Chapter 9 - Covalent Bonding

Transcription

1 Chapter 9 - Covalent Bonding I. Covalent Bonding: attractive force produced as a result of shared electrons. (pgs ) A. A is formed when two or more atoms bond covalently. B. Bonds form when there is a balance of attractive and repulsive forces between two atoms: C. Single Covalent Bond & Lewis Structures Single Covalent Bond also called bond (σ ) Orbitals overlap end to end, concentrating the electrons in a between the two atoms One pair of electrons is shared between the two atoms D. Multiple Covalent Bond & Lewis Structures Sharing than of electrons Carbon, nitrogen, oxygen, and sulfur most often form bonds Shared pairs may be needed to form a complete Double bond = two pairs of electrons shared Triple bond = three pairs of electrons shared The pi bond (П) forms when parallel p orbitals overlap. A pi bond always accompanies a sigma bond. For example: O= O For example: N=N This diagram shows how DOUBLE bonds form. First, a sigma (σ) bond forms (electron clouds overlap end-to-end). Then, two p orbitals overlap side-to-side to form a pi (П) bond. These diagrams show the same molecule above simplified. That strange looking shape is what a pi bond actually looks like! 1

2 Therefore, a DOUBLE BOND consists of sigma and pi bond. Simple structural formulas of the compound above: H 2 C = CH 2 or This diagram represents how triple bonds form. You need sigma and pi bonds to form a TRIPLE BOND.Simple structural formula of this compound: HC CH E. Strength of Covalent Bonds Strength of the covalent bond depends on how much separates the nuclei The distance between the two bonding nuclei at the position of maximum attraction is called For atoms to bond covalently, the distance between them must be just right ; not too close (electron clouds repel) or too far apart (no attractive forces between them). According to this diagram, what is the optimum distance for two hydrogen atoms to form a covalent bond? Bond length is determined by of the atoms and electron pairs are shared. Which molecule has the greatest bond length? F 2, O 2, N 2? Which molecule has the greatest bond strength? F 2, O 2, N 2? 2

3 It takes energy to break a bond. Bond dissociation energy is the amount of energy required to a specific covalent bond (usually kj/mole). Compare the energies of single, double and triple bonded carbon atoms. What do you conclude? Chemical reactions involve both breaking original bonds and the formation of new bonds. Bond formation RELEASES ENERGY. Exothermic reaction (feels HOT; net release of energy) occurs if more energy is released when new bonds form than was required to break the original bonds. Endothermic reaction (feels cold; net absorption of energy) occurs when more energy is required to break the original bonds than is released when new bonds form. Refer to this diagram H 1 is energy needed to break the bonds. H 2 is energy released when bonds form. Which is greater, H 1 or H 2? What kind of reaction is pictured here? Circle the correct answer: Endothermic or Exothermic II. Naming molecules (pgs ) A. Rules for writing formulas 1. Represent each kind of element in a compound with the correct symbol 2. Use subscripts to indicate the number of atoms of each element in the compound (If there is only one atom of a particular element, no subscript is used) 3. Write the symbol for the more metallic element first B. Naming binary molecular compounds (made of two nonmetals) REFER TO SEPARATE PRACTICE WORKSHEET! 1. Name the elements in the same order that they appear in the formula 2. Drop the last syllable (two syllables in some cases) in the name of the final element and add - ide. 3. add prefixes to the name of each element to indicate the number of atoms of that element in the molecule. (mono- is frequently omitted for the first element) 3

4 Do you know your prefixes? mono- (1), di- (2), tri-(3), tetra- (4), penta- (5), hexa- (6), hepta- (7), octa- (8), nona- (9), deca (10) examples: N 2 O = N 2 O 5 = NO = N 2 O 3 = C. Naming acids (REFER TO PRACTICE WORKSHEET) Anion ending Example Rule Acid name -ate HNO3 NO3 1- (nitrate) (stem)-ic acid nitric acid -ite H2SO3 SO3 2- (sulfite) (stem)-ous acid sulfurous acid -ide HCl Cl 1- (chloride) hydro-(stem)-ic acid hydrochloric acid III. Molecular Structures (pgs ) A. REVIEW (see page 140 in Chapter 5) Electron Dot Structures consist of the element s symbol (which represents the nucleus and core electrons) surrounded by dots which represent the atom s valence electrons. PRACTICE DRAWING SOME! Li B N F Be C O Ne 4

