Chapter 7 Covalent Bonds and Molecular Structure The

Chapter 7 Covalent Bonds and Molecular Structure

The Covalent Bond • Covalent bond – formed by the sharing of electrons between two nonmetal atoms – Forces involved in the bond • Electrostatic attraction between proton and electron • Electrostatic repulsion between electron and electron – When will a bond form?

Strengths of Bonds • Bond dissociation energy – the amount of energy required to break a bond – Energy increases as the length of the bond gets shorter

Bond length and covalent radius. Internuclear distance (bond length) Covalent radius

Strengths of Bonds – Single bonds > double bonds > triple bonds

Problem • Arrange the following bonds in order of increasing bond strength. – A. – B. – C. – D. – E. C-I < C-Br < C-Cl < C-F < C-Cl < C-Br < C-I < C-Cl < C-F C-I < C-Br < C-F< C-Cl none of these orders is correct

Problem • Select the strongest bond in the following group. – A. – B. – C. – D. – E. C-S C-O C=C C≡N C-F

Electron Dot Structures • Electron dot symbols – Aid in understanding the formation of bonds between atomic nuclei – Elemental symbol represents the type of element and all core electrons; the valence electrons are represented by dots around the symbol

Electron Dot Structures • A metal in an ionic loses its electrons to achieve an octet or pseudo-octet (transition elements) in its outermost shell • A nonmetal in an ionic compound gains electrons to achieve an octet in its outermost shell • Period 1 and 2 elements of a covalent compound share enough electrons to achieve an octet

Electron-Dot Structures • Ionic: • Covalent:

Electron Dot Structures • Using electron dot symbols build – H 2, H 2 O, CH 4, O 2, N 2, HCN, CO 2 • Least electronegative atom is often central (except H)

Naming Binary Molecular Compounds • Electronegativity – indicates how well an elements nuclei attract the electrons in a covalent bond

The Periodic Table and Electronegativity

Electron Dot Structures • Using electron dot symbols build – H 2, H 2 O, CH 4, O 2, N 2, HCN, CO 2 • Least electronegative atom is often central (except H)

Electron Dot Structures • Single bond: A covalent bond formed by sharing one electron pair. • Double bond: A covalent bond formed by sharing two electron pairs. • Triple bond: A covalent bond formed by sharing three electron pairs. • Single bonds are longer (weaker) than double bonds • Double bonds are longer (weaker) than triple bonds

Electron-dot Structures • Step 1: Count the total valence electrons. • Step 2: Identify the central atom - Often least electronegative • Step 3: Place all other atoms around the central atom • Step 4: Draw a single bond between each external atom and the central atom subtracting 2 electrons for each bond drawn from the total valence electrons.

Electron-dot Structures • Step 5: Distribute remaining valence electrons around the external atoms giving the external atoms an octet • Step 6: If valence electrons still remain, place them on the central atom in pairs • Step 7: Verify that each atom has an octet » Hydrogen needs only 2 electrons » Boron needs only 6 electrons

Electron-Dot Structures • Step 8: If the central atom does not have an octet, form a multiple bond by bringing a pair of electrons in from the external atom • Step 9: Calculate formal charge and minimize the formal charge if acceptable » Period 3 elements and greater can have expanded octets if one is necessary to minimize formal charge » Formal charge = # valence electrons for the atom – 1 for every dot on the atom – 1 for every line around the atom

Problems • • BF 3 PF 3 C 2 H 6 I 3+ NH 4+ SO 42 KCl. O 3

Electron-dot Structures and Resonance • How is the double bond formed in O 3? • The correct answer is that both are correct, but neither is correct by itself.

Electron-Dot Structures and Resonance • When multiple structures can be drawn, the actual structure is an average of all possibilities. • The average is called a resonance hybrid. A straight double-headed arrow indicates resonance.

Problem • • S 3 PO 43 CO 32 NO 2

Molecular Shapes: The VSEPR Theory • The approximate shape of molecules is given by Valence-Shell Electron-Pair Repulsion (VSEPR).

