Name: Hr: 8 Basic Concepts of Chemical Bonding

Transcription

1 Name: Hr: 8 Basic Concepts of Chemical Bonding 8.1 Chemical Bonds, Lewis Symbols, and the Octet Rule State the type of bond (ionic, covalent, or metallic) formed between any two atoms. Explain (using attractions and repulsions) why the formation of a bond lowers the potential energy of a molecule. Determine the number of valence electrons for any atom, and write its Lewis symbol. Recognize when the octet rule applies to the arrangement of electrons in the valence shell for an atom. 8.2 Ionic Bonding Write the Lewis symbol for ions. Explain the principle reason for the stability of ionic compounds is attraction between unlike charges. Explain the relationship between stability and lattice energy. Predict the strength of lattice energies using Coulomb s law. Describe how lattice energies can be determined using a Born-Haber cycle. O b j e c t i v e s Describe a single, double, and triple bond. 8.4 Bond Polarity and Electronegativity Explain the significance of electronegativity and relate electronegativity to an elements position on the periodic table. Predict trends in electronegativity using the periodic table. Predict the relative polarities of bonds (bond type) using the periodic table or electronegativity values. State the positive and negative end of any polar bond. 8.5 Drawing Lewis Structures Write the Lewis structure for molecules and ions containing covalent bonds using the periodic table. Define formal charge as the charge on an atom if all shared electrons are shared equally. 8.3 Covalent Bonding Describe the basis of the Lewis theory, and predict the valence of common nonmetallic elements based on their position on the periodic table. Be able to describe a covalent bond in terms of the sharing of electron density between bonded atoms. Describe the formation of covalent compounds using Lewis symbols. Be able to look at a Lewis structure and determine if it properly fits the Lewis model. Determine the formal charge of any atom in a Lewis Structure and use these formal charges to determine the best arrangements of atoms. State that the best structures have minimal formal charges and the more electronegative atoms have the negative formal charges. Contrast formal charge with oxidation states in which shared electrons are assigned to the more electronegative atom.

2 8.6 Resonance Structures Draw the Lewis structure for molecules and ions that exhibit resonance. 8.7 Exceptions to the Octet Rule Draw the Lewis structure for molecules and ions that have an odd number of electrons (NO and NO2), a deficiency of electrons (H, Be, and B), or expanded octets (e.g., SF6, XeF2, XeF4, IBr3, PCl5). 8.8 Strengths of Covalent Bonds Estimate the enthalpy change of a reaction using bond enthalpies. Explain that this method does not give exactly the same answer as Hess s Law because bond energies are average bond energies that differ slightly from molecule to molecule. Relate bond enthalpy to bond length and strength. Use a bond energy diagram to determine the bond length and bond energy of a bond.

3 Chapter 8 Notes Concepts of Chemical Bonding 8.1 Chemical Bonds, Lewis Symbols, and the Octet Rule Three basic types of bonds o Ionic Electrostatic attraction between ions - o Covalent Sharing of electrons - o Metallic Metal atoms bonded to several other atoms - All chemical bonds form because the electron-proton attraction the potential energy of the system as the bond forms. The distance between the bonded particles is a balance between the attractions of the oppositely charged particles and repulsions of the like charged particles. Lewis Symbols The electrons involved in bonding are called electrons. o Valence electrons are found in the incomplete, outermost shell of an atom. A visual method for showing valence electrons is to use dots around the symbol for the element. o The number of valence electrons available for bonding are indicated by. o These symbols are called Lewis symbols or Lewis electron-dot symbols. The Octet Rule Atoms tend to gain, lose or share electrons until they are surrounded by valence electrons; this is known as the octet rule. o An octet consists of full s and p subshells. o We know that s 2 p 6 is a noble gas configuration. o We assume that an atom is stable when surrounded by eight electrons (four electron pairs). Elements 1-5 are too to form a full octet. Some larger non-metallic (period ) atoms can expand their valence shell to 10 to 12 electrons. 8.2 Ionic Bonding When forming ionic bonds, like when NaCl is formed from sodium and chlorine, chlorine removes an electron from the sodium and the ionic bond forms. The reaction can be shown using Lewis electron-dot symbols as follows: The bracket around chloride emphasizes that all eight electrons are located on the Cl- ion. Energetics of Ionic Bond Formation Lattice energy The energy required to completely separate a of a solid ionic compound into its gaseous ions. o It is a measure of the strength of the attraction that is the principle cause of ionic compound stability. o The energy associated with electrostatic interactions is governed by law: Q1Q 2 E k d 2 Lattice energy, then, increases with the on the ions. It also increases with decreasing of ions.

