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OAT General Chem
Chemical Bonding, Stoichiometry, Solutions, and Reaction types
| Term | Definition |
|---|---|
| intramolecular bond | strong chemical bonds, that hold atoms together as molecules include ionic and covalent bonds |
| intermolecular forces | weaker bonds considerable importance in understanding physical properties of many substances - stronger forces hold molecules together more tightly more energy is required to weaken these bonds |
| octet rule | an atom tends to bond with other atoms until it has eight electrons in its outermost shell - exceptions: Hydrogen and Helium |
| ionic bonding | one or more electrons from an atom with a smaller ionization energy are transferred to an atom with a greater electron affinity, the ions are held together by electrostatic forces |
| covalent bonding | an electron pair is shared between two atoms |
| polar covalent bonds | - occurs between atoms with small differences in electronegativity generally in the range 0.4 to 1.7 - the bonding pair is not shared equally but is pulled more toward the element with the higher electronegativity - partial negative/positive |
| ionic bonds | - transfer of electrons from less electronegative to more electronegative - form crystal lattices - high melting/boiling points due to strong electrostatic forces - conduct electricity in liquid and aqueous states |
| covalent bonds | - two or more atoms with similar electronegativites interact - atoms achieve a noble gas configuration by sharing - generally contain discrete molecular units with weak intermolecular forces - low melting/boiling points - not conductive |
| bond order | the number of shared electron pairs between two atoms; - a single bond has a bond order of one - a double bond has a bond order of two - a triple bond has a bond order of three |
| bond length | the average distance between the two nuclei of the atoms involved in the bond - as the number of shared electrons pairs increase = the decrease in bond length - triple bond shorter than double, double is shorter than single |
| bond energy | the energy required to separate two bonded atoms - the strength of a bond increases as the number of shared electron pairs increases |
| bonding electrons | shared valence electrons of a covalent bond |
| nonbonding electrons | valence electrons not involved in the covalent bond |
| lone electron pairs | unshared electron pairs |
| lewis structure | represent the bonding and nonbonding electrons in a molecule |
| formal charge | formal charge = valence - 1/2 bonded e- - nonbonded e- OR formal charge = valence - (# of sticks + # of dots) |
| resonance structures | one or two more Lewis structure for a single molecule unable to be described fully with only one Lewis structure |
| exceptions to the octet rule | atoms found in or beyond the third period can have more than eight valence electrons, these atoms can be assigned more than four bond in Lewis structure |
| polar molecule | a molecule that has such separation of positive and negative charges |
| dipole moment | a measure of the polarity of a molecule - mu (u) = qr q = charge magnitude, r = distance between the two partial charges |
| nonpolar covalent bond | - occurs between atoms that have the same electronegativites - the bonding electron pair is shared equally such there is no separation if charge - occurs in diatomic molecules such as H2, Cl2, O2, and N2 |
| coordinate covalent bond | - the chared electron pair comes from the lone pair of one of the atoms in the molecule - typically found in Lewis acid-base compounds |
| lewis acid | a compound that can accept an electron pair to form a covalent bond |
| lewis base | a compound that can donate an electron pair to form a covalent bond |
| valence shell electron-pair repulsion theory | states that the 3-D arrangement of atoms surrounding a central atom is determined by the repulsion between the bonding and the nonbonding electron pairs in the valence shell |
| steps used to predict the geometric structure of a molecules using the VSEPR theory | 1. Draw the lewis structure 2. Count the total number of bonding and nonbonding electron pairs in the valence shell of the atom 3. arrange the electron pairs around the atom 4. determine the bond angle |
| hydrogen bonding | - specific particularly strong form of dipole-dipole interaction - when hydrogen is bound to either fluorine, oxygen, or nitrogen - hydrogen carries little electron density |
| dipole-dipole bonding | - energetically favorable because an attractive dipole force is formed - present in the solid and liquid phases - polar speices tend to have higher boiling points than nonpolar |
| London dispersion forces | - unequal sharing of electrons, causing rapid polarization and counterpolarization of the electron cloud and formation of short-lived dipoles - generally weaker - large molecules = greater dispersion - low temperatures noble gases liquefy |
| van der Waals forces | forces not due to the interactions of ions or hydrogen bonding |
| order of decreasing intermolecular force strength | dipole-ion > hydrogen bonding > dipole-dipole > LDF |