Molecular Shapes: The VSEPR Theory Molecular formula Step 1 Lewis structure Step 2 Electron-group arrangement Count all e- groups around central atom (A) – Single, Double and Triple bonds are all counted as 1 e- group Step 3 Bond angles Note lone pairs and double bonds Count bonding Step 4 and nonbonding egroups separately. Molecular shape (AXm. En)

Problem • Determine the shape of the molecules for which Lewis Structures have been developed.

Valence Bond Theory • If, in order for a bond to form, a pair of electrons must be shared, then how does C form molecules with 4 bonds? • Valence Bond Theory – hybrid orbitals

Valence Bond Theory Basic Principle A covalent bond forms when the orbtials of two atoms overlap and are occupied by a pair of electrons that have the highest probability of being located between the nuclei. Themes A set of overlapping orbitals has a maximum of two electrons that must have opposite spins. The greater the orbital overlap, the stronger (more stable) the bond. The valence atomic orbitals in a molecule are different from those in isolated atoms.

Valence Bond Theory Key Points The number of hybrid orbitals obtained equals the number of atomic orbitals mixed. The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed. Types of Hybrid Orbitals sp sp 2 sp 3 d 2

Valence Bond Theory The conceptual steps from molecular formula to the hybrid orbitals used in bonding. Step 1 Molecular formula Step 2 Lewis structure Step 3 Molecular shape and e- group arrangement Hybrid orbitals

Problems • Carbon uses ______ hybrid orbitals in Cl. CN. – A. – B. – C. – D. – E. sp sp 2 sp 3 d 2

The s bonds in ethane. both C are sp 3 hybridized s-sp 3 overlaps to s bonds sp 3 -sp 3 overlap to form a s bond relatively even distribution of electron density over all s bonds

The s and p bonds in ethylene (C 2 H 4) overlap in one position - s p overlap - electron density

The s and p bonds in acetylene (C 2 H 2) overlap in one position - s p overlap -

Polar Covalent Bonds: Electronegativity • Electronegativity – represents the ability of an atom to attract a shared pair of electrons • Higher the EN – the more the electrons in a bond will be pulled toward the atom – Most electronegative atom is F • EN ↓ down a group • EN↑ across a period from left to right w/ few exceptions

Polar Covalent Bonds: Electronegativity

Problem • Which of the following elements is the most electronegative? – A. – B. – C. – D. – E. S Ru Si Te Cs

Problem • Arrange calcium, rubidium, sulfur, and arsenic in order of decreasing electronegativity. – A. – B. – C. – D. – E. S > As > Rb > Ca S > As > Ca > Rb As > S > Rb > Ca As > S > Ca > Rb None of these orders is correct.

Polar Covalent Bonds: Electronegativity • % Ionic Character: As a general rule for two atoms in a bond, we can calculate an electronegativity difference (∆EN ): ∆EN = EN(Y) – EN(X) for X–Y bond. l If ∆EN < 0. 5 the bond is covalent. l If ∆EN 0. 5 - < 2. 0 the bond is polar covalent. l If ∆EN > 2. 0 the bond is ionic.

Problem • Select the most polar bond amongst the following. – A. – B. – C. – D. – E. C-O Si-F Cl-F C-I

Molecular Orbital Theory • The molecular orbital (MO) model provides a better explanation of chemical and physical properties than the valence bond (VB) model. – Atomic Orbital: Probability of finding the electron within a given region of space in an atom. – Molecular Orbital: Probability of finding the electron within a given region of space in a molecule.

Molecular Orbital Theory • Additive combination of orbitals (s) is lower in energy than two isolated 1 s orbitals and is called a bonding molecular orbital.

Molecular Orbital Theory • Subtractive combination of orbitals (s*) is higher in energy than two isolated 1 s orbitals and is called an antibonding molecular orbital.

Molecular Orbital Theory • Molecular Orbital Diagram for H 2:

Molecular Orbital Theory • Molecular Orbital Diagrams for H 2– and He 2:

Molecular Orbital Theory • Additive and subtractive combination of p orbitals leads to the formation of both sigma and pi orbitals.

Molecular Orbital Theory • Second-Row MO Energy Level Diagrams:

Molecular Orbital Theory • Bond Order is the number of electron pairs shared between atoms. • Bond Order is obtained by subtracting the number of antibonding electrons from the number of bonding electrons and dividing by 2.