The Covalent Bond Model
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1 1 Chapter 5 Chemical Bonding: The Covalent Bond Model Atoms with similar ionization energy and electronegativity DO NOT form ionic bonds. There is NO electron transfer! Electron pairs are shared to form a covalent bond.
2 2 A covalent bond is a chemical bond resulting from two attracting the same shared. Fig 5.1 Electron sharing can occur only when electron orbitals from two different atoms overlap. H + H H : H H H H 2 Each atom in H 2 has the electron configuration of
3 3 Ch 5.2 Lewis Structures for Molecular Compounds H + F : H : F : or H - F : bonding electrons nonbonding electrons F, with valence electrons, forms 1 covalent bond. O, with valence electrons, forms 2 covalent bonds. Nonbonding electrons are also called.
4 4 N valence electron 3 covalent bonds C valence electron 4 covalent bonds The number of covalent bonds formed by a nonmetallic element is directly correlated with the number of electrons it must share in order to obtain an of electrons.
5 5 Ch 5.3 Single, Double, and Triple Covalent Bonds Some atoms must share than one pair of electrons to obtain an octet of electrons. Single bond - 2 atoms share 1 pair of e Double bond - 2 atoms share 2 pairs of e Triple bond - 2 atoms share 3 pairs of e.. : N N : : N ::: N :.. : N N : 1 triple bond each N in N 2 has an
6 6 Carbon dioxide has double bonds : O : C O : : O :: C :: O : : O = C = O : each atom in CO 2 has a complete octet
7 7 Supplemental material Covalent bonds, bond energy, and bond length Single Double Triple bond energy increases bond length ethane ethylene acetylene
8 8 Supplemental material: Sigma and Pi Bonds ethylene double bond = 1 sigma (σ) bond & pi (π) bond σ bond π bond, 2 overlapping p orbitals (2 lobes/orbital) restricted motion acetylene triple bond = 1 σ bond & π bonds
9 9 Ch 5.4 Valence Electrons and Number of Covalent Bonds Formed Group VIIA - 7 valence electrons - 1 covalent bond & lone pairs.. : Cl 1 single bond.. Group VIA - 6 valence electrons - 2 covalent bonds & lone pairs : O : O 2 single bonds 1 double bond
10 10 Group VA - 5 valence electrons - 3 covalent bonds & lone pair : N : N : N 3 single 1 single 1 triple & 1 double Group IV - 4 valence electrons - 4 covalent bonds & lone pairs C C C C 4 single 2 single 2 double 1 single & 1 double &
11 11 Ch 5.5 Coordinate Covalent Bonds In a coordinate covalent bond electrons of a shared pair come of the two atoms in the bond.
12 12 Ammonia shares its with the hydrogen ion. H H + H N : H + H N H H no electrons H 8 electrons 1 coordinate covalent bond NH 3 + H + NH 4 + ammonia ion
13 13 Ch 5.6 Systematic Procedures for Drawing Lewis Structures Bottom Line All valence electrons must be shown. All atoms must have an octet of electrons. (common exceptions include H) Strategy 6 steps (slightly modified here) Example A Ammonia Step 1. Sum up all valence e (adjust for charge if necessary) N in VA = 5 valence e H in IA = 1 valence e 1 x 5e + 3 x 1e = valence e (there is no charge)
14 14 Step 2. Draw skeleton structure and connect atoms by covalent bonds. H H N H NH 3 the central atom is often written first Step 3. Subtract the number of e used in skeletal structure bonds from valence e of step 1. 8 valence e from step 1-6 e used in step 2 (2 e per bond) 2 e left Step 4. Count number of e needed in skeletal structure (step 2) to give each atom an octet (2 e for H). If that number = e left in step 3, distribute them. all H are H N needs 2 e for octet 2 e left in step 3 H N H add as lone pair on
