Gen Chem Review

Front 1

Nomenclature

Back 1

Binary, Polyatomic, Stock System

Front 2

Binary Ionic Compounds

Back 2

Metal + Non-metal Cation first, followed by anionMetal cations are name of elementNon-metal anions names are changed to "...ide"

Front 3

Binary Covalent Compounds

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Non-metal + Non-Metal* Add "...ide" to the root of the least metallic element.* use numeric prefixesEx) NO2 = nitrogen dioxide

Front 4

polyatomic ions

Back 4

acetate, ammonium, carbonate, chlorate, chlorite, chromate, cyanide, dichromate, bicarbonate, bisulfate, bisulfite, hydroxide, hypochlorite, nitrate, nitrite, oxalate, perchlorate, permanganate, phosphate, sulfate, sulfite

Front 5

stoichiometry

Back 5

Convert beginning measurement into moles of that elementMultiply moles of what you have by what you needTurn moles of what you need into ending measurement

Front 6

balancing equations

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1. Never touch subscripts2. Treat polyatomic as one unit3. Balance Metals4. Balance non-metalsSave hydrogen and oxygen for lastMake sure coefficients are at their lowest ratio

Front 7

percent yield equation

Back 7

Percent Yield= (Actual Yield / Theoretical Yield) x 100%

Front 8

Ideal gas law

Back 8

PV=NRT

Front 9

Boyles Law

Back 9

P1V1=P2V2

Front 10

Charles Law

Back 10

V1/T1 = V2/T2

Front 11

Combined Gas Law

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P1V1/T1 = P2V2/T2

Front 12

Hess' Law of constant heat summation

Back 12

Regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes.

Front 13

Specific Heat

Back 13

The heat required to raise the temperature of the unit mass of a given substance by a given amount (usually one degree)

Front 14

Enthalpy of a reaction

Back 14

Enthalpy of a reaction or energy change of a reaction DH, is the amount of energy or heat absorbed in a reaction. If the energy is required, DH is positive, and if energy is released, the DH, is negative.

Front 15

Heat of Formation

Back 15

Hreaction = SUM (Hproducts) - SUM (Hreactants)or if you prefer using symbolsHreaction = SHproducts - SHreactants

Front 16

Isotope

Back 16

Two or more forms of the same element that contain equal numbers of protons but different numbers of neutrons in their nuclei

Front 17

# sigma and pi bonds in a compound

Back 17

Sigma bonds are the first bonds, pi bonds are the second and third bonds. They are made from leftover p orbitals.

Front 18

Formal charge

Back 18

The charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms regardless of electronegativity.

Front 19

Hydrogen bond

Back 19

A weak bond between two molecules resulting from an electrostatic attraction between a proton in one molecule and an electronegative atom in the other.

Front 20

Absolute zero

Back 20

The temperature at which all substances have no thermal energy; 0 Kelvin or -273.15 Celsius

Front 21

Absorption series

Back 21

The series of discrete lines at characteristic frequencies representing the energy required to excite an electron from the ground state.

Front 22

Acid

Back 22

A species that donates hydrogen ions or accepts electrons--Lewis Acid.

Front 23

Acid Dissociation Constant (Ka)

Back 23

The equilibrium constant that measures the degree of dissociation of an acid under specific conditions.

Front 24

Acidic solution

Back 24

An aqueous solution that contains more H+ ions than OH- ions; pH<7 under standard conditions.

Front 25

Adiabatic Process

Back 25

A process that occurs without the transfer of heat into or out of the system.

Front 26

Alkali Metals

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Elements found in Group IA of the periodic table; highly reactive, readily losing one valence electron to form ionic compounds with nonmetals.

Front 27

Amphoteric Species

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A species capable of reacting as either as an acid or base, depending on the nature of the reactants.

Front 28

Amphiprotic Species

Back 28

A species that may either gain or lose a proton.

Front 29

Anion

Back 29

Negatively charged ion

Front 30

Anode

Back 30

The electrode at which oxidation occurs.

Front 31

Antibonding Orbital

Back 31

A molecular orbital formed by the overlap of two or more atomic orbitals; energy is greater than the energy of the combining atomic orbitals.

Front 32

Arrhenius Base

Back 32

A species that donates hydroxide ions (OH-) in aqueous solution.

Front 33

Arrhenius Acid

Back 33

A species that donates protons (H+) in an aqueous solution.

Front 34

Arrhenius Equation

Back 34

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.

Front 35

Atomic Mass Unit

Back 35

A unit of mass defined as 1/12 the mass of a carbon-12 atom; approximately equal to the mass of one proton or one neutron.

Front 36

Atomic Orbital

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Describes the region of space where there is a high probability of finding an electron.

Front 37

Atomic Radius

Back 37

The average distance between a nucleus and its outermost electron; usually measured as one-half the distance between two nuclei of an element in its elemental form.

Front 38

Aufbau Principle

Back 38

The concept that electrons fill energy levels in order of increasing energy, completely filling one sublevel before beginning to fill the next.

