Chemistry Study Guide- FINAL
Chemistry Study Guide- FINAL ch 1213
Popular in Chemistry 1
Popular in Chemistry
This 21 page Study Guide was uploaded by Madison Greer on Thursday April 21, 2016. The Study Guide belongs to ch 1213 at Mississippi State University taught by Erin Dornshuld in Summer 2015. Since its upload, it has received 59 views. For similar materials see Chemistry 1 in Chemistry at Mississippi State University.
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Date Created: 04/21/16
• ionization energy (IE)- the energy required to remove an electron from a neutral atom • electron afﬁnity (EA)- the energy change when a neutral atom attracts an electron to become a negative ion • Ionic compounds tend to form between metals and nonmetals when electrons are transferred from an element with a low ionization energy (the metal) to one with a high electron afﬁnity (the nonmetal) • compound- substance composed of two or more elements combined in a speciﬁc ratio and held together by chemical bonds • Maximum stability results when a chemical species is isoelectronic with a noble gas. • Example- Na+= isoelectronic with Ne; Cl-= isoelectronic with Ar • Lewis dot symbol- consists of an element’s symbol with dots. Each dot represents a valence electron • cation- lose electron (positive) • anion- gain electron (negative) • ionic bonding- refers to the electrostatic attraction that holds oppositely charged ions together in an ionic compound • chemical formula- denotes the constituent elements and the ratio in which they combine • lattice- a 3D array of oppositely charged ions • lattice energy- amount of energy required to convert a mole of ionic solid to its constituent ions in the gas phase • monatomic cation- named by adding the word ion to the name of the element • monatomic anion- named by changing the ending of the element’s name to -ide • In cases where a metal cation may have more than one possible charge, the charge is indicated in the name of the ion with a Roman numeral in parenthesis. • Lewis theory of bonding- where electrons are shared, not transferred; covalent bonds • molecule- a combination of at least two atoms in a speciﬁc arrangement held together by chemical forces • Law of deﬁnite proportions- different samples of a given compound always contain the same elements in the same mass ratio • Law of multiple proportions- if two elements can combine with each other to form two or more different compounds, the ratio of masses of one element that combines with the ﬁxed mass of the other element can be expressed in small whole numbers (works for simple compounds) • diatomic molecules- molecules that contain only two atoms • homonuclear- both atoms are the same; H , F , 2i ,2etc2 • heteronuclear- both atoms are different; HF, NaCl, etc… • polyatomic molecules- contain more than two atoms • chemical formula- denotes the composition of the substance • molecular formula- shows the exact number of atoms of each element in a molecule • structural formula- shows not only the elemental composition but also the general arrangements (and sometimes connectivity) • Some elements have two or more distinct forms known as allotropes. Example- oxygen (O 2 and ozone (O ) 3re allotropes of oxygen. empirical formula- written by reducing the subscripts to the smallest possible whole numbers • • binary molecular compounds- compounds with two elements • There are metals that will form covalent bonds rather than ionic. MEMORIZE B2H6 diborane SiH4 silane NH 3 ammonia MEMORIZE PH 3 phosphane H 2 water H 2 hydrogen sulﬁde • acid- a substance the produces hydrogen ions when dissolved in water • inorganic compounds- generally deﬁned as compounds that do not contain carbon • organic molecules- molecules that contain carbon and hydrogen • hydrocarbons- the simplest organic compounds that consist only of carbon and hydrogen MEMORIZE CH 4 methane C2H 6 ethane C3H 8 propane C 4 10 butane C 5 12 pentane C 6 14 hexane C 7 16 heptane C 8 18 octane C 9 20 nonane C10 22 decane • For alkanes, the names end in -ane. • functional group- groups that determine many chemical properties of a compound because that is typically where a chemical reaction occurs • polyatomic ions- a charged chemical species compound of either two or more atoms that are covalently bonded or a metal complex that can be considered to be acting as a single unit • MEMORIZE *look at notes for list of