Chapter 7 P a g e 1 COVALENT BONDING Covalent Bonds Covalent bonds occur between two or more nonmetals. The two atoms share electrons between them, composing a molecule. Covalently bonded compounds are also called molecular compounds. Structural Formula A structural formula uses lines to represent covalent bonds and shows how atoms in a molecule are connected or bonded to each other. It can also show the geometry of a molecule. The structural formula for H2O2 and CO2 are shown below: CHEMICAL BONDING I: LEWIS MODEL One of the simplest bonding theories is called Lewis theory. Lewis theory emphasizes valence electrons to explain bonding. Using Lewis theory, we can draw models, called Lewis structures. Also known as electron dot structures Lewis structures allow us to predict many properties of molecules. Molecular stability, shape, size, and polarity Gilbert Newton Lewis 1875-1946 Why do Atoms Bond? Chemical bonds form because they lower the potential energy between the charged particles that compose atoms. A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms. To calculate this potential energy, you need to consider the following interactions: Nucleus-to-nucleus repulsions Electron-to-electron repulsions Nucleus-to-electron attractions Valence Electrons and Bonding The column number on the periodic table will tell you how many valence electrons a main group atom has. Valence electrons are held most loosely. Chemical bonding involves the transfer or sharing of electrons between two or more atoms. Because of the two previously listed facts, valence electrons are most important in bonding. Lewis theory focuses on the behavior of the valence electrons.
P a g e 2 Lewis Structures of Atoms In a Lewis structure, we represent the valence electrons of maingroup elements as dots surrounding the symbol for the element. Also known as electron dot structures We use the symbol of the element to represent the nucleus and inner electrons. We use dots around the symbol to represent valence electrons. Pair the first two dots for the s orbital electrons. Put one dot on each open side for the first three p electrons. Then, pair the rest of the dots for the remaining p electrons. Lewis Bonding Theory Atoms bond because bonding results in a more stable electron configuration. More stable = lower potential energy Atoms bond together by either transferring or sharing electrons. Usually, this results in all atoms obtaining an outer shell with eight electrons. Octet rule There are some exceptions to this rule: The key to remember is to try to get an electron configuration like a noble gas. Covalent Bonding: Bonding and Lone Pair Electrons Electrons that are shared by atoms are called bonding pairs. Electrons that are not shared by atoms but belong to a particular atom are called lone pairs. Also known as nonbonding pairs. Single Covalent Bonds When two atoms share one pair of electrons, it is called a single covalent bond. Two electrons One atom may use more than one single bond to fulfill its octet. To different atoms H only duet Double Covalent Bonds When two atoms share two pairs of electrons the result is called a double covalent bond. Four electrons
P a g e 3 Triple Covalent Bonds When two atoms share three pairs of electrons the result is called a triple covalent bond. Six electrons Covalent Bonding: Model versus Reality Lewis theory predicts that the more electrons two atoms share, the stronger the bond should be. Bond strength is measured by how much energy must be added into the bond to break it in half. In general, triple bonds are stronger than double bonds, and double bonds are stronger than single bonds. However, Lewis theory would predict that double bonds are twice as strong as single bonds; the reality is that they are less than twice as strong. Lewis theory predicts that the more electrons two atoms share, the shorter the bond should be. When comparing bonds to like atoms Bond length is determined by measuring the distance between the nuclei of bonded atoms. In general, triple bonds are shorter than double bonds, and double bonds are shorter than single bonds. Polar Covalent Bond Covalent bonding between unlike atoms results in unequal sharing of the electrons. One atom pulls the electrons in the bond closer to its side. One end of the bond has larger electron density than the other. The result is a polar covalent bond. Bond polarity The end with the larger electron density gets a partial negative charge. The end that is electron deficient gets a partial positive charge. Electronegativity The ability of an atom to attract bonding electrons to itself is called electronegativity. Increases across period (left to right) and decreases down group (top to bottom) Fluorine is the most electronegative element. Francium is the least electronegative element. Noble gas atoms are not assigned values. Opposite of atomic size trend The larger the difference in electronegativity, the more polar the bond. Negative end toward more electronegative atom Electronegativity Difference and Bond Type If the difference in electronegativity between bonded atoms is 0, the bond is pure covalent. Equal sharing If the difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent. If the difference in electronegativity between bonded atoms is 0.4 to 1.9, the bond is polar covalent.
