Unit 7: Basic Concepts of Chemical Bonding. Chemical Bonds. Lewis Symbols. The Octet Rule. Transition Metal Ions. Ionic Bonding 11/17/15

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1 Unit 7: Basic Concepts of Chemical Bonding Topics Covered Chemical bonds Ionic bonds Covalent bonds Bond polarity and electronegativity Lewis structures Exceptions to the octet rule Strength of covalent bonds Chemical Bonds Whenever atoms or ions are strongly attracted to one another, we say that there is a chemical bond between them An ionic bond is an electrostatic attraction that exists between ions of different charges A covalent bond results from the sharing of electrons between two atoms A metallic bond is where metal atoms are bound to the surrounding atoms but electrons are free to move throughout the structure The Lewis symbol for an element consists of the chemical symbol for the element plus a dot for each valence e - Sulfur has the e - config. [Ne]3s 2 3p 4, so its Lewis symbol shows 6 valence e - Lewis Symbols The Octet Rule Atoms tend to gain or lose e - to achieve the same number of e - as the nearest noble gas to them Because all noble gases have 8 valence e -, the guideline that atoms tend to gain, lose, or share e - until they have 8 valence e - is referred to as the octet rule An octet of e - consists of full s and p subshells Transition Metal Ions Because of their d orbitals, transition metals often can form multiple ions Fe [Ar]4s 2 3d 6 Can lose 2 e - from 4s Fe 2+ [Ar] 3d 6 Can lose one more from 3d to have a stable half-full sublevel Fe 3+ [Ar]3d 5 Transition metals generally do not form ions with noble-gas configurations, which shows the limitations of the octet rule When sodium atoms transfer electrons to chlorine atoms, crystals of the ionic solid sodium chloride are formed In the sodium chloride crystal, each Na + ion is surrounded by six nearest-neighbor Cl - ions, and each Cl - ion is surrounded by six nearest-neighbor Na + ions 1

2 NaCl is an example of a very typical ionic compound because it consists of a metal with low ionization energy and a nonmetal with high electron affinity Reactions to form ionic compounds from their elements are generally very exothermic A measure of how much stabilization results from arranging oppositely charged ions in an ionic solid is given by the lattice energy Lattice energy is the energy required to completely separate one mole of solid ionic compound into its gaseous ions NaCl(s) Na + (g) + Cl - (g) H lattice = +788 kj/mol The magnitude of the lattice energy of a solid depends on the charges of the ions, their sizes, and their arrangement For a given arrangement of ions, the lattice energy increases as the charges on the ions increase and as their radii decrease Ionic radii do not vary over a very wide range, so the lattice energies depend primarily on the ionic charges Arrange the following ionic compounds in order of increasing lattice energy: NaF, CsI, and CaO 2

3 Which substance would you expect to have the highest lattice energy, AgCl, CuO, or CrN? One of the following pictures represents NaCl and one represents MgO. Which is which, and which has the larger lattice energy? Ionic Compounds There is no way to calculate lattice energies experimentally They can, however, be calculated using Hess s law The sequence of five steps that will be summed is best shown by constructing a Born-Haber cycle*** The Born-Haber Cycle for NaCl*** The Born-Haber Cycle for NaCl 1. Energy to vaporize sodium metal (standard heat of formation) 2. Energy to create monatomic chlorine (standard heat of formation) 3. Energy needed to ionize sodium gas (ionization energy) 4. Energy released when chlorine ionizes (electron affinity) 5. Energy released when gaseous ions form ionic crystal (negative lattice energy) ΔH f [ NaCl( s) ] = ΔH f [ Na( g) ] + ΔH f [ Cl( g) ] + I 1 ( Na) + E a ( Cl) ΔH lattice -411 kj = 107 kj kj kj kj - H lattice H lattice = 787 kj/mol 3

4 Use this data to estimate the lattice energies of NaBr and MgCl 2 In order for covalent bonds to exist, the attractive forces between molecules must exceed the repulsive ones Atoms are held together because the nuclei are mutually attracted to the electrons between them Covalent Bonding Lewis Structures The formation of covalent bonds can be shown using Lewis symbols Shared pairs of electrons are shown as dashes (single, double, or triple) Lewis Structures 3 + or + or 2 or Covalent Bonding For the nonmetals, the number of valence electrons on a neutral atom is the same as the group number You can predict that 7A elements, such as F, would form one covalent bond to have an octet; 6A elements, such as O, would from two bonds; 5A elements, such as N, would form three bonds; and 4A elements, such as C, would from 4 bonds These predictions are true in many compounds, but keep in mind that these are just guidelines and that there are many exceptions Multiple Bonds Sharing a single pair of electrons constitutes a single bond (single line drawn) When two pairs of electrons are shared between two atoms it is a double bond (two lines are drawn) When three pairs are shared it is a triple bond (three lines are drawn) 4

5 Multiple Bonds As a general rule, the distance between bonded atoms decreases as the number of shared electron pairs (bond order) increases Electronegativity The ability of an atom in a molecule to attract electrons to itself is called electronegativity We use electronegativity to estimate whether a given bond will be nonpolar covalent, polar covalent, or ionic Electronegativity is related to electron affinity and follows the same trends across the periodic table Electronegativity Bond Polarity A nonpolar covalent bond is one in which electrons are shared equally between two atoms In a polar covalent bond one of the atoms exerts a greater attraction for the bonding electrons and the sharing is unequal Bond Polarity We can use the difference in electronegativity to gauge the polarity of the bonding between atoms Examples F 2, = 0, nonpolar covalent HF, = 1.9, polar covalent LiF, = 3.0, ionic Which bond is more polar: B-Cl or C-Cl? P-F or P-Cl? Indicate in each case which atom has the partial negative charge 5

