chemistry notes CHE 1101
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This 10 page Class Notes was uploaded by Hannah Czajkowski on Sunday May 8, 2016. The Class Notes belongs to CHE 1101 at Appalachian State University taught by Jennifer Cecile in Spring 2016. Since its upload, it has received 12 views. For similar materials see General Chemistry1 in Chemistry at Appalachian State University.
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Date Created: 05/08/16
Chapter 5 Tuesday, April 05, 2016 1:16 PM Oxidation Reduction (redox) reactions Redox Reactions o Transfer or exchange of electrons between two species o To show both components occurring together o Oxidation- lose electrons OiL o Reduction- gain electrons RiG o Oxidation and reduction must occur together o No net change in the number of electrons in a redox reaction (both sides of the reaction must have same number of atoms and charge) o Reducing agent- species or substance that donates electrons o Oxidizing agent- species or substances that accept the electron Oxidation numbers o Number that is assigned to an element in a molecule or ion to reflect qualitatively its state of oxidation o Batteries, metabolism of food, combustion, bleach o Method of bookkeeping to track electrons o Change in oxidation number of an element shows that a redox reaction has occurred o Rules Sum of oxidation number in a neutral species is 0 Atoms of free elements have ON of 0 ON of metal ions in groups 1A 2A and Al have +1, +2, +3 respectively If a transition metal, () determines the oxidation number Nonmetals have negative numbers Bonded with nonmetal, oxygen is +1 Bonded with metal, oxygen is -1 Fluorine is -1 Group 7A is -1 Group 6A is -2 Group 5A is -3 Displacement reaction o Ions oxidize an element o The ions are then reduced o A+BX-->AX+B Using activity series to predict reactions o If M is below H, can displace H from solutions containing H+ o If M is above H, doesn’t react with non-oxidizing acids (HCl, H3PO4) o In general metal below replace ion above T est 4: Chapter 8 Tuesday, April 19, 2016 11:14 AM The Basics of Chemical Bonding Compounds o Types Ionic- bonding involves transfer of electrons Molecular- bonding involves sharing of electrons o Bonding only involves valance electrons Chemical Bonds- attraction between two atoms or ions o Types Ionic bond- electrostatic forces between two ions with opposite charge; result from attraction forces between oppositely charged particles Covalent bond- bond formed between two atoms that share electrons Metallic bond- bond formed only with metals in which each atom in the metal is bonded to several neighboring atoms Electron Configuration o Usually when ions form… Atoms of most representative elements (s and p blocks) tend to gain or lose electrons t obtain nearest Noble gas electron configuration Except He (two electrons), all noble gases have eight electrons in highest n shell o Octet Rule Atoms tend to gain or lose electrons until they have achieved outer (valence) shell containing octet of eight electrons Exceptions Ions or molecules with an odd number or electrons Ions or molecules with less than an octet- transition metals Ions or molecules with more than eight valance electrons Lewis Symbols- representation of molecules that show all electrons, bonding and non-bonding; the structures show the valance electrons as dot arranged around the atomic symbol o Book keeping method to keep track of electrons o Write chemical symbol surrounded by dots for each electron o Can use to diagram electron transfer in ionic bonding o Each H has two electrons through sharing o Can write shared pair of electrons as a line or dots to signify a covalent bond o Drawing Find the sum of valence electrons of all atoms in the polyatomic ion or molecules The central atom is the least electronegative element that isn't hydrogen. Connect the outer atoms to it by single bonds Fill the octets of the outer atoms Fill the octet of the central atom If you run out of electrons before the central atom has an octet, form multiple bonds until it does Polyatomic ions o Carry a net charge which is used for ionic bonding o The atoms within polyatomic ions are bound predominantly by covalent bonds Covalent bonding o Atoms share electrons o Electrostatic interactions in these bonds Attractions between electrons and nuclei Repulsion between electrons Repulsions between nuclei o Attraction of valence electrons of one atom by nucleus of other atom. The electron density shifts as distance between nuclei decreases, probability of finding either electron near either nucleus increases o Pulls nuclei closer together o Two quantities characterize this bond Bond length- bond distance; distance between two nuclei Bond energy- bond strength; amount of energy released when bond formed, decreasing PE or amount of energy must put in to break bonds The distance between bonded atoms decreases as the number of shared electron pairs increases o Nonmetals form more than one covalent bond o Single bond- one shared pair of electrons o Double bond- two shared pair of electrons o Triple bond-- three shared pair of electrons Electronegativity and Bond Polarity o Two atoms of same elements form bond- equal sharing of electrons o Two atoms of different elements form bond- unequal sharing of electrons o Polar bond- bond that carries partial positive and negative charges at opposite ends o Bond is dipole as two poles or two charges involved o Polar molecule- molecule has partial positive and negative charges at opposite ends of a bond o Polar covalent bonds- when two atoms share electrons unequally, a bond dipole results The greater the difference in electronegativity the more polar is the bond The dipole moment produced by two equal but opposite charges separated by a distance, r, is calculated u=Q*r It