polyatomic ions* • oxoanions- polyatomic ions that contain one or more oxygen atoms and one atom of another element; ex- chlorate, nitrate, sulfate • oxoacids- produce hydrogen ions and the corresponding oxyanions when dissolved in water; can be monoprotic or polyprotic • Oxyanions whose name ends in -ate, can be named as follows: • An acid based on an -ate ion is called -ic acid; ex- HClO i3 called chloric acid • An acid based on an -ite ion is called -ous acid; ex- HClO is2called chlorous acid • Preﬁxes in oxyanion names are retained in the names of the corresponding oxoacids; ex- HClO an4 HClO are called perchloric acid and hypochlorous acid • hydrate- a compound that has a speciﬁc number of water molecules within its solid structure • anhydrous- the compound no longer has water molecules associated with it; driven off all the water from the hydrate MEMORIZE CO 2 dry ice NaCl salt N2O nitrous oxide, laughing gas CaCO 3 marble, chalk, limestone NaHCO 3 baking soda MgSO 4 *H 2 Epsom salt Mg(OH) 2 milk of magnesia • molecular mass (molecular weight)- the mass in atomic mass units (amu) of an individual molecule. Because the atomic masses on the periodic table are averaged masses, the molecular mass is also and averaged value. • percent composition by mass- a list of percent-by-mass of each element in a compound • molecule= covalent bonds, molecular mass • ionic compound= ionic bonds, formula mass • molar mass- the mass in grams of one mole of a substance • *know how to do problems like Example 5.16 in notes* • octet rule- atoms will lose, gain, or share (valence) electrons to achieve a noble gas conﬁguration • Lewis structure- a representation of covalent bonding in molecules • single bond- 2 electrons, one electron pair • double bond- 4 electrons, 2 electron pairs • triple bond- 6 electrons, 3 electron pairs • bond length- the distance between two nuclei of two covalently bonded atoms in a molecule • shortest longest weakest strongest triple < double < single single < double < triple • pure covalent bond- neutral atoms held together by equally shared electrons • polar covalent bond- partially held together by unequally shared electrons • ionic bond- oppositely charged ions held together by electrostatic attraction • electronegativity- the ability of an atom to draw shared electrons toward itself • Six Steps for Drawing Lewis Structures for Molecules and Polyatomic Ions 1. Draw the skeletal structure of the compound. 2. Count the total number of valence electrons present; add electrons for negative charges and subtract electrons for positive charges. 3. For each bond in the skeletal structure, subtract two electrons from the total valence electrons. 4. Use the remaining electrons to complete octets of the terminal atoms by placing pairs of electrons on each atom. 5. Place any remaining electrons in pairs on the central atom. 6. If the central atom has fewer than 8 electrons, move one or more pairs from the terminal atoms to form multiple bonds between the central atom and the terminal atoms. • formal charge- used to determine the most plausible Lewis structure • formal charge “equation”- (# of valence electrons for that atom)-(# of bonds to that atom)-(# of dots around that atom)= formal charge • resonance structure- one of two or more equally valid Lewis structures for a single species that cannot be represented accurately with a single Lewis structure • Resonance structures are NOT rapidly changing structure; they are always equal. • expanded octet- the central atom has more than 8 electrons • free radicals- molecules with an odd number of electrons • coordinate covalent or dative bond- a bond created from a pair of electrons donated from one atom • Lewis acid- can accept a pair of electrons • Lewis base- can donate a pair of electrons • VSEPR (valence-shell electron pair repulsion) theory- assumes that each atom in a molecule will achieve a geometry that minimizes the repulsion between electrons in the valence shell of that atom • 180 degrees between 2 electron domains; linear • 120 degrees between 3 electron domains; trigonal planar • 109.5 degrees between 4 electron domains; tetrahedral • trigonal pyramidal- 5 electron domains • octahedral- 6 electron domains electron domain geometry- the arrangement of electron domains around the central atom • • molecular geometry- the arrangement of bonded atoms (ignore all lone electron pairs) • bond angle- the