P a g e 4 If the difference in electronegativity between bonded atoms is larger than or equal to 2.0, the bond is 100% ionic. Example 6 Classify the bond formed between each pair of atoms as covalent, polar covalent, or ionic. a. Sr and F b. Na and Cl c. N and O d. I and I e. Cs and Br f. P and O
P a g e 5 Writing Lewis Structures of Molecular Compounds 1. Write the correct skeletal structure for the molecule. Hydrogen atoms are always terminal. The more electronegative atoms are placed in terminal positions. 2. Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule. For polyatomic ions, consider the charge. Negative ions add electrons Positive ions subtract electrons 3. Distribute the electrons among the atoms, giving octets (or duets in the case of hydrogen) to as many atoms as possible. 4. If any atoms lack an octet, form double or triple bonds as necessary to give them octets. Move lone electron pairs from terminal atoms into the bonding region with the central atom. Note: Attempt these examples on a separate sheet of paper. Example 7 Write the Lewis structures for the following molecules: CO2, CO, NH3, SCl2, CH3SH (C and S central), HCOOH (both O bonded to C), C2H2, Cl2O, N2H2, N3H8 Example 8 Write the Lewis structures for the following ions: BrO, NO 2, O 2 2, NH 4 +, ClO 3, CO 3 2, CO +, NO 3, ClO 4, OCl Resonance and Formal Charge Lewis theory localizes the electrons between the atoms that are bonding together. Extensions of Lewis theory suggest that there is some degree of delocalization of the electrons; we call this concept resonance. Delocalization of charge helps to stabilize the molecule. When there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures. The actual molecule is a combination of the resonance forms a resonance hybrid. The molecule does not resonate between the two forms, though we often draw it that way. Look for multiple bonds or lone pairs. Example 9 Write the Lewis structures for NO2 and NO3. Include resonance structures.
P a g e 6 Formal Charge Formal charge is a fictitious charge assigned to each atom in a Lewis structure that helps us to distinguish among competing Lewis structures. In a Lewis structure, we calculate an atom s formal charge, which indicates the charge it would have if all bonding electrons were shared equally between the bonded atom. FC = # valence e [nonbonding e + ½ bonding e ] Sum of all the formal charges in a molecule = 0. In an ion, total equals the charge. Evaluating Resonance Structures Better structures have fewer formal charges. Better structures have smaller formal charges. Better structures have the negative formal charge on the more electronegative atom. Example 10 Write a Lewis structure that obeys the octet rule for each molecule or ion. Include resonance structures if necessary and assign formal charges to each atom. SeO 2, CO 3 2, ClO 3 Example 11 Use formal charge to determine which Lewis structure is better: Example 12 In N2O, nitrogen is the central atom and the oxygen atom is terminal. In OF2, however, oxygen is the central atom. Use formal charge to explain why. Example 13 Draw the Lewis structure (including resonance structures) for methyl azide (CH3N3). For each resonance structure, assign formal charges to all atoms that have formal charge. Exceptions to the Octet Rule There are three exceptions to the Lewis model: i. Odd-electron species ii. Incomplete octets iii. Expanded octet
P a g e 7 Odd-Electron Species Molecules and ions with odd number of electrons in their Lewis structures are called free radicals. o They have one unpaired electron. o Radicals are very reactive (unstable) because they want to attain an octet. o They are relatively rare compared to other molecules. For example, nitrogen dioxide, NO2 Incomplete Octets Some elements tend to form incomplete octets. o For example, boron (six electrons instead of eight) and beryllium For BF3, what if we form double bonds to increase the number of electrons around boron? When we assign formal charges to all the atoms, we get the following: What is wrong (if any) with the structure above? One of the ways BF3 can get an octet is through a chemical reaction. It gains electrons to complete its octet. Expanded Octets This means more than 8 electrons, up to 12 (and occasionally 14) electrons. o Elements in the third row (period) and beyond can exhibit expanded octets. o They have energetically accessible d-orbitals to accept the extra electrons. o Consider arsenic pentafluoride and sulfur hexafluoride: Example 14 Write the Lewis structures for XeF2, XeF4, and H3PO4.
