Presentation on theme: "Hybridization, Polarity, & Electronegativity"— Presentation transcript:
1Hybridization, Polarity, & Electronegativity Lots of multisyllabic words, I know
2Taking a Step back: Polarity Covalent Bondingatoms “bonded” through sharing of electronsDraw H2, Cl2, and H2OH2 and Cl2 – have an equal sharing of electronsNon Polar CovalentH2O– unequal sharing of electronsPolar Covalent
3Polar CovalentF2 vs. HFThe only time a bond is completely nonpolar is when identical elements are bonded together.In the above example, Fluorine pulls harder on the electrons than the Hydrogen… the Fluorine end of the molecule has more electron density than the hydrogen end.Fluorine is δ- and Hydrogen is δ+
4ElectronegativityIs the ability of an atom in a molecule to draw electrons to itselfIs a combination of Ionization Energy and Electron AffinityIs a MADE UP NUMBER
5ElectronegativityOn the periodic chart, electronegativity increases as you go……from left to right across a row.…from the bottom to the top of a column.
6Electronegativity and Polarity You can use the difference in the electronegativities of two elements to determine whether the bond they form will be nonpolar, polar, or ionicIf the difference in electronegativities is > 1.7, the bond is considered ionic.BondElectronegativity DifferenceType of BondH and H2.1 – 2.1 = 0Non-polar covalentN and H3.0 – 2.1 = 0.9Polar CovalentNa and Cl0.9 – 3.0 = 2.1Ionic
7Polar Covalent BondThe greater the difference in electronegativities, the more polar the bond is
8Polar Covalent BondWhen two atoms share electrons unequally, a bond dipole results.The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:= Qr
9Polar MoleculesPredicting whether or not a molecule will be polar or nonpolar depends on both the bonds and the geometryTo determine whether or not a molecule will be nonpolar:All terminal atoms are identicalAll of the terminal atoms are arranged symmetrically
10Polar Molecules Draw CH4 , BF3, Cl2CO, NH3 Which ones have a net dipole moment and which ones have no net dipole?CH4 and BF3 = no net dipoleCl2CO and NH3 = net dipole
11Overlap and BondingWe think of covalent bonds forming through the sharing of electrons by adjacent atoms.In such an approach this can only occur when orbitals on the two atoms overlap.
12HybridizationAtomic orbitals (s, p, d, f) cannot adequately explain the bonding in moleculesConsider CH4 … how do we get the tetrahedral shape out of the above orbitals
13Hybridization Consider beryllium: An averaging of atomic orbitals into a new set of “hybrid orbitals”Consider beryllium:In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals.
14Hybridization Consider beryllium: if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds.
15Hybrid OrbitalsMixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals.These sp hybrid orbitals have two lobes like a p orbital.One of the lobes is larger and more rounded as is the s orbital.
16Hybrid OrbitalsThese two degenerate orbitals would align themselves 180 from each other.This is consistent with the observed geometry of beryllium compounds: linear.
17Hybrid OrbitalsWith hybrid orbitals the orbital diagram for beryllium would look like this.The sp orbitals are higher in energy than the 1s orbital but lower than the 2p.
18Hybrid Orbitals Think about Boron What’s its electron configuration? How many bonds does it normally form?How?
26Practice ProblemsSpecify the electron-pair and molecular geometry for each underlined atom in the following list. Describe the hybrid orbital set used by this atom in each molecule or ion.BBr3CO2CH2Cl2CO3-2