4 2 Chapter 8 Calculation of Lattice Energies: The Born-Haber Cycle The Born-Haber cycle is a thermodynamic cycle that analyzes lattice energy precisely. Consider a Born-Haber cycle for the formation of NaCl(s) from Na(s) and Cl2 (g). The direct route is: Na(s) + ½Cl 2 (g) dnacl(s) H f = kj Alternatively, we can form: 1. Sodium gas formed, heat of vaporization (108 kj; endothermic), then 2. Chlorine gas is broken into atoms, heat of decomposition (122 kj; endothermic), then 3. Sodium ions form, ionization energy (ionization energy for Na, 496 kj; endothermic), then 4. Chloride ions form, electron affinity (electron affinity for Cl, 349 kj; exothermic), then 5. Form the ionic lattice, opposite of lattice energy (exothermic). 6. NaCl is formed, formation enthalpy (exothermic, kj). o Lattice energy = E 5 o E 6 + E 1 + E 2 + E 3 = (E 4 + E 5 ) o E 5 = E 6 + E 1 + E 2 + E 3 + E 4 o The lattice energy = (410.9 kj) + (108 kj) + (122 kj) + (496 kj) + (-349 kj) = +788 kj These phenomena also helps explain the octet rule. Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies. 8.3 Covalent Bonding In covalent bonds atoms electrons. There are several electrostatic interactions in these bonds: o Attractions between electrons and nuclei o Repulsions between electrons o Repulsions between nuclei Lewis Structures By sharing valence electrons, atoms in a molecular compound can achieve a electron configuration. For nonmetals, the valence electrons in a neutral atom is the same as the number. o Nitrogen, group 5A, has valence electrons and forms bonds. o Carbon, group 4A, has valence electrons and forms bonds. Single covalent bond - sharing of 1 of e - between atoms. Unshared pair - pair of electrons that is not being shared. Multiple Bonds Atoms of C, N, O, and sometimes S can form bonds. o Double covalent bond - sharing of pairs of e - between atoms. o Triple covalent bond - sharing of pairs of e - between atoms.

5 Concepts of Chemical Bonding Bond Polarity and Electronegativity Though atoms often form compounds by sharing electrons, the electrons are not always shared equally. Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the hydrogen end. Nonpolar covalent - sharing of electrons. Polar covalent - sharing of electrons. Electronegativity Electronegativity is the ability of atoms in a molecule to electrons to themselves. On the periodic chart, electronegativity as you go o from left to right across a row (greater ). o from the bottom to the top of a column (greater distance decreases attractive forces). Electronegativity and Bond Polarity The difference in electronegativity between two atoms in a bond determines the bond polarity o 0-0.4: covalent o > : covalent o 2.1: The greater the difference in electronegativity (ED), the more is the bond. 8.5 Drawing Lewis Structures Lewis structures are representations of molecules showing all electrons, bonding and nonbonding. 1. Find the of valence electrons of all atoms in the polyatomic ion or molecule. a. If it is an anion, one electron for each negative charge. b. If it is a cation, one electron for each positive charge. 2. The central atom is the electronegative element that isn t hydrogen. Connect the outer atoms to it by single bonds. 3. Fill the octets of the atoms. 4. Fill the octet of the atom. 5. If you run out of electrons before the central atom has an octet a. form bonds until it does. PCl 3 6. If you have extra electrons, place them as pairs on the atom. Formal Charge In some instances, there are several different Lewis dot structures that all obey the octet rule (we re not talking resonance that comes next). How do we know which is correct?