P a g e 8 Example 15 Write Lewis structures for each molecule or ion. Include resonance structures if necessary and assign formal charges to all atoms. If necessary, expand the octet on the central atom to lower formal charge. a. PO 4 3 b. I 3 c. AsF 6 d. Cl 3 PO Example 16 Draw the Lewis structure for urea, H2NCONH2, one of the compounds responsible for the smell of urine. Does urea contain polar bonds? Which bond (if any) in urea is most polar? [Hint: The central carbon atom is bonded to both nitrogen atoms and to the oxygen atom]. Example 17 Phosgene (Cl2CO) is a poisonous gas used as a chemical weapon during World War I. It is a potential agent for chemical terrorism today. Draw the Lewis structure for phosgene. Include all three resonance forms by alternating the double bond among the three terminal atoms. Which resonance structure is the best? Explain. CHEMICAL BONDING II: VSEPR THEORY VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY Properties of molecular substances depend on the structure of the molecule. Valence shell electron pair repulsion (VSEPR) theory is a simple model that allows us to account for molecular shape. o Electron groups are defined as lone pairs, single bonds, double bonds, and triple bonds. o VSEPR is based on the idea that electron groups repel one another through coulombic forces. Electron groups around the central atom will be most stable when they are as far apart as possible. We call this VSEPR theory. o Because electrons are negatively charged, they should be most stable when they are separated as much as possible. The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule. The Lewis structure predicts the number of valence electron pairs around the central atom(s). o Each lone pair of electrons constitutes one electron group on a central atom. o Each bond constitutes one electron group on a central atom, regardless of whether it is single, double, or triple. Electron Group Geometry There are five basic arrangements of electron groups around a central atom. Based on a maximum of six bonding electron groups Though there may be more than six on very large atoms, it is very rare. Each of these five basic arrangements results in five different basic electron geometries. In order for the molecular shape and bond angles to be a perfect geometric figure, all the electron groups must be bonds, and all the bonds must be equivalent.
P a g e 9 For molecules that exhibit resonance, it doesn t matter which resonance form you use since the electron geometry will be the same. Two Electron Groups: Linear Geometry When there are two electron groups around the central atom, they will occupy positions on opposite sides of the central atom. This results in the electron groups taking a linear geometry. The bond angle is 180. Three Electron Groups: Trigonal Planar Geometry When there are three electron groups around the central atom, they will occupy positions in the shape of a triangle around the central atom. This results in the electron groups taking a trigonal planar geometry. The bond angle is 120. Four Electron Groups: Tetrahedral Geometry When there are four electron groups around the central atom, they will occupy positions in the shape of a tetrahedron around the central atom. This results in the electron groups taking a tetrahedral geometry. The bond angle is 109.5. Five Electron Groups: Trigonal Bipyramidal Geometry When there are five electron groups around the central atom, they will occupy positions in the shape of two tetrahedra that are base to base with the central atom in the center of the shared bases. This results in the electron groups taking a trigonal bipyramidal geometry. The positions above and below the central atom are called the axial positions. The positions in the same base plane as the central atom are called the equatorial positions. The bond angle between equatorial positions is 120. The bond angle between axial and equatorial positions is 90. Octahedral Electron Geometry When there are six electron groups around the central atom, they will occupy positions in the shape of two square-base pyramids that are base to base with the central atom in the center of the shared bases. This results in the electron groups taking an octahedral geometry.
