Topic 8: Bonding : Topic 8: Bonding Zumdahl 7e, Chapter 8
I. Valence Electrons & Bonding : I. Valence Electrons & Bonding Recall that valence refers to the outermost electrons, the ones that are actually involved in bonding when chemical reactions occur
Representing valence electrons
Use either two dots or two x’s for a pair of unbonded electrons, and a line for two bonded electrons
The Octet Rule : The Octet Rule Most atoms would like to have an octet of valence electrons (most stable configuration, like noble gases)
All bonding and ion-forming is to that purpose
E.g. NaCl is formed when Na gives its 3s1 electron to Cl, thereby completing Cl’s 2p orbital; both now have octets
H, He, Li, and Be are exceptions; they would prefer duets (two electrons) because that’s all that fits in the s subshell; B prefers a sextet
Elements lower on the periodic table can fit ten or more electrons due to the availability of the d subshell
Lab #9: Molecular Models : Lab #9: Molecular Models Do this on binder paper, not your lab notebook
Pre-lab: Draw Lewis dot diagrams
Set up the table at home; we will be drawing several molecules together in class today
You will actually build the models on Thursday
The directions in the lab refer to electronegativity, which is similar to electron affinity
The closer to fluorine (4.0), the more electronegative
Oxygen (3.5), Cl (3.0), N (3.0)
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II. VSEPR Geometry : II. VSEPR Geometry VSEPR theory
Valence Shell Electron Pair Repulsion: electron pairs repel each other because of the negative charges; they will arrange themselves as far from each other as possible
b. 3D representation of molecules : b. 3D representation of molecules This cannot always be accurately represented in 2D (works for BF3, but not for CH4) We use a line to represent bonds in the plane of the page, a triangular wedge to represent a bond coming out of the page toward the viewer, and a dotted line/wedge to represent a bond going into the page away from the viewer
c. Electron vs. molecular geometries : c. Electron vs. molecular geometries Electron geometry is the basic arrangement of the electrons (bonded & unbonded) around the central atom
Linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral
Depends only on how many groups there are
Molecular geometry is the arrangement of the bonded electrons only around the central atom
Although the unbonded pairs still affect the arrangement, there is a different name for the resulting geometry if not all the electron groups are bonded
E.g. “Trigonal pyramidal” instead of tetrahedral if one of the electron groups is unbonded; “bent” if two are unbonded
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More Exceptions to the Octet Rule : More Exceptions to the Octet Rule Some elements in the third period or higher can have more than eight electrons because the d-orbital is available for bonding
Why n=3? Because at this point, the 3d is available; no such thing as 2d!
E.g. PF5, SF6, and XeF4
III. Polarity : III. Polarity Electronegativity
Electronegative elements attract electrons
Polar bonds
When electrons are equally shared, a nonpolar covalent bond is made
When electrons are unequally shared, a polar covalent bond is made
Usually labeled as “polar” when featuring one of the big four electronegative elements (FONCl)
c. Dipole moments : c. Dipole moments We can draw a vector to indicate the polar bonds in a molecule; the arrow points to the partially negative side
When all the vectors cancel out, the molecule is nonpolar (even though it has polar bonds)
When the vectors do not cancel out, the molecule has a dipole moment; it is a polar molecule
IV. Hybridization : IV. Hybridization Examination of the CH4 molecule shows that all bond lengths are the same
How does this work, if two of the valence electrons are from the s-subshell, and therefore have less energy than the other six electrons from the p-subshell?
Answer: the s and p orbitals blend together to make a hybrid orbital called the sp3 orbital The hybrid orbitals are “assembled” based on the number of electron groups (bonded & unbonded); linear electron geometry results in sp hybridization; octahedral electron geometry results in sp3d2 hybridization
Sigma and Pi Bonds : Sigma and Pi Bonds In the molecules HCN, by that reasoning, the bond is sp hybridization since it is linear (only two groups of electrons)
How do two orbitals make up the four bonds on carbon (1 single, 1 triple)?
Only the first bond is hybridized: called the sigma bond; this makes the initial H—C—N
Any additional bonds (to make a double or triple bond) are made with the remaining unhybridized orbitals (in this case, p orbitals): called the pi bonds
V. Bonding Summary : V. Bonding Summary Ionic bonds: electrons are donated by metals, accepted by nonmetals; the bond results from the attraction of oppositely-charged cations and anions (NaCl is actually Na+Cl-)
Covalent bonds
Nonpolar: electrons are shared equally (H – H)
Polar: electrons are shared unequally (H – F)
Coordinate/dative: electron deficient molecules share another molecule’s lone pair (F3B – NH3)
Network: giant lattice of nonmetal atoms held together by shared electrons (diamond, quartz)
Metallic: metal atoms share delocalized electrons (Cu)
VI. Intermolecular Forces : VI. Intermolecular Forces Covalent, ionic, and metallic bonds are intramolecular forces (within the molecule)
Attractions between molecules (what connects one water molecule with another to make liquid?) are intermolecular forces
Think: international relations, interstate freeways
Hydrogen bonding
In bonds between H and F, O, or N, the attraction of the electrons is so strong that the H atom becomes a naked proton
Almost an ionic bond!
Hydrogen Bonding cont’d. : Hydrogen Bonding cont’d. Molecules with this type of bond (H-F, H-O, or H-N) experience greater than normal attractions to each other because of the very strong dipole moment
This means that substances which undergo hydrogen bonding have:
Higher than usual melting/boiling points
Greater density
Greater viscosity
Water’s hydrogen bonding capabilities make capillary action, surface tension, and ice-skating possible
NOTE: the hydrogen bond is the force between the molecules, not the H-F/H-O/H-N bond
b. Dipole-dipole attractions : b. Dipole-dipole attractions Slightly polar molecules like HCl, HBr, etc. are still attracted to each other, but not as strongly (weaker partial charges) c. London dispersion forces Nonpolar molecules like H2S do not have a permanent dipole moment – so how do they ever become liquids or solids?
Temporary dipoles are caused by the constant movement of electrons, which can then induce dipoles in surrounding molecules
Influences on Physical Properties : Influences on Physical Properties E.g. Noble gases’ BP
He 4 K
Ne 27 K
Ar 87 K
Kr 121 K
Xe 165 K E.g. Hydrides of VA
NH3 240 K
PH3 186 K
AsH3 213 K
SbH3 248 K
E.g. Hydrides of IVA
CH4 112 K
SiH4 161 K
GeH4 183 K
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