6 4 Chapter 8 Assign formal charges to determine which is correct. o Count the number of electrons. o Add up the number of electrons to half of the number of (or bonding) electrons. o this value from the number of valence electrons. The best Lewis structure o is the one with the charges. o puts a charge on the most electronegative atom. 8.6 Resonance Structures This is the Lewis structure we would draw for ozone, O 3. But this is at odds with the true, observed structure of ozone, in which o both O-O bonds are the length. o both outer oxygens have a charge of -1/2. Lewis structure cannot accurately depict a molecule like ozone. We use multiple structures, resonance structures, to describe the molecule. In truth, the electrons that form the second C-O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. o They are not localized; they are. The organic compound benzene, C 6 H 6, has two resonance structures. It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring. 8.7 Exceptions to the Octet Rule There are three types of ions or molecules that do not follow the octet rule: o Ions or molecules with an number of electrons o Ions or molecules with than an octet o Ions or molecules with than eight valence electrons (an expanded octet)

7 Concepts of Chemical Bonding 5 Odd Number of Electrons Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons. o NO and NO 2 Less than an Octet Deficit Octet Elements 1-5 are stable with less than 8 electrons. o H: only 2 electrons ( bond) o Be: only 4 electrons ( bonds) o B: only 6 electrons ( bonds) o Consider BF 3 : By formal charge, giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine (F is more electronegative). This would not be an accurate picture of the distribution of electrons in BF 3. Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons. The lesson is: if filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don t fill the octet of the central atom. More than an Octet Expanded Octet This is the largest class of exceptions (and often easiest to see does it have more than 4 atoms on the outside?). o Atoms from the period (and below) can accommodate more than an octet. Examples: PCl 5, SF 4, AsF 6, and ICl 4. Elements from the third period and beyond have unfilled d orbitals that can be used to accommodate the additional electrons. o Size also plays a role. The the central atom, the larger the number of atoms that can surround it. o Minimal electron-electron repulsion. The size of the surrounding atoms is also important. Expanded octets occur often when the atoms bound to the central atom are the smallest and most electronegative (e.g., F, Cl, O). The only way PCl 5 can exist is if phosphorus has electrons around it. It is allowed to expand the octet of atoms on the 3 rd row or below. Presumably d orbitals in these atoms participate in bonding.

8 6 Chapter 8 Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens. o This eliminates the charge on the phosphorus and the charge on one of the oxygens. o The lesson is: when the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges, do so. 8.8 Strengths of Covalent Bonds Most simply, the strength of a bond is measured by determining how much energy is required to the bond. This is the bond enthalpy. The bond enthalpy for a Cl-Cl bond, D(Cl-Cl), is measured to be 242 kj/mol. This table lists the average bond enthalpies for many different types of bonds. Average bond enthalpies are positive, because bond breaking is an process. NOTE: These are bond enthalpies, not absolute bond enthalpies; the C-H bonds in methane, CH 4, will be a bit different than the C-H bond in chloroform, CHCl 3. Bond Enthalpies and the Enthalpies of Reactions Yet another way to estimate H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed. In other words, o H rxn = (bond enthalpies of bonds ) - (bond enthalpies of bonds ) Isn t this: reactants products Don t we always do products reactants? We aren t really doing ΔE = reactants products We are actually adding up all of the energies, and the energies on the bond enthalpy table are for bond. Bond forming should be and we are just making them negative when we subtract!

9 Concepts of Chemical Bonding 7 Bond Enthalpy Problem CH 4(g) + Cl 2(g) CH 3 Cl (g) + HCl (g) In this example, one C-H bond and one Cl-Cl bond are broken; one C-Cl and one H-Cl bond are formed. So, H = [D(C-H) + D(Cl-Cl)] - [D(C-Cl) + D(H-Cl)] = = = Estimating Reaction Enthalpies using Bond Enthalpies 1. It often helps to draw the structure because it is easier to see all of the bonds of each type that are formed and broken. 2. Multiply the of each bond times its bond energy. 3. Total the reactants and products 4. Energy = - Bond Enthalpy and Bond Length As number of bonds increases o the bond length. o the bond enthalpy. What is the bond length and bond energy for the H 2 molecule? If this diagram represents formation of C-C bond, draw the C=C bond formation on the same graph.