Lectures Molecular Bonding Theories 1) Lewis structures and octet rule

Slides:



Advertisements
Similar presentations
Covalent Bonding and Molecular Compounds.  Many chemical compounds are composed of molecules.  A molecule is a neutral group of atoms that are held.
Advertisements

8.1 Chemical Bonds, Lewis Symbols, and the Octet Rule
Chapter 8 Covalent Bonding
Quantum Mechanics & Molecular Structure Quantum Mechanics : Quantum mechanics is the foundation of all chemistry and biology. Statistical mechanics rests.
Organic Chemistry Organic Chemistry (10 lectures) Book:
Types of chemical bonds Bond: Force that holds groups of two or more atoms together and makes the atoms function as a unit. Example: H-O-H Bond Energy:
Molecular Orbitals: combine atomic orbitals (AOs) from all the atoms in a molecule into the same number of molecular orbitals. MOs have different shapes,
6/25/2015 One Point Quiz  One quiz per table, list everyone’s name  Agree on an answer  You have two minutes.
Types of chemical bonds Bond: Force that holds groups of two or more atoms together and makes the atoms function as a unit. Example: H-O-H Bond Energy:
Molecular Orbital Theory
Tentative material to be covered for Exam 2 (Wednesday, October 27) Chapter 16Quantum Mechanics and the Hydrogen Atom 16.1Waves and Light 16.2Paradoxes.
Advanced Chemistry Ms. Grobsky. Bonding is the interplay between interactions between atoms Energetically favored Electrons on one atom interacting with.
Daniel L. Reger Scott R. Goode David W. Ball Chapter 9 Chemical Bonds.
Chapter 5: Covalent Bonds and Molecular Structure
Molecular orbital theory Overcoming the shortcomings of the valence bond.
VSEPR Theory
Molecular Geometry and Chemical Bonding Theory
Today’s Quiz 1 1.What is ground-state electron configuration? 2.Define valence electrons and valence shell. 3.Explain the exceptions to the octet rule.
Ch. 3 HW- 3.18, 3.21, 3.32, 3.33, 3.38, 3.39, 3.43, 3.52, 3.53, 3.56, 3.59, 3.61.
Covalent Bonds – Valence Bond (Localized e - ) Model A covalent bonds is the intra-molecular attraction resulting from the sharing of a pair of electrons.
Chapter #10 Chemical Bonding. CHAPTER 12 Forces Between Particles  Noble Gas Configurations  Ionic Bonding  Covalent Bonding  VSEPR Theory and Molecular.
Chemical Bonding I: Basic Concepts
Theories of Bonding and Structure CHAPTER 10 Chemistry: The Molecular Nature of Matter, 6 th edition By Jesperson, Brady, & Hyslop.
Bonding & Molecular Structure Unit 6 Nazanin Ashourian Brittany Haynes.
Chapter 9 Covalent Bonding: Orbitals. Schroedinger An atomic orbital is the energy state of an electron bound to an atomic nucleus Energy state changes.
Covalent Bonding Orbitals Adapted from bobcatchemistry.
The Big Picture1 1.The importance of Coulombs Law: Atomic attraction Relative electronegativity Electron repulsion model for shapes of molecules Choice.
Chapter 8: Periodic Properties of the Elements Chemical Bonds ionic bond covalent bond metallic bond.
Lecture 1 Chemical Bonds: Atomic Orbital Theory and Molecular Orbital Theory Dr. A.K.M. Shafiqul Islam
Chemical Bonding I – Basic Concepts General Chemistry I CHM 111 Dr Erdal OnurhanSlide 1 Lewis Dot Symbols of Some Elements.
Chemical Bonding Unit 4.  Imagine getting onto a crowded elevator. As people squeeze into the confined space, they come in contact with each other. Many.
Chemistry 11 Resource: Chang’s Chemistry Chapter 9.
Chapter 9 Chemical Bonding I: Lewis Theory
BONDING THEORIES SCH4U Grade 12 Chemistry. Lewis Theory of Bonding (1916) Key Points:  The noble gas electron configurations are most stable.  Stable.
Chapter 6 Covalent Compounds. 6.1 Covalent Bonds  Sharing Electrons  Covalent bonds form when atoms share one or more pairs of electrons  nucleus of.
CHAPTER 8 Basic Concepts in Chemical Bonding. Introduction Attractive forces that hold atoms together in compounds are called chemical bonds. The electrons.
Chemical Bonding. Chemical Bonds A bond is a force that holds groups of two or more atoms together and makes them function as a unit. A bond is a force.
Electron Dot Formulas Chemistry 7(C). Lesson Objectives Draw electron dot formulas – Ionic compounds – Covalent compounds Electron Dot Formulas.
Atoms are the smallest units of chemical elements that enter into chemical reactions. ATOM.
Chapter 1 Lecture Introduction and Review Organic Chemistry, 8 th Edition L. G. Wade, Jr.
1 CHEMISTRY 161 Chapter 9 Chemical Bonding I
Carbon’s valence electrons?. Hybrid Orbitals  Mixing of valence shell orbitals to form new similar orbitals for bonding electrons.
Ch. 8 Covalent Bonding Pre AP Chemistry. I. Molecular Compounds  A. Molecules & Molecular Formulas  1. Another way that atoms can combine is by sharing.
2008, Prentice Hall Chemistry: A Molecular Approach, 1 st Ed. Nivaldo Tro Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA.
1 Slater’s Rules for the Determination of Effective Nuclear Charge (Z*) 1) Write out the electronic configuration of the element and group the orbitals.
1 Molecular Geometry and Hybridization of Atomic Orbitals.
Chemical Bonding Lewis Dot Diagrams VSEPR
5.1 Ionic Bonds: Chemical Bonding
Molecular Orbital Theory
Chapter 6 Table of Contents Section 1 Covalent Bonds
Chemical Bonding Review
Covalent Bonding and Molecular Compounds
Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals Chapter 10.
Chapter 1: Orbitals And Bonding
1.1 Atoms, Electrons, and Orbitals
Vocabulary words: Chapter 6: Chemical Bonding Study
Today’s Quiz What is ground-state electron configuration?
Valence Bond Theory.
Chapter 9 Chemical Bonding I: Lewis Theory
Chapter Eight Molecular Structure
Drawing Lewis Structures
Bonding.
Bonding TheorIES SCH4U Grade 12 Chemistry.
Chapter 6: Ionic Bonds and Some Main-Group Chemistry
Quantum Mechanics and Bonding
Chapter 7: Covalent Bonds and Molecular Structure
10.5 (exceptions) Tro's Introductory Chemistry, Chapter 10.
Covalent Bonding and Molecular Compounds
Bonding TheorIES SCH4U Grade 12 Chemistry.
Presentation transcript:

Lectures 5-6 Molecular Bonding Theories 1) Lewis structures and octet rule According to the Lewis theory, a single covalent bond between two atoms forms when they share one electron pair. A double bonds means two shared electron pairs etc. Formation of an ionic bond assumes that one atom loses an electron and another adds it to its shell. The octet rule. For the elements with unavailable d-orbitals (Li … Ne), there is a maximum of 8 electrons in their valence shell. If this number is not reached (BH3, BeH2 etc.), a compound tends to react with electron pair donors (OH2, NH3 etc) to get more shared electrons. Based on the number of electrons in the valence shell, each atom in a chemical species can be assigned a formal charge (the difference in the number of electrons in the element’s valence shell before and after formation of bonds; see + and – signs in the formula above). For the elements with available d orbitals the valence shell can include more than 8 electrons (AlF63-, SiF62-, PCl5, SF4, ClF5, IF7). In some cases atoms can share less (H2+, B2H6) or more (NO) than one electron pair per one bond. Sometimes a configuration with unpaired electrons is more stable than that with paired electrons (O2).

2) Valence Bond (VB) theory One of two major theories describing orbital structure of polyatomic species. Used in modern quantum chemical calculations along with Molecular Orbital (MO) theory. VB theory assumes that covalent bonds are formed when atomic orbitals on two adjacent atoms overlap and electrons on these orbitals are shared. Thus, as a rule, the bonds are two electron, two-center. VB theory gives rise to the concepts of hybridization and resonance. Allows for more exact calculation of bond energies than MO theory. Used extensively in organic chemistry on a qualitative level.

3) Valence bond theory. Origin of bonding in H2+ molecule ion When two nuclei approach each other, electron density decreases at the nuclei which is the region of the strongest nucleus – electron attraction. The overall potential energy Vg of the system increases (becomes more positive) so destabilizing the system. At the same time the kinetic energy Tg of the electron decreases in a much greater extent and stabilizes the system. In quantum mechanics, kinetic energy T of an electron is a function of the square of the gradient of its wavefunction y, T = f(grad(y)2). Good overlap of atomic orbitals leads to lower gradients and thus to lower T what is critical to make a strong bond.

4) Valence bond theory. H2 molecule Consider the wavefunction ycov, g of a system of two hydrogen atoms A and B with electrons 1 and 2 of opposite spin. There are two ways of distributing the electrons, one per atoms, which correspond to y’ and y”. Combining y’ and y” (exchanging electrons by their location) we increase the space each of the electrons 1 and 2 can move within what leads finally to an extra energy gain (exchange energy). When two hydrogen atoms approach each other, their atomic orbitals overlap and the total energy of the system decreases mainly as a result of sharing (exchanging) electrons between two atoms. The H-H bond energy calculated with ycov, g is 72.4 kcal/mol vs 5.7 kcal/mol with either y’ or y” only.

5) Valence bond theory. H2 molecule The experimental value of H-H bond energy is 109.5 kcal/mol. The H-H bond energy of 72.4 kcal/mol calculated with the wavefunction ycov, g can be improved to 87.2 kcal/mol if Z* is used instead of Z to account for the electron shielding. Another way to improve the description of the bonding in the dihydrogen molecule with the resulting H-H bond energy of 92.7 kcal/mol is to consider the ionic contribution (H+H-) with a weight factor l: The covalent and ionic structures are said to be in a resonance with each other:

6) Valence bond vs MO theory In contrast to MO theory, in VB theory electron pairs are localized in a restricted space near two atoms. This leads to smaller error related to accounting for electron correlation and more exact value of bond energies. At the same time in the case of polyatomic molecules with many resonance structures possible application of VB theory becomes quite complicated. MO theory which is simpler is used instead.

7) Resonance A concept suggested to overcome some limitations of Lewis and VB theories and to complement them. Uses multiple canonical (Lewis) structures to describe bonding when a single structure does not allow to account for some features of a species. 1) A favorable resonance structure should have maximum number of covalent bonds.

8) Resonance 2) In resonance structures atoms should not change their connectivity. 3) Structures with separated charges contribute less than those without them. 4) A resonance structure contributes more if sign of a developing charge matches elements’ electronegativity (see above). 5) All resonance structures should have the same number of unpaired electrons.