Lecture Notes on Chemical Bonding - General Chemistry | CHEM 152, Study notes of Chemistry

Download Lecture Notes on Chemical Bonding - General Chemistry | CHEM 152 and more Study notes Chemistry in PDF only on Docsity! Lecture 21: Chemical Bonding • Reading: Zumdahl 13.1-13.3 • Outline – Types of Chemical Bonds – Electronegativity – Bond Polarity and Dipole Moments Chemical Bonds • In broad terms, a chemical bond is a term used to characterize an interaction between two atoms that results in a reduction in energy for the system relative to the isolated atoms. • The degree of energy reduction or “stabilization” is given by the energy required to break the bond (known as the “bond energy”) Coulomb’s Law • For charges with the opposite sign, V is negative and becomes more negative as r becomes smaller, i.e. its an attractive force. • For charges with the same sign V is positive and becomes larger (more positive) as r becomes smaller, i.e. its a repulsive force. • Red: two electrons (repulsion). • Blue: one electron and one proton (attraction). Coulomb’s Law • Electrons are attracted to positively charged nuclei. • If an electron on one atom can simultaneously get close to the nuclei of another atom, that is often favorable, i.e. lower energy. • If the electron is between the nuclei, the attraction of the electron to the two nuclei can be more important than the repulsion between the nuclei. Chemical Bonds (cont) • The electrons in the highest energy orbitals have the most to gain by being shared between two atoms. • So the highest energy electrons (valence electrons) of an atom will form bonds. • The atomic orbitals involved in bonding will be those occupied by those highest energy electrons, and the unoccupied orbitals of lowest energy. Chemical Bonds (ionic) • Ionic bonding example: NaCl (Sodium Chloride) • When NaCl is heated to the point of melting, one can demonstrate that the resulting fluid conducts electricity. • This observation demonstrates that the solution (molten NaCl) contains charged species. Those species are Na+ and Cl−. Chemical Bonds (ionic) • The melting point of NaCl is 801ºC. • Demonstrate conduction in molten salt using a mixture of NaNO2 and KNO3, which has a lower melting point. • Operational definition: any compound which conducts electricity when melted is an ionic compound Chemical Bonds (ionic) • This example was for a collection of NaCl dimers in the gas phase….solid NaCl is a bit different. = Na+ = Cl- • Each ion is surrounded by six ions of the opposite charge. Ionic vs. Covalent Bonding • NaCl is an example of “ionic bonding”. In this case an amount of charge approaching that for an electron is “transferred” from one atom to its bonding partner. • Ionic bonding is one limit in the spectrum of bonding. The second limit is a bond in which electrons are “shared” rather than being transferred. • Electrons are shared in a covalent bond. Chemical Bonds (covalent) • The H-H example we saw previously is an example of covalent bonding. For H-H the electrons are shared equally by the two atoms. Electronegativity • Electronegativity can be defined in many ways. Pauling model is the most widely used. • Idea: compare the bond energy of an “HX” molecule to that of the average of an HH bond and an XX bond: Expected energy = [(H-H energy)(X-X energy)]1/2 ∆ = (H-X)experimental − (H-X)expected ∆ > 0: ionic character∆ = 0: covalent Electronegativity (cont.) • Pauling used this approach to develop a scale, where F = 4.0 (flourine has largest electronegativity). F = 4 Cl = 3.2 O = 3.4 C = 2.6 H = 2.2 Na = 0.9 Electronegativity (cont.) • The key idea is this: the greater the electronegativity difference between two atoms, the more ionic the bond. • Example: Which of the following compounds is expected to demonstrate intermediate bonding behavior (i.e., polar covalent). Cl-Cl O-H 1.2 Na-Cl 0 2.3∆elect Dipole Moments (cont.) • When the centers of negative and positive charge are separated, we say that the molecule has a dipole moment. Dipole Moments (cont.) • The dipole moment (µ) is defined as: µ = QR Charge magnitude Separation distance R + center Dipole Moments (cont.) • The units of dipole moment are generally the Debye (D): 1 D = 3.336 ×10−30 C.m • Example, the dipole moment of HF is 1.83 D. What would it be if HF formed an ionic bond (bond length = 92 pm)? µ = (1.6 × 10−19 C)(9.2 × 10−11 m) = 1.5 × 10−29 C.m × (1D/3.336 × 10−30 C.m) = 4.4 D Dipole Moments (cont.) • Molecular geometry is a critical factor in determining if a molecule has a dipole moment: Which compound is polar? = F = H = C Dipole Moments (cont.) • Molecular geometry is a critical factor in determining if a molecule has a dipole moment: No net dipole moment. Dipoles add as vectors!