 and be stable); and molecules with an odd number of electrons
74
Open System
A system that allows for the exchange of energy and matter across its boundaries
75
Oxidation
A reaction in which a species loses electrons.
76
Oxidation-Reduction Half-Reaction
A hypothetical equation showing only the species that is oxidized or reduced in a oxidization-reduction reaction and the correct number of electrons transferred between the species in the complete, balanced equation.
77
Oxidizing Agent
A species that is reduced in the process of oxidizing another species.
78
Paramagnetic
An atom or a substance that contains unpaired electrons and is consequently attracted to a magnet
79
Partial Pressure
The pressure contribution of a single gas in a container holding a mixture of gases, as given by: PA = XA(Ptotal), where XA is the mole fraction of gas A and Ptotal is the total pressure of the mixture.
80
Pauli Exclusion Principle
States that no two electrons in an atom can have the same set of four quantum number values
81
Percent Composition
The percentages of the elements making up a compound (usually by mass in amu)
82
Percent Yield
A ratio (%) of the actual mass of product yielded to the theoretical yield of product mass.
83
pH
Scaled value used to measure the acidic strength of a solution, calculated by taking the negative logarithm of the proton concentration in a solution:
pH = -log[H+]
84
Phase Diagram
A pressure-temperature plot showing the conditions under which a substance exists in pure phase or in equilibrium between different phases
85
Photon
A particle of light energy with a value equal to hf, where h is Planck's constant and f is the frequency of radiation.
86
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 partially positive and partially negative poles.
87
Rate-Determining Step
The slowest step in a reaction mechanism; determines the overall rate of the reaction.
88
Rate Law
An experimentally determined mathematical expression showing the rate of a reaction as a function of the concentrations of its reactants.
89
Reaction Mechanism
A "play-by-play' showing the individual steps of a reaction, including the formation and destruction of any reaction intermediates that may occur.
90
Reaction Order
The sum of the exponents in a rate law, where each exponent provides the reaction order with respect to its reactant.
91
Reaction Quotient
The ratio of the concentrations of the products to the concentrations of the reactants at any point during a reaction, where each reactant and product in the expression is raised to the power of its stoichiometric coefficient. Commonly denoted by Q. Usually compared to the equilibrium constant to determine the direction a reaction will proceed to regain equilibrium.
92
Reaction Rate
The measure of how quickly reactants are consumed and products are formed, commonly expressed in terms of mol L^-1s^-1.
93
Reducing Agent
A species that is oxidized in the process of reducing another species.
94
Reduction
A reaction in which a species gains electrons.
95
Reduction Potential
A measure of the tendency of a species to be reduced, commonly used in identifying the anode and cathode of an electrochemical cell.
96
Resonance Structures
Alternate Lewis structures of the same molecule that show the delocalization of electrons within that molecule; Lewis structures that contribute to a resonance-stabilized system. Resonance structures have the same atomic connectivity but differ in the distribution of electrons.
97
Reversible Reaction
A process that will proceed bidirectionally to form both product and reactant.
98
Solubility
A ratio that measures how much solute can dissolve in a solvent at a given temperature.
99
Solute
A compound, commonly an ion, dissolved in a solvent to create a solution.
100
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.
101
Solvation
The process of forming a cagelike network of solvent molecules around a solute in a solution.
102
Solvent
A medium, commonly a liquid, into which a solute is dissolved to create a solution.
103
Spontaneous Reaction
A reaction that will proceed or occur on its own without additional energy input from its surroundings.
104
Standard Temperature and Pressure (STP)
273K (0C) and 1 atm (760mmHg)
105
System
The part of the universe under consideration that is separated by some real or imaginary boundary from its surroundings.
106
Theoretical Yield
The expected amount of product yielded in a reaction according to the reactant's stoichiometry.
107
Titrand
A solution of unknown concentration to which small volumes of a solution of known concentration are added to reach the equivalence point.
108
Titrant
A solution of known concentration added in small volumes to a solution of unknown concentration to reach the equivalence point.
109
Titration
An analytical procedure in which a solution of known concentration is slowly added to a solution of unknown concentration to the point of molar equivalency, thereby providing the concentration of the unknown solution
110
Transition Elements
The elements found in the B groups of the Periodic Table. These elements contain partially filled d subshells.
111
Transition State
A high-energy complex in which old bonds are partially broken and new bonds are partially formed. Charges existing only in the transition state are designated as partial charges.