Front 39

Autonization

Back 39

The process by which a molecule (usually water) spontaneously dissociates into cations and anions.

Front 40

Bohr Model

Back 40

The model of the hydrogen atom in which electrons assume certain circular orbits around a positive nucleus.

Front 41

Boiling point elevation

Back 41

The amount by which a given quantity of solute raises the boiling point of a liquid; a colligative property. ΔT = i Kb mi=Van't Hoff Factor (degree of dissociation for the solute) for substances that ionize in water (ie separate into ions). Substances which do not ionize have a Van't Hoff Factor of 1. Kb=molal boiling point elevation constant (ebullioscopic constant)m=molality of the solution.

Front 42

bond length

Back 42

The average distance between two nuclei in a bond; as the number of shared electron pairs increases, the bond length decreases. (so single bond>double bond>triple bond in terms of increasing to decreasing bond length).

Front 43

bond order

Back 43

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, and a triple bond has a bonder order of 3.

Front 44

bonding orbital

Back 44

A molecular orbital formed by the overlap of two or more atomic orbitals; energy is less than that of the combining orbitals.

Front 45

Boyle's Law

Back 45

The law stating that at constant temperature, the volume of a gaseous sample is inversely proportional to it's pressure.

Front 46

Buffer

Back 46

A solution containing a weak acid and its salt (or a weak base and its salt) that tends to resist changes in pH.

Front 47

Cathode

Back 47

The electrode at which reduction takes place.

Front 48

Cation

Back 48

Positively charged ion

Front 49

Charle's Law

Back 49

The law stating that the volume of a gas at constant pressure is directly proportional to its absolute (kelvin) temperature.

Front 50

Common Ion Effect

Back 50

A shift in the equilibrium of a solution due to the addition of ions of a species already present in the reaction mixture.

Front 51

Conjugate Acid-Base Pair

Back 51

The relationship between a Bronsted-Lowry acid and its deprotonated form, or a Bronsted-Lowry base and its protonated form.

Front 52

Diffusion

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The random motion of gas or solute particles across a concentration gradient, leading to uniform distribution of the gas or solute throughout the container.

Front 53

Dipole-Dipole Interactions

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The attractive forces between two dipoles; magnitude is dependent on both the dipole moments and the distance between the two species.

Front 54

Dipole

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A species containing bonds between elements of different electronegativities, resulting in an unequal distribution of charge.

Front 55

Dipole Moment

Back 55

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.

Front 56

Dissociation

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The separation of a single species into two separate species; usually used in reference to salts or weak acids or bases.

Front 57

Double-Displacement Reaction

Back 57

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.

Front 58

Electrolysis

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The process in which an electrical current is used to power an otherwise nonspontaneous decomposition reaction.

Front 59

Electronegativity

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A measure of the ability of an atom to attract the electrons in a bond; commonly measured with the Pauling scale.

Front 60

Empirical Formula

Back 60

The simplest whole-number ratio of the different elements in a compound.

Front 61

Enthalpy

Back 61

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.

Front 62

Entropy

Back 62

A property related to dispersion of energy through a system or degree of disorder in that system; the change in entropy (delta S) in the course of a reaction is the difference between the entropies of the products and the reactants.

Front 63

Equilibrium

Back 63

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.

Front 64

Exothermic Reaction

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A reaction that gives off heat to the surroundings (negative delta H, delta H<0) as the reaction proceeds.

Front 65

First order reaction

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A reaction in which the rate is directly proportional to the concentration of only one reactant.

Front 66

Gibbs Free Energy

Back 66

The energy of a system available to do work. The change 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.

Front 67

Metalloids

Back 67

-also called semimetals-EN and ionization energy inbetween metals and nonmetals

Front 68

IMF and relative strength

Back 68

Weakest: dispersal or London or Van der Waals forces-temporary dipoles (the reason noble gases can liquify)Intermediate: dipole-dipole interactions-more permanent dipoles (not relevant with gases)Strongest: hydrogen bonding (10% of covalent bond)

Front 69

zero order reaction

Back 69

temp matters, concentration doesnt matter

A (T) vs Time

Front 70

1st order reation

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rate is proportional to reactant

changes with concentration

Concentration VS Time

Front 71

2nd order reaction

Back 71

depends on collision between two reactant molecules

1/A VS Time

Front 72

Hess' Law

Back 72

enthalpy changes of a reaction are additive

Front 73

Solubility

Back 73

Ksp

The solubility of non-gaseous solvents increases with temp, while that of gaseous temps decreases with increased temp

Front 74

latent heat

Back 74

heat absorbed or radiated during a change of phase at a constant temperature and pressure

Front 75

conduction

Back 75

Heat transfer by direct contact. Requires things to touch

Front 76

convection

Back 76

Heat transfer by flowing current. Need the physical flow of matter

Front 77

radiation

Back 77

Heat transfer by electromagnetic radiation. Does not need the physical flow of matter, can occur through vacuum