Chapters and 7.4 plus 8.1 and 8.3-5: Bonding, Solids, VSEPR, and Polarity

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1 Chapters and 7.4 plus 8.1 and 8.3-5: Bonding, Solids, VSEPR, and Polarity Chemical Bonds and energy bond formation is always exothermic As bonds form, chemical potential energy is released as other forms of energy (usually heat and/or light) Bonds form because the nucleus of one atom is attracted to the electron cloud of a nearby atom If the atoms are too far apart, the attraction is too weak to have any effect Once the atoms reach some critical distance, they will attract each other and move toward some minimum optimal distance If the atoms get too close, the nuclei begin to repel strongly and the atoms move apart Chemical bonds all chemical bonds have two positively charged nuclei mutually attracted to electron density between the nuclei Bond strength is inversely related to bond length Long bonds tend to be weaker Short bonds tend to be stronger Types of chemical bonds (listed strongest to weakest but overlap does occur) Bond Type Identifying Characteristic Metallic sea of freely moving e Ionic e are transferred major Polar covalent e shared unevenly form bond types Nonpolar covalent e shared evenly molecules Hydrogen bonds e nearly naked p+ (N, O, or F) Dipole dipole bonds molecule with one and one + end intermolecular bonds London dispersion forces momentary uneven e distribution Valence electrons and ion formation Positive Ion Formation (cations) Metals tend to lose all their valence e forming cations (because they have low ENC) The resulting cations generally have full s and p sublevels in the next lower energy level These cations do not lose e in this lower energy level because the ENC jumps way up Transition metals tend to lose only the highest energy level s electrons Such configurations are known as pseudo-noble gas configurations Negative Ion Formation (anions) Nonmetals tend to gain e forming anions (because they have high ENC) They gain e until they have full s and p sublevels They do not gain e in the next higher energy level because the ENC is negative there This explains why both metals and nonmetals tend to acquire an outermost energy level that resembles a noble gas configuration (an octet) or pseudo-noble gas configuration Only the first energy level is full with just two e (go on to the next page)

2 Ionic bonds and ionic compounds Ionic bonds form when positive cations are attracted to negatively charged anions Because of the high ENC difference between metals and nonmetals, compounds formed by the reaction of a metal with a nonmetal will be ionic compounds (see Periodic Table below) nonmetal metalloid metal Example of ionic bond formation using LED formulas: NaCl + Na. +. Cl : Na : Cl : Lattice energy energy required to separate 1 mole of ions in an ionic compound Increases as the charge on individual ions increases Decreases as bond length increases Covalent bonds occur when the electronegativity difference between bonding atoms is so slight that the electrons tend to be shared Uneven sharing results in a polar covalent bond Even sharing results in a (mostly) nonpolar covalent bond Very even sharing (ΔEN = 0) results in a strictly nonpolar covalent bond (pure covalent bond) Identifying the major bond types Metals are easy to identify, just check the periodic table. Ex.e (s), Ni (s) or Hg (l) For ionic, polar covalent and nonpolar covalent bonds, you must calculate the electronegativity differences (ΔEN) using pages 10 and 11 in the NYS Chemistry Regents Tables or page 265 in the Glencoe text ΔEN Bond Type Mostly ionic bond 1.7 Mostly polar covalent bond 0.4 Mostly covalent bond (nonpolar covalent bond) 0 Pure covalent bond Examples: NaCl Na = 0.9 Cl = 3.2 ΔEN = = 2.3 ionic bond HCl H = 2.2 Cl = 3.2 ΔEN = = 1.0 polar covalent bond PH 3 P = 2.2 H = 2.2 ΔEN = = 0 nonpolar covalent bond HF H = 2.2 F = 4.0 ΔEN = = 1.8 very polar covalent bond!

3 Molecules are groups of atoms held together by covalent bonds (polar covalent, nonpolar covalent, or a mix of the two covalent types) Intermolecular forces weak chemical bonds that hold a molecule to another nearby molecule Hydrogen bonds especially strong dipole-dipole forces that occur when hydrogen polar covalently bonds to N, O, or F atoms and then the resulting nearly naked proton formed is strongly attracted to the lone electron pair of the N, O, or F of another molecule (or a N, O, or F of another part of a large molecule) Dipole-dipole forces due to shape and polar bond character, many molecules form permanent dipoles (one end of the molecule is partially positive and the other end of the molecule is partially negative. The dipoles align so that opposite charges will be near each other. These forces are not nearly as strong as hydrogen bonds for several reasons. First, they are usually significantly longer (300 pm for dipole-dipole forces compared to 169 pm for hydrogen bonds). Secondly, the dipole-dipole interaction for HCl is much weaker than a true hydrogen bond because the larger electron cloud in chlorine (period 3) can partially cover the proton whereas period 2 elements (N, O, and F) have much smaller and more defined electron clouds which cannot re-cover the proton. van der Waals forces (sometimes called London dispersion forces) due to the fact that electrons are in constant motion in molecules, sometimes the electrons can form an momentary dipole if they become dispersed unevenly. The small buildup of charge on one molecule can cause electrons on a nearby molecule to migrate toward the slightly positive end of the first molecule thereby inducing another momentary dipole in the second molecule. Even though van der Waals forces can constantly shift and reform, their transient nature causes them to be much weaker than permanent dipole-dipole interactions. Factors that will affect intermolecular bonds and, therefore, state (solid, liquid, or gas) Temperature: High temperatures produce gases (molecules move too fast for bonds to form) Strength of the bond: Hydrogen bonded H 2 O is a liquid at room temperature Dipole dipole bonded H 2 S is a gas at room temperature Molecular mass: High mass molecules move slower allowing more time for bonds to form Shape: Long, flat molecules provide more surface area making stronger bonds (pentane is a liquid) Rounder molecules have less surfaced area making weaker bonds (neopentane is a gas) Polarizability: more e or e that disperse more readily create stronger van der Waals forces

