Topic 4: Chemical Bonding 4.0 Ionic and covalent bonds; Properties of covalent and ionic compounds 4.1 Lewis structures, the octet rule. 4.2 Molecular geometry: the VSEPR approach. Molecular polarity. 4.3 Valence bond theory, hybridization and geometry, multiple bonds. 4.4 Molecular orbital theory.
Types of Bonding: Covalent o Sharing of electrons o Discrete molecular species o Generally nonconductive, Low boiling point o Usually between a two or more non-metals Ionic o Discrete Anions and Cations held together by Electrostatic Attraction o Empirical Formula only / No discrete units o High Boiling points o Usually between one or more metal and one or more non-metal Metallic o Delocalized bonding / electron sharing o Low to high boiling points o Malleable
Key Concept: Coulomb s Law: Describes the energy of interaction between charged species. V = Q 1 Q 2 4πε 0 r Q 1 Q 2 = 2.31 x 10-19 J nm ( ) r Where: V is in Joules r is the distance between the charge centers in nm Q 1 and Q 2, are the numerical unit charges ε 0 is a constant called the Vacuum Permittivity For Example: for Sodium Chloride V = 2.31 x 10-19 J nm ( ) (+1) (-1) 0.276 nm = - 8.37 x 10-19 J Note: negative signs indicate an attraction, ion pair has less energy.
Covalent Bonds are the result of sharing of electrons, usually result in a completion of unfilled electronic shells or in other words becoming isoelectronic with a noble gas.
Lewis Dot symbols and Valance electrons: Lewis Dot Symbols are a way of showing valence e- symbolically. Valence e- are those electrons outside of the filled [Noble gas core] electrons. Except for Helium, the number of valence electrons correspond to the group number. H He F Cl O S IA THE PERIODIC TABLE OF THE ELEMENTS VIII 1 H 1.008 IIA VIIIA Group Number IIIA IVA VA VIA VIIA 3 Li 6.941 11 Na 22.99 19 K 39.10 4 Be 9.012 26 Fe 55.85 Atomic Number Symbol* Atomic Weight *Synthetic elements are hollow faced. The most stable isotope is shown. 12 Mg 24.31 IIIB IVB VB VIB VIIB VIIIB IB IIB 20 Ca 40.08 21 Sc 44.96 22 Ti 47.88 23 V 50.94 24 Cr 52.00 25 Mn 54.94 26 Fe 55.85 27 Co 58.93 28 Ni 58.69 29 Cu 63.55 30 Zn 65.38 5 B 10.81 13 Al 26.98 31 Ga 69.72 6 C 12.01 14 Si 28.09 32 Ge 72.59 7 N 14.01 15 P 30.97 33 As 74.92 8 O 16.00 16 S 32.06 34 Se 78.96 9 F 19.00 17 Cl 35.45 35 Br 79.90 2 He 4.003 10 Ne 20.18 18 Ar 39.95 36 Kr 83.80
Ionic Compounds and the Ionic Bond: An ionic bond is the electrostatic force that holds ions together in an ionic compound. The most common elements to form ionic compounds are Group 1A and 2A metal atoms (cations) and group 6A and 7A (anions). Example: Consider the following Reaction Li + F Li + + - F 1s 2 2s 1 1s 2 2s 0 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6
Covalent Bonding: H + H H H H H A Covalent bond occurs when two electrons are shared by two atoms. A Covalent Compound is one that contains only covalent bonds. F + F F F In this case in the molecule F 2 there is one covalent bond and each F has 3 pairs of lone pairs. A Lewis Structure is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms. Only valence electrons are shown in a Lewis structure
Covalent Bonding and Lewis Structures: Lewis Octet rule: An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons. In other words, a covalent bond forms when there are not enough electrons for each individual atom to have a complete octet. In a single bond, two atoms are held together by one electron pair. Many compounds are held together by multiple bonds, that is, bonds formed when two atoms share two or more pairs of electrons. A double bond occurs when two atoms share two pairs of electrons. A triple bond when two atoms share three pairs of electrons. For Example: CO 2 and C 2 H 4
For Example: N 2 and C 2 H 2 Multiple bonds are shorter than single covalent bonds. Bond length is defined as the distance between the nuclei of two covalently bonded atoms in a molecule For a given pair of atoms triple bonds are shorter than double bonds which are shorter than single bonds The shorter multiple bonds are also more stable.