angle between two adjacent A-B bonds KNOW THIS!!! total number number of type of electron- molecular example of electron lone pairs molecule domain geometry domains geometry 3 1 AB trigonal planar bent SO 2 2 4 1 AB 3 tetrahedral trigonal NH 3 pyramidal 4 2 AB 2 tetrahedral bent H2O 5 1 AB 4 trigonal seesaw- SF4 bipyramidal shaped 5 2 AB 3 trigonal T-shaped ClF 3 bipyramidal 5 3 AB 2 trigonal linear IF2 bipyramidal 6 1 AB 5 octahedral square BrF 5 pyramidal 6 2 AB 4 octahedral square planar XeF 4 • Loan pairs take up more space than bonds. • VSEPR (valence-shell- electron pair repulsion)- assumes that each atom in a molecule will achieve a geometry that minimizes the repulsion between electrons in the valence shell of that atom • electron domains • lone electron pairs • single bonds • double bonds • triple bonds • 2 electron domains= linear= 180 degrees • 3 electron domains= trigonal planar= 120 degrees • 4 electron domains= tetrahedral= 109.5 degrees • 5 electron domains= trigonal bipyramidal • 6 electron domains= octahedral • electron domain geometry- the arrangement of electron domains around the central atom • molecular geometry- the arrangement of bonded atoms (ignore all lone electron pairs) • bonds angle- the angle between two adjacent A-B bonds • Electron Domain and Molecular Geometries and Molecules with Lone Pairs on the Central Atom total number number of type of electron- molecular of electron domain example domains loan pairs molecule geometry geometry 3 1 AB trigonal planar bent SO 2 2 4 1 AB 3 tetrahedral trigonal NH 3 pyramidal 4 2 AB 2 tetrahedral bent H2O trigonal seesaw- 5 1 AB 4 bipyramidal shaped SF 4 trigonal 5 2 AB 3 T-shaped ClF3 bipyramidal 5 3 AB 2 trigonal linear IF2- bipyramidal square 6 1 AB 5 octahedral pyramidal BrF5 6 2 AB 4 octahedral square planar XeF 4 • Strong Weak Ion-Ion Ion-Dipole Hydrogen Bond Dipole-Dipole Dipole-Induced Dipole Dispersion • polar bond- atoms bonded together are different • non polar bond- atoms held together are the same • structural isomers- molecules with the same molecular formula but bonded in either a different way or arranged in a different way • trans- different sides • cis- same sides; boiling point is higher • intermolecular forces- attractive forces between neighboring molecules; collectively known as van der Waals forces • dipole-dipole interactions- attractive forces that act between polar molecules • hydrogen bond- an important type of dipole-dipole interaction; occurs between molecules that contain H bonded to a small, highly electronegative atom such as N, O, or F • dispersion forces- also known as London dispersion forces; attractive forces that arise from temporary dipoles • ion-dipole interactions- the Coulombic attraction between ions (positive or negative) and polar molecules • induced dipole- a dipole is induced by a neighboring charge • valence bond theory- atoms share electrons when an atomic orbital in one atom overlaps with an atomic orbital of another • Energy always wants to ﬂow downhill. • p-orbitals=90 degrees • hybridization- the mixing of atomic orbitals Number of Electron Domains and Hybrid Orbitals on a Central Atom number of electron domains on 2 3 4 5 6 central atom hybrid orbitals sp sp 2 sp3 sp d sp d2 • sigma (σ) bond- all single bonds are sigma bonds; double and triple bonds each contain a single sigma bond • pi (π) bond- single bonds contain no pi bonds; double bonds contain a single pi bond; a triple bond contains two pi bonds • molecular orbital theory- another bonding theory which is based on the premise that atomic orbitals combine to form molecular orbitals that are a ‘property’ of the entire molecule • bonding molecular orbital- constructive combination of the two 1 s orbitals gives rise to a molecular orbital that lies along the intermolecular axis • anti bonding molecular orbital- destructive combination of the two 1 s orbitals gives rise to a molecular orbital that lies along the intermolecular axis, but does not lie between the two nuclei • bond order- the higher the bond order, the more stable the molecule is chemical reaction- a process that neither creates not destroys atoms, but rearranges them in chemical compounds • chemical equation- uses chemical symbols to denote what occurs in a chemical reaction • reactant- each chemical species that appears to the left of the arrow • product- each species that appears to the right of the arrow • aqueous- dissolved in water • stoichiometric coefﬁcients- used to achieve a balanced chemical equation • KNOW HOW TO BALANCE CHEMICAL EQUATIONS • The Greek letter delta denotes that heat has been added to the reaction. • combination reaction- two or more reactants combine to form a single product • decomposition reaction- two or more products form from a single reactant • double replacement reaction- two compounds react and the positive ions and the negative ions of the two reactants switch places to form two new products • acid/base reaction- an acid and a base produce water and a salt • combustion reaction- a substance hat burns in the presence of oxygen; products will ALWAYS be carbon dioxide and water • combustion analysis- the experimental determination of an empirical formula can be carried out using this method • limiting reagent— the reactant that is completely consumed • excess reagents- reactants that are still present after the limiting reagent is consumed • theoretical yield- using stoichiometry to determine the amount of product formed in a reaction • actual yield- the product that was actually obtained • percent yield- the ratio of theoretical yield and actual yield multiplied by 100% • KNOW HOW TO FIND LIMITING REAGENT AND EXCESS REAGENT • KNOW HOW TO CALCULATE THEORETICAL YIELD AND PERCENT YIELD • solution- a homogenous mixture of two or more substances; saltwater • solvent- the substance present in the largest amount; water • solute- the substance in a solution that is not the solvent; salt • electrolyte- a substance that dissolves in water to produce ions; conducts electricity • non-electrolyte- a substance that dissolves in water that does not produce ions; does NOT conduct electricity • dissociation- the process by which an ionic compound breaks apart into its constituent ions • ionization- the process by which a molecular compound forms ions when it dissolves • Acids and bases are molecules that ionize in water. • acid- a substance that produces H+ ions when dissolved in water • base- a substance that produces OH- ions when dissolved in water • strong electrolyte- a compound that dissociates completely in water • !!!MEMORIZE!!! Strong Acids Strong Bases HCl LiOH HBr NaOH HI KOH HNO 3 RbOH HClO 3 CsOH HClO 4 Ca(OH) 2 H 2O 4 Sr(OH) 2 Ba(OH) 2 • weak electrolyte- compound that does NOT completely dissociate in water Strong acids and strong bases are strong electrolytes. • • A double arrow in a chemical equation means that the dissolved parts are able to “bounce” back and forth to create the reactant even though it is dissolved in water; NOT completely dissociated/dissolved in water; also known as dynamic chemical equilibrium. • ionic compound- metal plus a nonmetal • precipitate- an insoluble solid product that separates from a solution • precipitation reaction- corresponding chemical reaction that forms a precipitate • hydration- when an ionic substance dissociates in water, the water molecules remove the individual ions from the 3D crystal lattice and surround them • insoluble- very slightly soluble • !!!MEMORIZE!!! Water- soluble Compounds Insoluble Exceptions Compounds containing an alkali metal (ﬁrst column on periodic table) cation (Li+, Na+, K+, Rb+, Cs+) or the ammonium ion (NH )4+ Water- soluble Compounds Insoluble Exceptions Compounds containing the nitrate ion (NO ), 3- acetate ion (C H O ), or chlorate ion (ClO ) - 2 3 2 3 Compounds containing the chloride ion (Cl-), Compounds containing Ag+, Hg 22, or Pb 2+ bromide ion (Br-), or iodide ion (I-) Compounds containing Ag+, Hg 22, Pb , Ca ,2+ Compounds containing the sulfate ion (SO ) 42- 2+ 2+ Sr , or Ba • !!!MEMORIZE!!! Water-insoluble Compounds Soluble Exceptions Compounds containing the carbonate ion (CO ), 32- phosphate ion (PO ), chromate ion (CrO ), or 2- Compounds containing Li+, Na+, K+, Rb+, Cs+, or 4 4 NH 4+ sulﬁde ion (S )- Compounds containing Li+, Na+, K+, Rb+, Cs+, or Compounds containing the hydroxide ion (OH-) Ba 2+ • metathesis=double replacement reaction • molecular equations- equations not explicitly writing out the individual ions • ionic equations- represent any compound that exists completely or predominantly as ions in the solution • net ionic equations- only show the ions directly involved in the reaction • spectator ions- in an ionic equation, ions that are found on both sides of the arrow (equal sign) • Arrhenius acid- a substance that