CHEM 111 1nd EditionExam # 3 Study Guide Lectures: 17 - 24Lecture 17: Octet Exceptions, Bond Strength; 4.8-9 (October 8) Understand free radicals.Free radicals include reactive molecules that have an odd number of valence electrons and contain atoms with incomplete octets. Atoms of elements in the third row of the periodic table with Z> 12 and beyond may expand their valence shells to accommodate more than an octet of electrons.Define and explain the relationships between bond order, bond energy, and bond length.Bond order is the number of bonding pairs in a covalent bond. Bond energy is the energy change that accompanies the breaking of 1 mole of a particular covalent bond in the gas phase. As the bond order between two atoms increases, the bond length decreases and the bond energy increases.Lecture 17: VSEPR; 5.1-2 (October 10)What is the effect of minimizing repulsion between valence electron pairs? Minimizing repulsion between pairs of valence electrons (the VSEPR model) results in the lowest energy orientations of bonding and nonbonding electron pairs and accounts for the observed geometries of molecules. How does the steric number of a molecule affect its shape?The shape of a molecule can be determined by its steric number (the sum of the number of bonded atoms and lone pairs around a central atom) and the electron-group geometry, or arrangement of atoms and lone pairs. Molecules with SN = 2 and no lone pairs on the central atom have a linear electron-group geometry and linear molecular geometry, while the electron group geometries of molecules with steric numbers 3 to 6 are trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral, respectively. Understand molecular geometry structures and know their names. Molecular geometries called angular (bent), trigonal pyramidal, seesaw, T-shaped, square pyramidal, and square planar results from different combinations of bonded atoms and lone pairs about a central atom. The observed bond angles in molecules deviate from the ideal values as a result of unequal repulsions between lone pairs and bonding pairs of electrons. Lecture 18: Molecular Polarity; 5.3 (October 13)Explain bond dipoles.Two covalently bonded atoms with different electronegativities have partial electrical charges ofthe opposite sign, creating a bond dipole. If the individual bond dipoles in a molecule do not offset each other, the molecule is polar. If they do offset each other, the molecule is nonpolar. A polar molecule has a permanent dipole moment (u), which is a quantitative measure of thepolarity of the molecule. Lecture 19: Hybrid Orbitals; 5.4 (October 15)Understand the effects of overlapping orbitals. According to valence bond theory, the overlap of half-filled orbitals results in covalent bonds between pairs of atoms in molecules. Molecular geometry is explained by the mixing, or hybridization, of atomic orbitals to create hybrid atomic orbitals. Mixing one s and three p orbitals forms four sp3 hybrid orbitals. Overlap between sp3orbitals and other atomic or hybrid orbitals results in up to four sigma (σ) bonds and a tetrahedral orientation of valence electrons. Mixing one s and two p orbital forms three sp2 hybrid orbitals. Overlap between sp2 and other atomic or hybrid orbitals results in up to three σ bonds and trigonal planar orientation of valence electrons. Mixing one s and one p orbital forms two sp hybrid orbitals. Overlap betweentwo sp hybrid orbitals results in up to two σ bonds oriented linearly to one another. Mixing one s orbital, three p orbitals, and one d orbitals yields five equivalent sp3d hybrid orbitals with lobes that point toward the vertices of a trigonal bypyramid. Overlap between sp3dorbitals and other atomic or hybrid orbitals results in up to five σ bonds. Mixing one s orbital, three p orbitals, and two d orbitals gives six equivalent sp3d2 hybrid orbitals that point toward the vertices of an octahedron. Overlap between sp3d2 orbitals and other atomic or hybrid orbitals results in up to six σ bonds. What kind of covalent bonds form pi (π) bonds?Covalent bonds, in which the electron density is greatest either above and below or in front of and behind the bonding axis, are pi bonds. Lecture 20: London Forces; 6.1 (October 17)Understand polarizability. When does it increase/decrease?All atoms and molecules experience London dispersion forces due to their polarizability and the existence of temporary dipoles even in particles that have no permanent dipole. Polarizability increases with increasing particle size; so larger atoms and molecules experience stronger London dispersion forces than small ones. Stronger London dispersion forces lead to higher boiling points and greater viscosities. What are hydrocarbons? What’s the difference between them and constitutional isomers?Hydrocarbons are compounds whose molecules contain only hydrogen and carbon atoms. Alkanes are hydrocarbons in which each carbon atom is bonded to four other atoms. Constitutional isomers have the same molecular formulas but different connections between the atoms in their molecules.Lecture 21: Polar Interactions; 6.2 (October 20)Define and explain the strongest dipole interactions.Ions in aqueous solutions interact with water molecules through ion-dipole interactions, forming a sphere of hydration around the ion. Molecules with permanent dipoles interact through a combination of London dispersion forces and dipole-dipole interactions. Thestrongest dipole-dipole interactions are hydrogen bonds, which form between H atoms bonded to N, O, and F atoms and other N, O, and F atoms. Polar organic compounds contain polar functional groups. Among them are carbonyl groups, which contain C=O bonds. When the carbonyl carbon atom is bonded to two other carbon atoms in a molecular structure, the compound is called a ketone. Other polar functional groups include the hydroxyl group (OH) in alcohols. Lecture 22: Solubility; 6.3 (October 22)What happens when polar solutes dissolve in polar solvents?Polar solutes dissolve in polar solvents when the dipole-dipole interactions between solute and solvent molecules offset the interactions that keep either solute molecules or solvent moleculestogether. The limited solubility of nonpolar solutes in polar solvents is a result of dipole-induceddipole interactions. Explain the difference between hydrophilic and hydrophobic.Hydrophilic substances are more soluble in water than are hydrophobic substances, which are more