is measured in debyes (D) o 0-0.5 nonpolar o 0.5-1.7 polar o >1.7 ionic Electronegativity- ability of atoms in a molecule to attract electrons to itself o EN increases from left to right across the period o EN decreases from top to bottom down group Resonance Structures o Two Lewis structures for a compound may be equivalent o Resonance structures illustrate the multiple structures of the molecule o Electron locations are different o Ozone An average of two resonance structures Two bond in ozone are in between single and double bonds o The electrons that form the second C-O bond in the double bonds below do not always sit between that C and O, but rather can move among the the oxygens and the carbon o They are not localized but rather are delocalized Assigning Formal Charge o For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms o Subtract that from the number of valence electrons for that atom- the difference is its formal charge Chapter 9 Tuesday, April 19, 2016 11:14 AM Theories of Bonding and Structure Stereochemistry- Study of the three dimensional shapes of molecules Molecular shape o Shape of a molecule plays an important role in its reactivity o By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule Covalent Bonding Theories o Valence electron pair repulsion theory Abbreviated VSEPR Developed by Gillespie, 1950 Regions of high electron density around the central atom are arranged as far apart as possible to minimize repulsions Five basic geometries based on the number of regions of high electron density around the central atom Eleven molecular geometries (described later) give the arrangement of atoms in space Covalent bonds are formed by the overlap of atomic orbitals Atomic orbitals on their central atom can mix and exchange their character with other atoms in a molecule through hybridization Hybrid orbitals have the same shapes as predicted by VSEPR Sigma bonds Head to head over lap Cylindrical symmetry of electron density about the internuclear axis Sigma bonds are always sigma bonds, because sigma overlap is greater, resulting in a stronger bond and more energy lowering Pi bonds Side to side overlap Electron density above and below the internuclear axis Can never occur alone, must have a sigma bond Multiple bonds In a multiple bond one of the bonds is sigma and the other is pi o Valence Bond Theory Uses hybrid atomic orbitals Developed by L. Pauling, 1930 and 1940 o Covalent bonds form through the sharing of electrons by adjacent atoms. This bonding can only occur when orbitals on the two atoms overlap Approach too molecular structure prediction o Draw a Lewis dot structure ID the central atom Show bonding pairs and lone pairs of electrons on central atoms o Count regions of high electron density on the central atom Include both bonding and lone pairs in the counting o Determine the electronic geometry around the central atom VSEPR is a guide to the geometry o Determine molecular geometry around the central atom Ignore lone pairs of electrons o Adjust molecular geometry for effect of any lone pairs o Find the hybrid orbitals on the central atom o Repeat if greater than one central atom in molecule o Determine molecular polarity from entire molecule geometry using electronegativity difference Predicting shape o Electron pairs, whether they be bonding or nonbonding, repel each other o By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule o Electron pairs are called electron domains o In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain Structural geometries o Each molecule has two types of geometries Electronic geometries- locations of high electron density regions around the central atoms 2 electron domains, Linear, 180 3 electron domains, Trigonal planar, 120 4 electron domains, Tetrahedral, 109.5 5 electron domains, Trigonal Bipyramidal, 120 and 90 6 electron domains, octahedral, 90 Molecular geometries- the arrangement of atoms around the central atoms around the central atoms; no electron pairs Sometimes shape is not given by the electron domain geometry The molecular geometry describes the positions of only the atoms in the molecules, not the nonbonding pairs VSEPR Repulsion strengths o Lone pair to lone pair is the strongest repulsion; bond angle decreases o Long pair to bonding pair is intermediate repulsion o Bonding pair to bonding pair is weakest repulsion o Electron in multiple bonds repel more than electrons in single bonds o Lp/lp>lp/bp>bp/bp Nonbonding pair and bond angle o Nonbonding pairs are physically larger than bonding pairs o Therefore their repulsions are greater; this tends to decrease bond angles in a molecule Multiple bonds and bond angles o Double and triple bonds have greater electron density on one side of the central atom than do single bonds o Bond angles are affected Larger molecules o Look at geometry about a particular atom rather than the geometry of the molecule as a while Polar molecules and geometries o Molecular geometry affects molecular polarity o Bond dipoles cancel or reinforce each other o For polar molecules- on polar bond or one lone pair of electrons on central atom Bonds and lone pairs cannot be symmetrically arranged such that their polarities cancel Molecules are usually polar if all atoms attached to central atm are not the same or there are one or more lone pairs of electrons on the center atom o Nonpolar molecules Symmetrical molecules Non polar because dipoles cancel All five shapes are symmetrical when all domains attached t them are composed of identical atoms Resonances o When writing a lewis structure for species like the nitrate ion, we draw resonance structures to more accurately reflect on the structure of the molecule or ion o For NO3 each of the four atoms in the nitrate ion has a p orbital o The p orbitals on all three oxygens overlap with the