| carbon-carbon bonding | based on length and energy levels as well as hybridization - C-C longest length; lowest bond energy - C=C middle length; middle bond energy - C_=C shortest length; highest bond energy |
| mole | weight of sample (g)/molar weight (g/mol) |
| gram equivalent weight | molar mass / # of equivalents per mole (the # of hydrogen ions) |
| equivalents | weight of compound / gram equivalent weight |
| law of constant composition | all samples of a given compound will contain the same elements in identical mass ratios ex: every sample of H2O will contain 2 atoms of hydrogen and 1 atom of oxygen and 1 gram of hydrogen for every 8 g of oxygen |
| empirical formula | the simplest whole number ratio of the elements in the compund |
| molecular formula | gives the exact number of atoms of each element in the compound |
| percent composition | (mass of element in formula/ formula weight of compound ) x 100 |
| limiting reagent | mass of element in sample x mole/molecular weight |
| theoretical yield | the amount of product that can be predicted from a balanced equation |
| actual yield | the amount of experimental product |
| percent yield | (actual yield / theoretical yield) x 100 |
| solvation | interaction between solute and solvent molecules - possible when attractive forces between solute and solvent are stronger than between solvent particle - nonionic solutes involves van der Waal forces - like solute dissolve in like solvents |
| solubility | - maximum amount of that solute that can be dissolved in a particular solvent at a particular temp - the solubility of liquids or solids increases with high temp - solubility of gas will increase with low temp and high pressure |
| precipitate | if more solute is added the excess solute will come out and form a collection at the bottom of the container |
| crystallization | when a dissolved solute comes out of solution and forms crystals |
| supersaturated | solutions that contain more solute than found in saturated solutions - formed by manipulating temperature or pressure - addition of more solute will cause excess solute in the solution to separate |
| percent composition by mass | (solute mass / mass of the solution) x 100 |
| mole fraction (X) | (# of moles / total # of moles in solution) |
| molarity | # of moles / liter of solution |
| molality | # of moles solute / kg of solvent |
| normality | molarity x (# of equivalent weight of solute/ liter of solution) |
| dilution | MiVi = MfVf |
| soluble salts in aqueous solutions (with exceptions) | - all salts of alkali metal ions (Li+, Na+, K+, Rb+, etc) - all salts of ammonium ion (NH4+) - all salts with chloride, bromide, and iodide (except Ag, Pb, Hg2) - all salts of sulfate ion (SO4) (except Ca2+, Sr2+, Ba2+, Pb2+) |
| insoluble salts in aqueous (with exceptions) | - all metal oxides (exception of alkali metal oxides, CaO, SrO, and BaO) - all hydroxides containing OH- (except alkali metal hydroxide, Ca(OH)2, Sr(OH)2,and Ba(OH)2 - all salts with carbonates (CO3), phosphates (PO4), sulfides (S2-), and sulfites (SO3) |
| electrolytes | electrical conductivity of aqueous solutions is governed by presence and concentration of ions in solution |
| strong electrolytes | - a solute dissociates completely into its constituent ions ex: ionic compounds (NaCl and KI) - molecular compounds with highly polar covalent bonds that disspcoate into ions ex: HCl in water |
| weak electrolytes | ionizes or hydrolyzes incompletely in aqueous solution and only some of the solute is present in ionic form ex: acetic acid, weak acids, ammonia, weak bases |
| nonelectrolytes | compounds do not ionize at all in aqueous solution retaining their molecular structure in solution which limits their solubility ex: many nonpolar gases and organic compounds |
| combination reactions | reactions in which two or more reactants form one product S (s) + O2 (g) ---> SO2 (g) |
| decomposition reactions | a compound breaks down into two or more substances as a result of heating or electrolysis *(^) = heat* 2 HgO (s) --^--> 2 Hg (l) + O2 (g) |
| electrolysis | a specific process that causes the decomposition of a compound by passing an electric current through the reactant |
| single-displacement reactions | occur when an atom or ion of one compound is replaced by an atom of another element Zn (s) + CuSO4 (aq) ---> Cu (s) + ZnSO4 (aq) |
| net ionic equations | net equations written in ionic form Zn (s) + Cu2+ (aq) + SO4 2- (aq) ----> Cu (s) + Zn2+ (aq) + SO42- (aq) |
| spectator ions | do not take part in the overall reaction but simply remain in solution throughout ex: SO4 |
| net ionic reaction | showing only species that actually participate in the reaction Zn (s) + Cu2+ (aq) ----> Cu (s) + Zn2+ (aq) |
| double-displacement reactions | elements from two different compounds displace each other to form two new compounds CaCl2 (aq) + 2 AgNO3 (aq) ----> Ca(NO3)2 (aq) + 2 AgCl (s) |
| neutralization reactions | specific type of double displacement that occurs when an acid reacts with a base to produce a solution of a salt and water HCl (aq) + NaOH (aq) ----> NaCl (aq) + H2O (l) |
Created by:
Jalisa.bland
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