Comparison of Ionic and Covalent Properties:
Electronegativity: the ability of an atom to attract toward itself the electrons in a chemical bond Elements with high electronegativity have a greater tendency to attract electrons than do elements with low electronegativity. Electronegativity is related to electron affinity and ionization energy. Fluorine has a high electron affinity and a high ionization energy and thus has a high electronegativity. Sodium has a low electronegativity Electronegativity is a relative concept. An element's electronegativity can be measured only in relation to the electronegativity of other elements. En > 2 Ionic compound En < 2 Covalent
Electronegativity Differences result in polar molecules:
Lewis Structure Corrections; PCl 5 (40 VE), Cl Cl Cl P Cl Cl SF 4 (34 VE) F F S F F
VSEPR: Valance Shell Electron Pair Repulsion: The Structure around a given atom is determined principally by minimizing electron pair repulsion. Generic Structure type: Use letter A for the central atom, X for any atoms associated with it and E for any electron pairs. Examples: F Be F O C O C N - Be (AX 2 ); F (AXE 3 ) C (AX 2 ); O (AXE 2 ) C (AXE); N (AXE)
Classifications: Type Structure Orbital Hybrid Angle(s) AX 2 ; AXE Linear sp 180 Type Structure Orbital Hybrid Angle(s) AX 3 Trigonal Planar sp 2 120 AX 2 E Bent sp 2 <120 AXE 2 Linear sp 2 180 Type Structure Orbital Hybrid Angle(s) AX 4 Tetrahedral sp 3 109.5 AX 3 E Trigonal Pyramidal sp 3 <109.5 AX 2 E 2 Bent sp 3 <<109.5 AXE 3 Linear sp 3 180 Type Structure Orbital Hybrid Angle(s) AX 5 Trigonal Bipyramidal sp 3 d 90 & 120 AX 4 E See-Saw sp 3 d 90 & 120 AX 3 E 2 T- shaped sp 3 d 90 AX 2 E 3 Linear sp 3 d 180 AXE 4 Linear sp 3 d 180
Classifications (continued): Type Structure Orbital Hybrid Angle(s) AX 6 Octahedral sp 3 d 2 90 AX 5 E Square Pyramidal sp 3 d 2 90 AX 4 E 2 Square Planar sp 3 d 2 90 AX 3 E 3 T-Shaped sp 3 d 2 90 AX 2 E 4 Bent sp 3 d 2 <90 AXE 5 Linear sp 3 d 2 180
AX 2 ; AXE AX 3 AX 2 E AXE 2 AX 4 AX 3 E AX 2 E 2 AXE 3
AX 5 AX 4 E AX 3 E 2 AX 2 E 3 AXE 4 AX 6 AX 5 E AX 4 E 2 AX 3 E 3 AX 2 E 4 AXE 5
Assignment: Draw all the Example Lewis Structures listed with Formal Charge.
Topic 5: Chemical Bonding 5.0 Ionic and covalent bonds; Properties of covalent and ionic compounds 5.1 Lewis structures, the octet rule. 5.2 Molecular geometry: the VSEPR approach. Molecular polarity. 5.3 Valence bond theory, hybridization and geometry, multiple bonds. 5.4 Molecular orbital theory.
Dipole Moments and Polarity: In molecules containing atoms with different electronegativities, there is an unequal sharing of electron density. In other words, a polar covalent bond is formed. This can be represented by: H F δ + δ - H F Dipole Moment (µ): Is a quantitative measure of polarity. Magnitude of charge Q multiplied by the distance between charges. µ = Q x r µ is normally expressed in Debye units (D) 1 D = 3.336 x 10-30 C m
Dipole Moments and Polarity: Dipoles are vector quantities that are additive to produce an overall molecular dipole moment. For Diatomic molecules from atoms of different electronegativity there is always a dipole moment. For polyatomic molecules the resultant dipole moment is a vector sum of the individual bond dipoles. As a result not all molecules with polar covalent bonds have a net molecular dipole. O C O CO 2 is a non-polar molecule EN: O (3.5); C (2.5) N (3.0); H (2.1); F (4.0)
Dipole Moments and Polarity: Other Examples: C 2 Cl 2 H 2 (connectivity ClHC 2 HCl)
Molecular Orbital Theory: The previous MO model, Valence Bond Theory, does not account for all of the observed properties of molecules, for example the paramagnetic characteristics of Oxygen. Rather than considering an interaction of Atomic Orbitals centered on atoms, MO theory develops a new set of molecular orbitals. For H 2 formation: In a sigma molecular orbital, σ, (bonding or antibonding) the electron density is concentrated symmetrically around a line between the two nuclei of the bonding atoms. A single covalent bond (such as H-H or F-F) is almost always a σ bond.
Bond Order: Fundamentally, a molecule forms because it has lower energy than the separated atoms. In the simple MO model this is reflected by the number of bonding electrons (those that achieve lower energy in going from the free atoms to the molecule) versus the number of antibonding electrons (those that are higher in energy in the molecule than in the free atoms). If the number of bonding electrons is greater than the number of antibonding electrons in a given molecule, the molecule is predicted to be stable. The quantitative indicator of molecular stability (bond strength) for a diatomic molecule is the bond order: the difference between the number of bonding electrons and the number of antibonding electrons, divided by 2. Bond Order = (# Bonding Electrons) (# Antibonding Electrons) 2
Combinations of p-orbitals: Consider the shape of p orbitals and how they might be able to combine.
In a pi molecular orbital (bonding or antibonding), the electron density is concentrated above and below an imaginary line joining the two nuclei of the bonding atoms. Two electrons in a p molecular orbital form a pi bond
Summary of Diatomic Molecular orbital Formation:
Measuring Paramagnetic Properties:
Rules Governing Molecular Electron Configuration and Stability: The number of molecular orbitals formed is always equal to the number of atomic orbitals combined The more stable the bonding molecular orbital, the less stable the corresponding antibonding molecular orbital The filling of molecular orbitals proceeds from low to high energies. In a stable molecule, the number of electrons in bonding molecular orbitals is always greater than that in antibonding molecular orbitals Each molecular orbital can accommodate up to two electrons with opposite spins in accordance with the Pauli exclusion principle Electrons enter molecular orbitals with parallel spins (Hund's rule) The number of electrons in the molecular orbitals is equal to the sum of all the electrons on the bonding atoms Bond Order = ½ ((# e- in Bonding MO s)-( # e- in Antibonding MO s))
Examples: Comment on Stability of H 2 +, H 2, He 2 +, He 2 Recall:
Heterogenous Diatomic Molecular Orbitals: NO (11 VE) NO + or CN - (10 VE)