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chem notes week 12

by: Andrea Scota

chem notes week 12 CHE 106 - M001

Andrea Scota
GPA 3.7

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Chemistry notes for week 12
General Chemistry Lecture I
R. Doyle
Class Notes
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This 7 page Class Notes was uploaded by Andrea Scota on Sunday November 29, 2015. The Class Notes belongs to CHE 106 - M001 at Syracuse University taught by R. Doyle in Fall 2015. Since its upload, it has received 69 views. For similar materials see General Chemistry Lecture I in Chemistry at Syracuse University.

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Date Created: 11/29/15
Pink- mentioned in class Chem Notes Week 12 TEXTBOOK CHAPTER 9 Molecular Geometry and Bonding  The shape and size of molecules- their molecular architechture- are defined by angles and distances between the nuclei of their component atoms Molecular Shapes (9.1)  Lewis structures do not indicate shapes of molecules- they simply show the numbers and types of bonds  Band angle: determines the shape of the molecule; angles made by lines joining nuclei of atoms in a molecule  Bond angles and lengths determine the shape and size of a molecule  Molecules with single central atom (A) surrounded by n (number) of atoms B with the general formula AB n o Number of shapes possible depends on the value of n, many different geometric shapes  Shapes of most AB mnlecules are derived from 5 basic geometric shapes, all of which are highly symmetric arrangements of n B atoms around central A atom  We can derive additional shaped, such as bent and trigonal pyramidal, by starting with one if these 5 basic arrangements  When A is a representative element in the s or p block, we can predict their shapes on AB nodel due to a valence-shell electron-pair repulsion (VSEPR) model The VSEPR Model (9.2)  An electron domain defines a region in ehich an electron is most likely to be found (regions around a central atom) o Bonding pairs: electrons involved in making bonds in molecules o Lone pairs (nonbonding pairs): electrons not involved in making bonds (they belong to only one atom), define the electron domains  According to VSEPR model, electron domains orient themselves to minimize electrostatic repulsions (they remain as far apart as possible)  In general; each nonbonding pair, single bond, or multiple bond produces a single electron domain around the central atom in a molecule  Shaped of different AB molecules or ions depends on the number of n electron domains surrounding the central atom  Electron geometry: arrangement of electron domains about the central atom of AB nolecule or ion  Molecular geometry: arrangement of only the atoms in a molecule or ion (any nonbonding/lone pairs are not included)  In determining the shape of a molecule, first use the VSEPR model to find the electron domain geometry and then predict the molecular geometry o When all the electron domains in molecules exist as bonds then the electron domain geometry and the molecular geometry are equal (aka, when there are no lone pairs) o When 1+ electron domain involved nonbonding lone pairs, we must remember that molecular geometries involve only electron domains due to bonds o General Steps:  1) Draw Lewis structure of molecule or ion and count the number of electron domains  2) Determine the electron domain geometry so that the repultions of the electrons are minimized  3) Use arrangement of bonded atoms to determine the molecular geometry, counting only the bonding electron domains  Bond angles decrease as the number of nonbonding lone electron pairs increase  Bonding pair of electrons is attracted to 2 nuclei of bonded atoms, but lone pair electrons are only attracted to one— so, since the bonding pair is attracted to 2, its electron domain is more spread out in space o Nonbonding lone electron pairs take up more space than bonding pairs  Electron domains for nonbonding electron pairs exert greater repulsive forces on adjacent electron domains and end to compress bond angles  Electron domains from multiple bonds exert slightly greater repulsions than those of single bonds  Molecules with 5 or 6 electron domains around the central atom have molecular geometried based on either trigonal biparyamidal (5 domains) or octahedral (6 domains) electron domain geometries o Trigonal: electrons can be in axial (straight up/down) or equatorial (form equilateral triangle) position o Octahedral: all bond angles form 90 degree angles and have 6 equivalent vertecies Molecular Shape and Molecular Polarity (9.3)  Recall: bond polarity