Chemical Bonding

Study Guide 1. Chemical Bonding a. Link www.youtube.com/watch?v=Ac7G7xOG2Ag 2. Octet Rule a. Atoms often gain, lose or share electrons to achieve the same number of electrons as the noble gas closest to them on the Periodic Table. The noble gases have very stable electron arrangements. Because all noble gases except He have eight valence electrons, many atoms undergoing reactions end up with eight valence electrons. This observation has led to a guideline known as the octet rule: Atoms tend to gain, lose and share electrons until they are surrounded by eight valence electrons. 3. Lewis Symbols a. Consists of the element’s chemical symbol plus a dot for each valence electron b. The dots are placed on the four sides of the symbol with each side being able to accommodate up to two electrons. All four sides are equivalent, which means the choice of where to place two electrons, or one electron, is arbitrary. c. The formation of covalent bonds can be represented with Lewis Symbols. d. In forming the H-H covalent bond, each hydrogen atom acquires a second electron, achieving the stable, noble gas electron configuration of Helium. e. For Cl2, by sharing the bonding electron pair, each Cl atom has eight electrons in its valence shell, thus achieving the noble gas configuration of Ar. 1. How to Draw Lewis Structures f. Step 1: Add up all the valence electrons for every atom in the molecule. For any anions, add one electron to the total number for each negative charge. For cations, subtract one electron from the total number for each positive charge. g. Step 2: Write the symbols for all the atoms in the molecule, showing which atoms are attached to which. Then connect them with a single bond. Chemical formulas are often written in the order in which the atoms are connected. In many polyatomic ions, the first atom in the formula is the central atom in the Lewis structure. Usually, but not always, the central atom is the less electronegative. h. Step 3: Complete the octets around all the atoms bonded to the central atom, except for hydrogen, which only wants two electrons around it. i. Step 4: Place leftover electrons on the central atom, even if doing so results in more than an octet of electrons around it. j. Step 5: If there are not enough electrons to give the central atom an octet, try multiple bonds. 4. Ionic Bonding k. Ionic substances generally result from the interaction of metals on the left side of the Periodic Table with non-metals on the right side. NaCl(s) ΔH = -410.9 kJ Ex. Na(s) + ½ Cl2(g) i. NaCl is composed of Na+ and Cl- ions arranged in a three dimensional array l. Forms when Na atoms transfer their single 3s electrons to chlorine m. Each ion ends up with an octet of electrons. Na like Ne and Cl like Ar 5. Ionic Formulas a. In order to predict the chemical formulas of ionic compounds we have to: b. Step 1: Determine what charge each element will preferentially form as a cation or anion, based on its location on the Periodic Table. For example: Na will form Na+. Al will form Al+3 c. Step 2: Write the cation on the left and the anion on the right in the formula. d. Step 3: If necessary, add subscripts next to the cation and anion to ensure that their total combined charges equal zero. i. For example Ax+ + ByAy B x 6. Lattice Energy a. The attraction between positively and negatively charged ions makes ionic compounds very stable. These compounds stabilities are measured by their lattice energy, which is the energy required to completely separate one mole of a solid ionic compound into its individual gaseous ions. i. Ex. NaCl(s) Na+(g) + Cl-(g) ΔH = +788 kJ/mol ii. Notice this process is extremely endothermic b. Lattice Energy, Eel, can be described by the following equation: i. Eel = [KQ1Q2]/d ii. Where K is a constant, Q1 and Q2 are the charges on the individual ions, and d is the distance between their nuclei iii. Thus the lattice energy increases as the charges on the ions increase, and as their radii decrease 7. Electronegativity a. An element’s thirst for electrons. The more an element wants to steal electrons to feel like a noble gas, the more electronegative it is. b. Increases as you move up and to the right of the Periodic Table. 8. Polarity a. When two bonded atoms in a covalent molecule have an electronegativity difference between them, the more electronegative one will ‘hog’ the electrons. b. This forms a partial charge in the molecule called a diploe. Covalent molecules with diploes are called polar molecules. The degree of uneven sharing of electrons between two atoms is called polarity. The greater the difference in electronegativity, the more polar a bond is. c. Measure polarity in units of Debyes (D). The greater the polarity, the larger the number of Debyes. We can say that bonds with dipoles have dipole moments. Drawn in two different ways. 9. Bond Enthalpy a. Measure of how strong a bond is. b. Can use bond enthalpies to estimate the reaction enthalpies of different reactions. Have to figure out what kinds of bonds are broken and what kinds are formed in a reaction. Then just do the math to calculate the reaction enthalpy. c. The full ΔH Rxn is then calculated by subtracting the combined enthalpies of the bonds formed from the combined enthalpies of the bonds broken. 10. Formal Charges a. For some molecules, you can draw more than one Lewis Structure that gives all the atoms a full octet. When this happens consider it a blend of all options. Ex. CO2 O=C=O or O-C=O i. CO2’s structure is somewhere in between these two, with the more stable form contributing more. ii. Which is more stable? 1. Assign formal charges to each atom in the molecule. The Lewis structure that contributes most is the most stable. a. Generally the structure in which the atoms all have formal charges that are closest to zero is most stable b. For Lewis structures that have charged atoms, the more stable structure will be the one that gives negative charges to the more electronegative atoms. b. How to calculate Formal Charges i. Formal Charge = (# of valence electrons) – (# of electrons assigned to the atom) 1. All lone pairs of electrons are assigned to the atom on which they are found. 2. For any bond, half of the bonding electrons are assigned to each atom in the bond.