5 B. There are a variety of ways to represent molecules: Ball and stick model Space filling model C. Maybe we should take a time out, and review the rules for drawing Lewis Structures. (Refer to various Lewis structures worksheets.) How to draw Lewis structures : 1. Calculate the total number of electrons. 2. Arrange the atoms ; show valence electrons as. 3. The number of electron dots must exactly the number of valence electrons calculated in # Change each pair of shared electrons to. 5. Most atoms except and follow the 6. If single bonds don't work, try or bonds. 7. Be aware of exceptions to the octet rule. 8. If carbon is in the formula, it generally acts as the central atom to which all other atoms are attached. OR, it forms chains (or sometimes rings) of carbons all attached to each other, and then the other atoms attach to the carbon chain or ring. 9.Hydrogen is always a terminal (or end) atom. It can be attached to ONLY one other atom! 1. Some equations showing how simple COVALENT Lewis structures form: H Cl + + H Cl H 2 + Cl 2 2HCl Br 2 H 2 O 5

6 2. These are Lewis structures for compounds with MULTIPLE BONDS: Example O2 Oxygen has valence electrons, total Example N2 Nitrogen has valence electrons, total O O N N 3. How do Lewis structures for IONIC compounds differ? Na metal + Cl NaCl 4. Polyatomic ions are groups of atoms covalently bonded but with an overall charge. So, you can use Lewis Structures to represent these ions too! D.Resonance Structures Sometimes MORE than one Lewis structure is needed to represent a molecule or a polyatomic ion. This happens when there are both single AND double bonds Drawing the structure. Lewis The position Structures of the atoms does not change, but the location of the double bond does. (see below) You may need to draw two or more structures to correctly represent the formula, and separate each of them with this symbol: Just like yellow and blue make green, the Resonance actual structure Structures is really an average Example: of all of its in possible ozone resonance the extreme structures. possibilities The bond have lengths one are shorter than double a single and one bond single but longer bond. than The a double resonance bond. structure The true bond length is an average of all of the bonds in the resonance structures. has two identical bonds of intermediate character. O O Resonance in Benzene Drawing O We O write Lewis resonance Structures Ostructures for benzene in which there are single bonds between Oeach pair of C atoms and Common examples: O 3, NO 3-, SO 2- the 6 additional electrons 4, NOare 2, and delocalized benzene. over the entire Resonance in Benzene ring: Benzene consists of 6 carbon atoms in a hexagon. Each C atom is attached to two other C atoms and one hydrogen atom. There are alternating double and single bonds between the C atoms. Experimentally, the C-C bonds in benzene are all the same length. Practice Experimentally, one here: benzene is planar. Drawing Lewis Structures Benzene belongs to a category of organic molecules This is the shorthand called aromatic compounds (due to their odor). or method of drawing benzene 6

7 There are three classes of exceptions to the octet rule: Molecules with an odd number of electrons; Molecules in which one atom has less than an octet; Molecules in which one atom has more than an octet. E. Exceptions to the Octet Rule 1.Odd number of electrons Odd Number of Electrons Few examples. Generally molecules such as ClO 2, NO, and NO 2 have Exceptions an odd number of to electrons. the Octet N O N O Rule 2.Less than Prentice an Hall octet 2003 (watch for elements Chapter 8 like Boron or Beryllium) Less than an Octet Relatively rare. Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. Most typical example is BF 3. Exceptions to the Octet 3.More Formal than an charges octet (also indicate called that expanded the Lewis octet ) structure with an Watch incomplete for elements octet like is S, more P, Cl, important Rule I, Xe. than the ones with double bonds. More than an Octet This is the largest class of exceptions. Atoms from the 3 rd period onwards can accommodate more than an octet. Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density. V M IV. Molecular shape (pgs ) A.The molecular geometry or shape of a molecule can be determined from the Lewis structure of a molecule. The model (abbreviation for Valence Shell Electron Pair Repulsion ) is used to determine the shape of the molecule. Im plain English, this model simply says that atoms align themselves in a molecule to minimize the repulsive forces of their electron clouds. B.These are the seven molecular shapes you must memorize (page 260): 7