15 15 Example B Hydrogen cyanide HCN (written in order bonded) Step 1 total valence e = = 10 e (no charge) Step 2 Step 3 Step 4 Step 5 Step 6 H C N 10 e available - e used in step 2 6 e left H has 2 e = complete C in step 2 needs 4 e N in step 2 needs 6 e need 10 e to complete all octets We are short e For each 2 e short, share one more pair by forming multiple bonds. Add 2 more bonds. H C N Distribute remaining e to complete octets. 10 e available (step 1) - e used (step 5) 2 e remain H C N :
16 16 Example C (drill problem) Chlorite ion ClO 2 - valence e : 7 + 2(6) + 1 = 20 e - 4 e used 16 e left : O Cl complete this structure : O : emcompleteember the charge! O needs Cl needs e needed
17 17 Ch 5.7 Bonding in Compounds with Polyatomic Ions Present Both ionic and covalent bonds are present..... _ K + : O Cl : : O :
18 18 Supplemental Material Resonance Structures A single Lewis structure does not always adequately represent a substance and the concept of resonance is used to describe the bonding in such molecules. Resonance structures are two or more Lewis structures that represent the same ion or molecule equally well. Examples: ion NO 2 - ozone O 3 ion CO 3 2- benzene
19 19 Ch 5.8 Molecular Geometry is the 3-D arrangement of atoms within a molecule Molecular models of methane
20 20 We can predict the molecular geometry using a molecule s Lewis structure and the Valence Shell Electron-Pair Repulsion theory ( theory) Electrons repel each other Electrons tend to be as far apart as possible Electron pairs control 2 electron pairs or VSEPR groups 3 electron pairs or VSEPR groups 4 electron pairs or VSEPR groups Figure 5.8
21 21 Steps in predicting molecular geometry draw a Lewis structure determine the number of VSEPR groups bonding and nonbonding pairs count equally single, double, and triple bonds count equally as one group predict which arrangement of VSEPR groups minimizes repulsion A. Molecules with two VSEPR groups are linear.... : O = C = O : 2 VSEPR groups around the central atom bond angle
22 22 B. Molecules with three VSEPR groups are either trigonal planar or angular H 2 CO (formaldehyde) : O : H C H 3 VSEPR groups Trigonal planar bond angle ~
23 23 NO nitrite ion : O N O : VSEPR groups around N 2 bonding, 1 nonbonding bent or angular shape bond angle ~120 o
24 24 C. Molecules with four VSEPR groups have three possible geometries: Tetrahedral (no nonbonding e pairs) Trigonal pyramidal ( nonbonding e pair) Angular or bent ( nonbonding e pairs) o 107 o o
25 Molecules with more than one central atom 25
26 26 Ethane 4 VSEPR groups for carbon atom Tetrahedral arrangement around each carbon atom. ball-and-stick model of Ethane
27 27 Ch 5.9 Electronegativity Electronegativity is a measure of the relative attraction that an atom has for the electrons in a bond.
28 28 Ch 5.10 Bond Polarity pure covalent polar covalent ionic Differences in Electronegativity covalent bonds (similar values) < < covalent bonds < < Borderline area < 2.0 (ionic or polar covalent) bonds > 2.0 Depicting bond polarity
29 29 Ch 5.11 Molecular Polarity depends on two factors: Bond polarity Molecular A. Polar bond + unsymmetrical distribution of electronic charge = molecule
30 30 B. Polar bond + symmetrical distribution of electronic charge = molecule Polarity cancels Polarity cancels in a trigonal planar molecule with 3 identical atoms or groups attached. Polarity cancels in a tetrahedral molecule with 4 identical atoms or groups attached.
31 Examples of polar and nonpolar molecules (all contain bonds) NH 3 is trigonal pyramidal, not planar 31
32 32 Ch 5.12 Naming Binary Molecular Compounds Different compounds exist for most pairs of. Examples of N-O compounds: NO NO 2 N 2 O 3 N 2 O 4 N 2 O 5 N 2 O 3 dinitrogen trioxide 1. prefix + full name of least electronegative nonmetal 2. prefix + stem name of more electronegative nonmetal + suffix of ide Table 5.1 Common numerical prefixes for 1-10
33 33 Table 5.2 Accepted Common Names Memorize? Yes Yes Yes No Yes
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