4 The four types of solids and their properties Solid Type Sketch Properties Examples Metallic Shiny luster, conducts heat and electricity, malleable and ductile. High melting and boiling points. Copper, zinc. (Name any two metals.) Ionic Network Molecular Usually has a sheen (but not shiny), conducts in liquid state but not in solid state, brittle. High melting and boiling points. Usually has a sheen (but not shiny), does not conduct in either liquid or solid state, brittle high melting and boiling points often decomposing. Usually has a sheen (but not shiny), usually do not conduct in either liquid or solid state, brittle, low melting and boiling points often exist as gases and liquids. Sodium chloride, cupric sulfate. (Any salt.) Quartz, diamond. (Ruby, sapphire, graphite.) Ice [H 2 O (s) ] is hydrogen bonded, iodoform [CHI 3 (s) ] is a dipole-dipole solid, and dry ice [CO 2(s) ] is held by van der Waals forces (or London dispersion forces). (go on to the next page)

5 Lewis Electron Dot (LED) formulas and structures Electron Dot Notation shows only the valence electrons using dots (or x or whatever). O. : or. F : Electron Dot Formulas : or H : O : H Lewis Structures F : or H O H A few rules for Lewis structures Hydrogen atoms will almost never be central atoms Carbon atoms will almost always be central atoms Most atoms will form an octet (end up with 8 e in their valence shells) Some metals will act as if they are losing e so they are deficient (less than 8 e ): Be, B, Al Be F : (The Be is ionic, Be 2+, but takes the same shape as if molecular) Only atoms in period 3 or higher (4, 5, 6 and 7) can have extended octets (more than 8 e ) Xe F : (The Xe has 10 e, not just 8 with extras in the empty 5d orbitals) Only nonmetals will form multiple bonds and only then if they are e deficient O = O Valence Shell Electron Pair Repulsion (VSEPR) theory [see chart next page] The shapes of molecules can be predicted using VSEPR theory and LED structures Create the LED structure Identify the Molecule Type PH 3 is AB 3 E Use the 3D VSEPR sketch Remove atoms from VSEPR for lone e pairs Polar and nonpolar molecules Usually a combination of asymmetry and polar bonds results in polar molecules Examples O = C = O Even though the bonds are polar, the linear molecule will be nonpolar Even though the bonds are nonpolar, the shape results in a polar molecule

6 The chart below shows VSEPR geometries and actual molecule geometries for up to six electron pairs about a central atom Number of e - Pairs VSEPR 3D VSEPR Sketch Molecule Type Molecule Geometry Hybridization Examples 2 linear AB 2 linear sp BeF 2, CO 2 3 planar AB 2 E bent, 120 sp 2 (NO 2 ) AB 3 planar sp 2 BF 3, SO 3 4 tetrahedral AB 2 E 2 bent, 109 sp 3 H 2 O AB 3 E pyramid sp 3 NH 3 AB 4 tetrahedral sp 3 CH 4 5 bipyramid AB 2 E 3 linear sp 3 d XeF 2 AB 3 E 2 T-shaped sp 3 d ClF 3 AB 4 E AB 5 distorted tetrahedron bipyramid AB 2 E 4 linear sp 3 d 2 sp 3 d SF 4 sp 3 d PCl 5 6 octahedral AB 3 E 3 T-shaped sp 3 d 2 AB 4 E 2 square planar sp 3 d 2 XeF 4 AB 5 E square pyramid sp 3 d 2 ClF 5 AB 6 octahedral sp 3 d 2 SF 6 Solubility Since only nonpolar molecules can produce van der Waals forces, only another nonpolar molecule can dissolve a nonpolar molecule Only a dipole can attract other permanent charges, so only dipoles can dissolve ionic compounds and other polar molecules Scientists usually express these observations by the rule: Like dissolves like

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