P a g e 10 It is called octahedral because the geometric figure has eight sides. All positions are equivalent. The bond angle is 90. The Effect of Lone Pairs The actual geometry of the molecule may be different from the electron geometry. Lone pair electrons typically exert slightly greater repulsion than bonding electrons, affecting the bond angles. A lone electron pair is more spread out in space than a bonding electron pair because a lone pair is attracted to only one nucleus while a bonding pair is attracted to two nuclei. In general, electron group repulsions vary as follows: Lone pair lone pair > lone pair bonding pair > bonding pair bonding pair Derivatives of the Tetrahedral Geometry When there are four electron groups around the central atom, and one is a lone pair, the result is called a trigonal pyramidal shape, because it is a triangular-base pyramid with the central atom at the apex. When there are four electron groups around the central atom, and two are lone pairs, the result is called a tetrahedral-bent shape. Consider ammonia, NH3: Consider water, H2O:
P a g e 11 Derivatives of the Trigonal Bypramidal Geometry Lone pairs on central atoms with five electron groups will occupy the equatorial positions because there is more room. The result is called the seesaw shape (aka distorted tetrahedron). When there are two lone pairs around the central atom, the result is T-shaped. When there are three lone pairs around the central atom, the result is a linear shape. The bond angles between equatorial positions are less than 120. The bond angles between axial and equatorial positions are less than 90. Linear = 180 axial to axial. Seesaw T-Shaped Linear Derivatives of the Octahedral Geometry When there are lone pairs around a central atom with six electron groups, each even number lone pair will take a position opposite the previous lone pair. When one of the six electron groups is a lone pair, the result is called a square pyramid shape. The bond angles between axial and equatorial positions are less than 90. When two of the six electron groups are lone pairs, the result is called a square planar shape. The bond angles between equatorial positions are 90. Square Pyramidal Square Planar Representing Three Dimensional Structures on Paper One of the problems with drawing molecules is trying to show their dimensionality. By convention, the central atom is put in the plane of the paper. Put as many other atoms as possible in the same plane and indicate with a straight line. For atoms in front of the plane, use a solid wedge. For atoms behind the plane, use a hashed wedge.
P a g e 12 Example 18 Determine the molecular geometry of NO3. Example 19 Determine the molecular geometry of CCl4. Example 20 Suppose that a molecule with six electron groups were confined to two dimensions and therefore had a hexagonal planar electron geometry. If two of the six groups were lone pairs, were would they be located? a) Positions 1 and 2 b) Positions 1 and 3 c) Positions 1 and 4 Example 21 Predict the geometry and bond angles of PCl3. Example 22 Predict the geometry and bond angles of ICl4. Predicting the Shape of Larger Molecules Many molecules have larger structures with many interior atoms. We can think of them as having multiple central atoms. we describe the shape around each central atom in sequence. Consider the amino acid, glycine:
P a g e 13 Example 23 Predict the geometry about each interior atom in methanol (CH3OH) and make a sketch of the molecule. Predicting Polarity of Molecules Draw the Lewis structure, and determine the molecular geometry. Determine whether the bonds in the molecule are polar. o If there are no polar bonds, the molecule is nonpolar. Determine whether the polar bonds add together to give a net dipole moment. Adding Dipole Moments (Vector Addition) 1. The H Cl bond is polar. The bonding electrons are pulled toward the Cl end of the molecule. The net result is a polar molecule.
P a g e 14 2. The O C bond in CO2 is polar. The bonding electrons are pulled equally toward both O ends of the molecule. The net result is a nonpolar molecule. 3. The H O bond in H2O is polar. Both sets of bonding electrons are pulled toward the O end of the molecule. Because the molecule is bent, not linear, the net result is a polar molecule. Example 24 Determine whether NH3 is polar. Example 25 Determine whether CF4 is polar. Example 26 Determine whether CClF3 is polar. Polarity and Solubility in Water Like dissolves like. Polar molecules are attracted to other polar molecules. Because water is a polar molecule, other polar molecules dissolve well in water. And ionic compounds as well. Water and oil do not mix because water molecules are polar and the molecules that compose oil are generally nonpolar. Some molecules have both polar and nonpolar parts for example, soaps.