112
Triple Point
A point on a phase diagram at which a substance exists in equilibrium between all three phases.
113
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.
114
VSEPR Theory
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
115
Water Dissociation Constant
An expression of the autoionization of water into H+ and OH- at a certain temperature, given by the product of the ions' molar concentrations. Denoted by Kw and equal to 10^-14 at 25C. Kw = [H+][OH-]
116
Acid Dissociation Constant
An equilibrium expression used to measure acid strength, given by the ratio of the product of the products' molar concentrations to the product of the reactants' molar concentration, with each term raised to the power of its stoichiometric coefficient. Denoted Ka.
117
Adiabatic Process
A process in which no heat is transferred to or from the system by its surroundings.
118
Arrhenius Definition
A definition of acids as producers of excess H+ and bases as producers of excess OH- in aqueous solutions.
119
Closed System
A system that allows for the exchange of energy, but not matter, across its boundaries.
120
Combination Reaction
A reaction in which two or more reactants combine to form a product (such as A + B => C).
121
Common Ion Effect
States that the molar solubility of one salt is reduced when another salt, having a common ion, is brought into the same solution
122
Constant-Volume Calorimeter
An apparatus commonly referred to as a bomb calorimeter; used to measure that amount of heat absorbed or released during a reaction.
123
Disproportionation (Dismutation)
An oxidation-reduction reaction in which the same species is both oxidation and reduced.
124
Effective Nuclear Charge
The resulting positive nuclear charge an outer electron senses after accounting for the shielding effect of inner core electrons. Abbreviated Zeff. Increases from left to right from bottom to top on the Periodic Table.
125
Effusion
The movement of gas from one compartment to another through a small opening under pressure.
126
Equilibrium Constant
The ratio of the concentrations of the products to the concentrations of the reactants at the point of equilibrium, where each reactant and product in the expression is raised to the power of its stoichiometric coefficient. Commonly denoted by Keq.
127
Faraday's Constant
Denoted by F. It equals
9.65 x 10^4 C/mol.
Commonly used in the formula:
It = nF, where I is current, t is time (in seconds), and n is moles of electrons transferred.
128
Graham's Law
Law stating that the rates at which two different gases effuse or diffuse are inversely proportional to the square root of their molar masses:
r1 = Square root of M2
--- ------
r2 M1
129
Heisenberg Uncertainty Principle
The quantum mechanical concept 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
130
Ideal Gas
A hypothetical gas containing particles with zero volume and with no attractive intermolecular forces.
131
Ion Product
The product of the molar concentrations of dissociated ions in solution at any point in a dissociation reaction, where each ion is raised to the power of its stoichiometric coefficient. Denoted IP.
132
Metalloids
Elements that have properties between those of metals and nonmetals; includes B, Si, Ge, As, Sb, Te, Po, and At.
133
Molality
Concentration of a solution calculated by moles of solute per kilograms of solvent, often denoted by m.
134
Nernst Equation
An equation used to determine a cell's electromotive force (emf) when conditions are not standard.
Ecell = Eocell - (0.0592/n)logQ where Ecell is the emf, Eocell is the emf under standard conditions, n is the number of electrons transferred in the oxidation-reduction reaction, and Q is the reaction quotient.
135
Neutralization Reaction
A reaction in which an acid and a base are combined to form a salt (and often water as well).
136
Normality
Concentration of a solution calculated by gram equivalent weight of solute per liter of solution, often denoted by N.
137
Process Function
A variable that depends on the path taken to get from one state to another. Includes work and heat.
138
Raoult's Law
States that the vapor pressure of a solvent is proportional to the mole fraction of the solvent in the solution; provides an explanation for the boiling point elevation seen in solutions.
139
Single-Displacement Reaction
A chemical reaction in which an atom or ion of one compound is replaced by another atom or ion (such as
A + BC => B + AC).
140
Solubility Product Constant
The product of the molar concentrations of dissociated ions in a saturated solution, where each ion is raised to the power of its stoichiometric coefficient.
Denoted Ksp.
141
Specific Heat
The amount of heat required to raise one gram of a substance by one degree Celsius.
142
State Function
A variable that depends only in the current state of a system, and not the path taken to get there. Includes pressure, density, temperature, volume, enthalpy, internal energy, Gibbs free energy, and entropy.
143
Strong Acid
An acid that will completely dissociate in aqueous solution (such as HCl, HI, HNO3, and HClO4).