increases H+ concentration when added to water • Arrhenius base- a substance that increases OH- concentration when added to water • Brønsted acid- a proton donor • Brønsted base- a proton acceptor • monoprotic- an acid with one ionizable hydrogen • diprotic- an acid with two ionizable hydrogens • triprotic- an acid with three ionizable hydrogens • Most of the time we call acids with more than one ionizable hydrogen, polyprotic • neutralization reaction- a reaction between an acid and a base the produce water and a salt • salt- an ionic compound • oxidized- lose electrons • reduced- gain electrons • oxidizing agent- the species that accepts the electrons • reducing agent- the species that donates the electrons • oxidation number (oxidation state)- the charge an atom would have if electrons were transferred completely • Elements that show an increase in oxidation number are oxidized. • Elements that show a decrease in oxidation number are reduced. • KNOW HOW TO ASSIGN OXIDATION NUMBERS • single displacement reaction- when an element or ion moves out of one compound and into another • activity series- a list of metals (and hydrogen) arranged from bottom to top in order of increasing ease of oxidation • An element in the activity series will be oxidized by the ions of any element that appears below it. • Elements will NOT be oxidized by elements that appear above it • concentration- the amount of solute dissolved in a given quantity of a solvent or solution • molarity (M)- deﬁned as the number of moles of solute per liter of solution • dilution- make a concentrated solution less concentrated • KNOW HOW TO FIND MOLARITY • KNOW HOW TO FIND MOLARITY FROM A DILUTION • gravimetric analysis- an analytical technique based on the measurements of mass • KNOW HOW TO CALCULATE PERCENT BY MASS • titration- a standard solution is added gradually to another solution of unknown concentration until the chemical reaction between the two solutions is complete • standard solution- a solution of accurately known concentration • equivalence point- the point in the titration at which all the acid (or base) has been neutralized • endpoint- the point at which the color changes • disproportionation reaction- occur when one element undergoes both oxidation and reduction • KNOW HOW TO DO ACID/BASE NEUTRALIZATION TITRATION • When chemical reactions occur, there is a change in energy. • system- the speciﬁc part of the universe that is of interest to us and is usually deﬁned as the substances involved in chemical and physical changes • surroundings- everything that isn't the system • joule- kg m / s 2 • universe= system+surroundings • thermal energy- the energy that comes from heat • heat- the transfer of thermal energy between two bodies that are different temperatures • thermochemistry- the study of the heat (transfer of thermal energy) associated with chemical reactions • endothermic- heat is supplied to the system; energy ﬂows into the system • exothermic- heat is given off by the system; energy ﬂows out of the system • thermodynamics- the scientiﬁc study of the interconversion of heat and other kinds of energy • open system- can exchange mass and energy with its surroundings • closed system- can transfer energy but not mass to the surroundings • isolated system- cannot exchange either mass or energy with the surroundings • state of the system- the values of all relevant macroscopic properties such as composition, energy, temperature, pressure, and volume • state functions- properties that are determined by the state of the system regardless of how the condition was achieved • ΔT- change in temperature • ﬁrst law of thermodynamics- states that energy can be converted from one form to another, but it can never be created or destroyed • internal energy (ΔU)- ΔU= U ﬁnal Uinitial • internal energy has two components • kinetic energy- comes from the motions of particles (electrons, nucleus, molecule, etc.) • potential energy- comes from the repulsive/attractive interactions between the particles • ΔU sys+ ΔU surr 0 • ΔU sys= -ΔU surr • Work is NOT a state function. • q- heat released or absorbed by system • w- work done on the system or done by the system • q- positive for an endothermic process and negative for exothermic process • w- positive for work done on the system and negative for work done by the system • KNOW HOW TO DETERMINE AMOUNT OF WORK DONE • pressure-volume work- work