P orbital on the central nitrogen o So the pi electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion Energy and Chemical Change Thermochemistry o The study of energies given off by or absorbed by reactions o Study of heat transfer or heat flow o ????H is directly proportional to the amount of reactant or product present o ????H changes sign when a reaction is reversed o ????H for a reaction has the same value regardless of the number of steps Energy ( E )- ability to do work or to transfer heat Kinetic energy- energy in motion, KE=1/2(m)(v^2) Potential energy- energy an object possesses by virtue of its position of chemical composition (stored energy) Chemical energy o PE possessed by chemical o Stored in chemical bonds o Breaking bonds requires energy, forming bonds releases energy Potential energy- dependent upon the mass of the object, the gravitational constant (9.8 m/s^2) and the height of the object compared to a reference height o PE=mgh 1st law of thermodynamics o Energy can neither be created nor destroyed o Can only be converted from one form to another o Total energy of universe is constant o Total energy = potential energy + kinetic energy o (+) q Heat absorbed by system (in) (-) q Heat released by system (out) (+) w Work done on system (in) (-) w Work done by system (out) o Endothermic ????E=+ o Exothermic ????E=- Energy Units o SI unit of energy is Joule (J) o Energy needed to raise the temperature of 1g water by 1 degree C 1 cal = 4.184 J o A nutritional Calorie o 1 Cal = 1000 cal = 1 kcal Temperature- proportional to average kinetic energy of objects particles; higher average kinetic energy means higher temperature and faster moving molecules Heat- total amount of energy transferred between objects; heat transfer is caused by a temperature difference; always passes spontaneously from warmer objects to colder objects; transfers until both are the same temperature Heat flow o System- the component that we are interested in studying o Surroundings- everything else that exchanges heat with the system o Calorimetry may be used to measure the heat flow in a reaction Reaction mixture is the system Calorimeter is the surroundings q-rxn = -q calorimeter o Heat capacity- the amount of heat required to raise the temperature of an object by I K Specific heat = q / m????T Specific heat of water is 4.18 J/g°C o Molar heat capacity- the heat capacity of 1 mole of a substance o Specific heat capacity- the heat capacity of 1 gram of a substance 1C Energy Transfer o Work- energy used to move an object over some distance o w = F x d o Where w is work, F is force, and D is the distance over which the force is exerted (push or pull exerted on an object) o Heat is the transfer f energy between two objects o Heat flows form warmer objects to cooler objects o Energy is the capacity to do work or transfer heat Internal Energy o Is the sum of all kinetic and potential energies of all components of the system; called E o The change in internal every,????E, is the final energy of the system minus the initial energy of the system ????E= Efinal - Einital o When energy is exchanged between the system and the surroundings, it is exchanged as either heat or work ????E = q + w o If ????E>0, E final > E initial Therefore, the system absorbed energy from the surroundings, endergonic o If ????E<0, E final < E initial Therefore the system released energy into the surroundings Work o We can measure the work done by the gas if the action is done in a vessel that has been fitted with a piston (constant pressure) o w=-P????V Enthalpy o When a process takes place at constant pressure and the only work done is this pressure volume work, we can account for heat flow during the process by measuring the enthalpy of the system o Enthalpy is a state function, the internal energy plus the product of pressure and volume o When the system changes at constant pressure, the change in enthalpy is ????H is ????H=????(E+PV) ????H=????E+P????V o We know ????E=q+w and w=-P????V, we can substitute these into the enthalpy expression ????H=q o Endo thermic Reactants + heat-->products o Exothermic Reactants --> products + heat o The change in enthalpy ????H is the enthalpy of the products minus enthalpy of the reactants ????H = H products - H reactants o The quantity ????H is called the enthalpy of reaction, or the heat of reaction o Enthalpy is an extensive property (dependent on mols in the reaction) o ????H for a reaction in the forward direction is equal in size, but opposite in sign, to ????H for the reverse reaction o ????H for a reaction depends on the state of the products and the state of the reactants o When we reverse a reaction we change the sign of ????H o Change in enthalpy depend on state Forming bonds o Attraction of atoms means they move closer together and decrease the potential energy of reacting system o Energy released Breaking bonds o Atoms that are attracted to each other are forced apart and increase the potential energy of reacting system o Energy absorbed Bond strength o Measure of how much energy is needed to break bond or how much energy is released when bond is formed. o Large amount of energy to break = stronger bond o Small amount of energy to break = weak bond Enthalpies of Formation o An enthalpy of formation ????Hf is defined as the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms o The standard state of a substance in its pure form at 1 atm and 25C o The standard enthalpy indicates enthalpy at standard conditions o The stand enthalpy of formation is the change in enthalpy for a reaction that forms 1 mole of a compound from its elements with all substances in their standard states ????H = ????n????H f(products) - ????m????H f (reactants) where n and m are the stoichiometric coefficients
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