is a measure of how equally electrons in a bond are shared between 2 atoms of a bond. As the difference in electronegativity between the 2 atoms increases, so does the polarity of the bond  Recall: dipole moment of diatomic molecule is the measure of the amount of charge separation is in the molecule  For molecule with 2+ atoms, the dipole moment depends on both the polarities of the individual bonds and the geometry of the molecule  Bond dipole: dipole moment due only to the two atoms in that bond  Sipole moment of a bond depends on the vector sum of the dipole moments associated with the individual bonds o Bond dipoles/dipole moments are vector quantities because they have both magnitude and direction  Certain molecular shapes, such as linear AB and2trigonal planar AB , 3 lead to the cancellation of the bond dipole, producing a nonpolar molecule whose overall dipole moment is 0  Shapes, such as bent AB an2 trigonal pyramidal AB , bond3dipoles do not cancel out and the molecule will therefore be polar (nonzero dipole moment) Covalent Bonding and Orbital Overlap (9.4)  The idea of atomic orbitals and Lewis’ notion of electron pair bonds leads to a model of chemical bonding called valence-bond theory  Valence-bond theory: bonding electron pairs are concentrated in regions between atoms, and nonbonding electron pairs lie in directed regions of space  Covalent bonds are formed when atomic orbitals on neighboring atoms overlap one another o Overlap region is one of greater stability for 2 electrons because of simultaneous attraction to the 2 nuclei o The greater the overlap the stronger the bond formed o Increased overlap brings electrons and nuclei closer together Hybrid Orbitals (9.5)  Recall: atomic orbitals are mathematical functions that come from the quantum mechanical model for atomic structure  To explain molecular geometries, we often assume that the atomic orbitals of an atom mix to form new hybrid orbitals o The shape of a hybrid orbital is different from the shapes of the original atomic orbitals  Hybridization is the process of mixing atomic orbitals o The number of hybrid orbitals is equal to the number of atomic orbitals that are mixed  Mixing s and p orbitals leads to the hybridization which leads to hybrid atomic orbitals that have a large lobe directed to overlap with orbitals on another atom to make a bond  According to valence-bond model, a linear arrangement of electron domains implies sp hybridization  Hybrid orbitals can also accomidate nonbonding pairs  Particular mode of hybridization can be associated with each of 3 electron-domain geometries: o Linear = sp o Trigonal planar = sp 2 3 o Tetrahedral = sp  Procedure: o 1) Lewis Structure o 2) From structure, find the electron-domain geometry o 3) Specify the hybridization needed to achieve the geometry 3  EX: H 2 is a tetrahedral so sp hybridization  Bonding in hypervalent molecules (those with more than octet of electrons) are not as readily discussed in terms of hybrid orbitals Multiple Bonds (9.6)  In covalent bonds, the electron density that is concentrated along the line connecting the nuclei forms a sigma () bond (head to head, single bonding) o Passes through the middle of the overlap region  Bonds can be formed by the sideways overlap of p orbitals (in multiple bonds) called a pi () bond (happen when there are double or triple bonds) o Overlap regions lie above and below the intermuclear axis o Electron density is not concentrated on the internuclear axis o Generally weaker than  bonds  In almost all cases, single bonds are  bonds  A double bond consists of one  bond and one  bond  A triple bond consists of one  bond and two  bonds 2  p (pee-pie) bond: the unhybridized 2p atomic orbital of an sp hybridized atom (p orbital that can be involved in forming a  bond) o the orbitals that overlap to form  bonds  The formation of a  bond requires that molecules adopt a specific orientation o EX: two CH gro2ps in C H mus2 l4e in the same plane— as a result the presence of the  bonds introduces rigidity into molecules  Localized bonding electrons in molecules means the  and  electrons are associated totally with two atoms that form the bond  Molecules with 2+ resonace structures are not localized (like C H ) 6 6  Cannot describe bonds in C H a6 i6dividual bonds between neighboring atoms o We say it has six electron  system delocalized among the six carbon atoms  Delocalized electrons means that  bonds are pread among several atoms  Delocalization gives molecules a certain stability and gives organic molecules their colors Molecular Orbitals (9.7)


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