8 Bonding Domains Nonbonding Domains Molecular Geometry Example C.Atomic orbitals can mix together in a process called New orbitals form. Hybrid orbitals have different shapes and energies than the original orbitals that were combined The number of hybrid orbitals formed is the same as the number of original atomic orbitals that were mixed Formation of hybrid orbitals provides new bonding opportunities and helps explain molecular geometries. 8

9 sp hybrid orbitals form when one and one orbital are combined to form lopsided sp hybrid orbitals. When shown together, the two new orbitals are 180 to each other. So, sp hybridization results in molecular geometry sp 2 hybrid orbitals form when one s and TWO orbitals are combined to form lopsided sp 2 hybrid orbitals. When shown together, the two new orbitals are 120 to each other. So, sp 2 hybridization results in molecular geometry sp 3 hybrid orbitals form when one s and THREE orbitals are combined to form lopsided sp 3 hybrid orbitals. When shown together, the three new orbitals are about to each other. So, sp 3 hybridization creates molecular geometry. 9

10 V. Electronegativity and Polarity A. Electronegativity - measure of the electron-attracting power of an atom when it bonds covalently. An element with large electronegativity does not share electrons equally in a covalent bond. In a molecule, the atom with the greatest electronegativity will hog the electrons closer to its side of the molecule(pgs ) In this picture, the red color indicates that the electrons are more likely to be found on the atom on the right because that element is more electronegative. The electronegativity scale below ranks the elements electronegativities: Explain the general trend of electronegativity on the periodic table shown below: Highest = F, 4.0 Lowest = Cs, 0.7 metals-low electronegativity nonmetals-high electronegativity inert gases-not assigned because they don't bond Differences in electronegativities are used to determine the polarity of bonds. Simply find the mathematical difference in the bonding atoms electronegativities F F = NONPOLAR electronegativity difference =? = 0 Nonpolar Covalent Bonds: Electronegativity of the atoms are similar; electron attracting power is fairly equal (for example in diatomic molecules). Little or no difference in electronegativity means atoms share electrons fairly evenly between them as shown in the picture above. What are the seven diatomic molecules? Write them (remember BrINClHOF?): H F Bond = POLAR Polarity COVALENT and (a dipole) Electrons are still being shared but unevenly. Electronegativity A larger difference in electronegativity creates a POLAR bond. There is greater electron density on the F than the Electronegativity and Bond Polarity H (4.0 > 2.1). This means the electron density is greater by the Fluorine in the molecule below. The Fluorine end of an HF molecule is partially negative (note the symbols) and the Hydrogen end is partially positive. The positive end (or Another pole) name for in this a is a polar DIPOLE and bond it can also is be represented symbolized like this: 10 Bond Polarity and Electronegativity Prentice Hall 2003 Bond P Electr Electronega There is no sharp distinction b The positive end (or pole) in + and the negative pole -. Electronegativity Electronegativity and Bond Polarity and Bond Polarity There is no sharp distinction between bonding types. + and the negative pole -. In symmetrical molecules like CO 2, the individual bonds can be polar but because they point in opposite directions, they essentially cancel each other out. The molecule OVERALL acts as if it Hall 2003 is NONPOLAR. Chapter 8 Bond Polarity and Electronegativity is no sharp distinction between bonding types. There is no sharp distinction between bonding types. The positive (or pole) in a polar bond is represented ositive + and the end negative pole (or -. pole) in a polar bond is represented d the negative pole -. The arrow points to the more electronegative element (the one that hogs the electrons closer to itself) Chapter