done by a constant-pressure process work= PΔV • • q is NOT a state function but q Iv. • enthalpy- H • enthalpy- H= U+PV • U- internal energy • P- pressure • V- volume • change in enthalpy- ΔH • q = ΔH p • q is NOT a state function but q Ip • enthalpy of reaction- the difference between the enthalpies of the products and the reactants • thermochemical equations- the chemical equations that show the enthalpy changes as well as the mass relationships Guidelines for writing, interpreting, and manipulating thermochemical equations A + B —> C ΔH = 3.0 kJ/mol 2A + 2B —> 2C ΔH = 6.0 kJ/mol C —> A + B ΔH= -3.0 kJ/mol • calorimetry- the measurement of heat ﬂow (energy changes) • speciﬁc heat (s)- the amount of heat required to raise the temperature of 1 g of the substance by 1 degree Celsius • heat capacity (C)- the amount of heat required to raise the temperature of a object by 1 degree Celsius • speciﬁc heat of water- 4.184 J/ g C • qp- constant pressure • qp= ΔH • qv- constant volume • qv= ΔU • KNOW HOW TO FIND q, s, m, or ΔT. • KNOW HOW TO DO BOMB CALORIMETRY • KNOW HOW TO DO HESS’S LAW • standard enthalpy of formation (ΔH )- fhe heat change that results when 1 mole of a compound is formed from its constituent elements in the standard state at sea level • KNOW HOW TO CALCULATE STANDARD ENTHALPY OF FORMATION FOR A REACTION • standard conditions- 1 atm at 25 C o • standard enthalpy of reaction (ΔH rxn)- the enthalpy of reaction carried out under standard conditions • bond enthalpy- the enthalpy change associated with breaking a particular bond in 1 mole of gaseous molecules • KNOW HOW TO CALCULATE BOND ENTHALPY • Light gas particles move faster than heavy gas particles, which move slower. • KNOW THE 4 BASIC ASSUMPTIONS OF KINETIC MOLECULAR THEORY • root mean square (rms) speed (u rms)- the speed of a gas molecule with the average kinetic energy in a gas molecule • diffusion- the mixing of gases due to random motion and frequent collisions • effusion- the escape of gas molecules from a container to a region of vacuum • Graham’s law- states that the rate of diffusion or effusion of a gas is inversely proportional to the square root of its molar mass • pressure- a force applied per unit area • newton (N)- the SI unit of force • pascal (Pa)- the SI unit of pressure • barometer- an instrument that is used to measure atmospheric pressure • standard atmospheric pressure- 1 atm • Boyle’s law- states that the pressure of a ﬁxed amount of gas at a constant temperature is inversely proportional to the volume of the gas • Charles’ and Gay-Lussac’s law- states that the volume of a gas maintained at constant pressure is directly proportional to the absolute temperature of the gas • Avogadro’s law- states that the volume of a sample of gas is directly proportional to the number of moles in the sample at constant temperature and pressure • combined gas law- relates the properties of pressure, temperature, volume, and number of particles into one equation • ideal gas equation- describes the relationship among the four variables P, V, n, and T • ideal gas- a hypothetical sample of gas where the particles are assumed to have zero interaction with each other • gas constant (R)- the proportionality constant and its value and units depend on the units in which P and V are expressed • standard temperature and pressure • pressure- 1 atm • temperature- 0 C (273.15 K) • van der Waals equation- useful for gases that do not behave ideally • compressibility factor- Z • Dalton’s law of partial pressures- states that the total pressure exerted by a gas mixture is the sum of the partial pressures exerted by each component of the mixture • lattice energy- the energy change associated with converting 1 mole of ionic solid to its constituent ions in the gas phase • Born-Haber cycle- the method of determining lattice energy KNNOW TTHESE EQUUATIONS ANDDWHHENNANND HOW TTO USE THEM urms= √3RT/M root mean square (rms) speedrms uArms/uBrms √M BM A Graham’s law rate is proportional to 1/√M barometer equation P=hdg Boyle’s law P1V1=P 2 2 KNOOW THESSE EQUATIONSSANND WHEENANND HOW TTOUSSE THEM Charles’ and Gay-Lussac’s law V1/T1=V 2T2 Avogadro’s law V 1n1=V 2n2 combined gas law P1V1/n1 1= P2V2/n2 2 ideal gas equaiton PV=nRT ideal gas equation (with density) d=PM/RT compressibility factor Z=PV/RT Dalton’s law PtotalΣ Pi mole fractions Xi n/i total
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