Presentation on theme: "Lots of multisyllabic words, I know. Taking a Step back: Polarity Covalent Bonding atoms bonded through sharing of electrons Draw H 2, Cl 2, and H 2 O."— Presentation transcript:
Lots of multisyllabic words, I know
Taking a Step back: Polarity Covalent Bonding atoms bonded through sharing of electrons Draw H 2, Cl 2, and H 2 O H 2 and Cl 2 – have an equal sharing of electrons Non Polar Covalent H 2 O– unequal sharing of electrons Polar Covalent
F 2 vs. HF The 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+
Electronegativity Is the ability of an atom in a molecule to draw electrons to itself Is a combination of Ionization Energy and Electron Affinity Is a MADE UP NUMBER
Electronegativity On the periodic chart, electronegativity increases as you go… …from left to right across a row. …from the bottom to the top of a column.
Electronegativity 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 ionic If the difference in electronegativities is > 1.7, the bond is considered ionic. BondElectronegativity Difference Type of Bond H and H2.1 – 2.1 = 0Non-polar covalent N and H3.0 – 2.1 = 0.9Polar Covalent Na and Cl0.9 – 3.0 = 2.1Ionic
Polar Covalent Bond The greater the difference in electronegativities, the more polar the bond is
Polar Covalent Bond When 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: = Q r
Polar Molecules Predicting whether or not a molecule will be polar or nonpolar depends on both the bonds and the geometry To determine whether or not a molecule will be nonpolar: All terminal atoms are identical All of the terminal atoms are arranged symmetrically
Polar Molecules Draw CH 4, BF 3, Cl 2 CO, NH 3 Which ones have a net dipole moment and which ones have no net dipole? CH 4 and BF 3 = no net dipole Cl 2 CO and NH 3 = net dipole
Overlap and Bonding We 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.
Hybridization Atomic orbitals (s, p, d, f) cannot adequately explain the bonding in molecules Consider CH 4 … how do we get the tetrahedral shape out of the above orbitals
Hybridization 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.
Hybridization 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.
Hybrid Orbitals Mixing 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.
Hybrid Orbitals These two degenerate orbitals would align themselves 180 from each other. This is consistent with the observed geometry of beryllium compounds: linear.
Hybrid Orbitals With 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.
Hybrid Orbitals Think about Boron Whats its electron configuration? How many bonds does it normally form? How?
Practice Problems Specify 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. a) BBr 3 b) CO 2 c) CH 2 Cl 2 d) CO 3 -2
Practice Problems Answers Electron-Pair MolecularHybrid GeometryGeometryOrbital Set a) Trigonal PlanarTrigonal Planarsp 2 b) LinearLinearsp c) Tetrahedraltetrahedralsp 3 d) Trigonal planartrigonal planarsp 2