P a g e 15 THE VALENCE BOND (VB) THEORY Valence Bond theory (VB) approaches chemical bonding based on an extension of the quantummechanical model (perturbation theory). When orbitals on atoms interact, they make a bond. These orbitals are hybridized atomic orbitals, a kind of blend or combination of two or more standard atomic orbitals. When two atoms approach each other, the electrons and nucleus of one atom interact with the electrons and nucleus of the other atom. If the energy of the system is lowered because of the interactions, a chemical bond forms. When two atoms with half-filled orbitals approach each other, the half-filled orbitals overlap and the electrons align with opposite spins (spin-pair). This results in a net energy stabilization and hence a chemical bond. A bond can also result from the overlap of a completely filled orbital with an empty orbital. The geometry of the overlapping orbitals determines the shape of the molecule. When two atoms approach each other, the electrons and nucleus of one atom interact with the electrons and nucleus of the other atom. If the interaction lowers the energy (negative interaction energy), a chemical bond forms. If the interaction raises the energy (positive interaction energy), a chemical bond does not form. When the atoms are far apart, the energy is nearly zero because they are not interacting. As they get closer, the energy is lowered (becomes negative) and is minimum at the optimal overlap. At this minimum energy, the overlap has the ideal bond length. If the atoms get too close, there will be mutual repulsion of the two positively charge nuclei. Orbital Diagram for H2S
P a g e 16 Hybridization of Atomic Orbitals The overlap of standard half-filled orbitals does not adequately explain the bonding in many other molecules. Consider the bonding between hydrogen and carbon: Based on the electron configurations of carbon and hydrogen, the compound will be CH2 with a bond angle of 90. However, CH2 does not exist! Instead, CH4 is observed. Hybridization is particularly important in carbon, which tends to form four bonds in its compounds and therefore always hybridizes. Hybrid Orbitals The number of standard atomic orbitals combined = the number of hybrid orbitals formed. Combining a 2s with a 2p gives two 2sp hybrid orbitals. H cannot hybridize! Its valence shell has only one orbital. The number and type of standard atomic orbitals combined determines the shape of the hybrid orbitals. The particular kind of hybridization that occurs is the one that yields the lowest overall energy for the molecule. sp 3 Hybridization One s orbital and three p orbitals are mixed. Atom with four electron groups around it. Tetrahedral geometry 109.5 angles between hybrid orbitals Atom uses hybrid orbitals for all bonds and lone pairs. Example, NH3, CH4, H2O sp 2 Hybridization One s orbital + two p orbitals are mixed. Hybrid orbitals will overlap on axis with orbitals from other atoms. Trigonal planar geometry 120 bond angles between hybrid orbitals.
P a g e 17 Unhybridized p orbital will overlap sideways, or side by side, with an unhybridized p orbital of another atom. Sigma Bonds and pi Bonds A sigma (σ) bond results when the interacting atomic orbitals point along the axis connecting the two bonding nuclei. Either standard atomic orbitals or hybrids s to s, p to p, hybrid to hybrid, s to hybrid, etc. A pi (π) bond results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei. Between unhybridized parallel p orbitals The interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore, σ bonds are stronger than π bonds. Bond Rotation Because the orbitals that form the s bond point along the internuclear axis, rotation around that bond does not require breaking the interaction between the orbitals. But, the orbitals that form the p bond interact above and below the internuclear axis, so rotation around the axis requires the breaking of the interaction between the orbitals.
P a g e 18 sp Hybridization Atom with two electron groups; for example, C2H2 Linear shape 180 bond angle Atom uses hybrid orbitals for s bonds or lone pairs and uses nonhybridized p orbitals for p bonds Usually will for two s bonds and two p bonds.
P a g e 19 sp 3 d Hybridization Atom with five electron groups around it. Trigonal bipyramid electron geometry Seesaw, T-shape, linear 120 and 90 bond angles Use empty d orbitals from valence shell. sp 3 d 2 Hybridization Atom with six electron groups around it Octahedral electron geometry Square pyramid, Square planar 90 bond angles Use empty d orbitals from valence shell to form hybrid. Example 27 Write a hybridization and bonding scheme for a. BrF3 b. XeF4 c. HCN d.