6 Dipole Moments Whenever two electrical charges of equal magnitude but opposite sign are separate by a distance a dipole is established The quantitative measure of the magnitude of a dipole is called its dipole moment (µ) Dipole Moments There are three types: A permanent dipole occurs when two atoms in a molecule have substantially different electronegativity; one atom attracts electrons more than another, becoming more negative, while the other atom becomes more positive A molecule with a permanent dipole (accompanied by asymmetry within the molecule) is called a polar molecule An instantaneous dipole occurs due to chance when electrons happen to be more concentrated in one place than another in a molecule, creating a temporary dipole An induced dipole can occur when one molecule with a permanent dipole repels another molecule's electrons, inducing a dipole moment in that molecule We can represent the polarity of molecules in two ways The symbol δ is the lower-case Greek letter delta and denotes partially in chemistry Dipole Moments Dipole Moments Polarity helps determine many of the properties of substances that we observe on the macroscopic level Polar molecules align themselves with each other and with ions These interactions account for many properties of solids, liquids, and solutions Drawing Lewis Structures Indicate the dipole moment on each molecule BeH 2 HBr H 2 O The following guidelines can be used to create Lewis structures for covalent compounds 1. Sum up valence electrons from all atoms For anions, add an electron for each negative charge For cations, subtract an electron for each positive charge 2. Arrange element symbols and connect them with a single bond Formulas are often written in the order that the atoms are connected (HCN) or sometimes the central atom is written first (CO 3 2- ) The central atom is usually the least electronegative 6

7 Drawing Lewis Structures 3. Complete the octets of the atoms bonded to the central atom Remember that hydrogen can only have two electrons 4. Count electrons on the structure, place any leftover on the central atom Do this even if it gives the central atom more than an octet 5. If there are not enough electrons to give the central atom an octet, use unshared pairs of electrons on outer atoms to create double or triple bonds 6. For polyatomic ions, place large brackets around the structure and write the charge outside the brackets on the upper right-hand side 7. Draw all resonance structures when necessary Draw the Lewis structure for HBr Draw the Lewis structure for PCl 3 Draw the Lewis structure for CH 2 Cl 2 Draw the Lewis structure for HCN Draw the Lewis structure for NO + 7

8 Draw the Lewis structure for C 2 H 4 Draw the Lewis structure for BrO 3 - Draw the Lewis structure for ClO 2 - Draw the Lewis structure for PO 4 3- Formal Charge*** Formal charge is the charge that an atom in a molecule would have if all atoms had the same electronegativity (all bonding pairs were shared equally) It is used to determine which of several possible Lewis structures is correct To calculate formal charge, assign electrons to the atom as follows: All of the unshared (nonbonding) electrons are assigned to the atom on which they are found Half of the bonding electrons are assigned to each atom in the bond Subtract from the number of valence electrons found on the neutral atom Formal Charge*** Example CN - (C = -1, N = 0) Notice that the sum of the formal charges always equals the charge on the molecule As a general rule, the most stable Lewis structure will be the one that: The atoms bear the formal charges closest to zero Any negative charges reside on the more electronegative atom 8

9 Using formal charges, determine which of the following structures is more likely: CO 2 with two double bonds, or a single bond and triple bond Using formal charges, determine which of the following structures is more likely: NCS - with two double bonds, or two variations of single bond and triple bond Using formal charges, determine which of the following structures is more likely: NCO - with two double bonds, or two variations of single bond and triple bond Resonance Structures Resonance structures are structures that have the same placement of atoms with different placement of electrons that are equally stable The actual structure is a hybrid mixture with fractional bonds Draw the resonance structures for ozone and determine the bond order for each bond Draw the resonance structures for a nitrate ion and determine the bond order for each bond 9

10 Which has shorter bonds, SO 3 or SO 3 2-? Draw two resonance structures for the formate ion, HCO 2 - Resonance in Benzene Exceptions to the Octet Rule Although they are very rare, some molecules have an odd number of valence electrons to share These molecules break the octet rule Example: NO Exceptions to the Octet Rule Exceptions to the Octet Rule Some elements can have less than an octet in a molecule This is most commonly encountered with B and Be Some elements in the third period and beyond can expand their valence shells by utilizing unoccupied d orbitals This gives them the ability to have more than an octet 10

11 Draw the Lewis structures for SF 4 and ICl 4 - Bond Enthalpy Bond enthalpy is the enthalpy change, ΔH, for the breaking of a particular bond in one mole of gaseous substance Cl 2 (g) à 2Cl(g), ΔH = D(Cl Cl) = 242 kj/mol Always a positive value Multiple bonds (shorter bonds) have higher bond enthalpies Bond Enthalpy Calculating ΔH rxn from Bond Enthalpies The enthalpy change for a reaction, ΔH rxn, can be calculated using bond enthalpies ΔH rxn = Σ[bond enthalpies of bonds broken reactants)] - Σ[bond enthalpies of bonds formed (products)] Using bond enthalpies, calculate the change in enthalpy for the following reactions: CH 4 (g) + Cl 2 (g) à CH 3 Cl(g) + HCl(g) 2C 2 H 6 (g) + 7O 2 (g) à 4CO 2 (g) + 6H 2 O(g) N 2 H 4 (g) à N 2 (g) + 2H 2 (g) 11