The Basic – Bonding and Molecular Structure Chapter 1 The Basic – Bonding and Molecular Structure

Organic Chemistry and Life

1.1 The Development of Organic Chemistry as Science Organic compounds: compounds that could be obtained from living organisms The scientific study of the structure, properties, composition, reactions, and preparation (by synthesis or by other means) of chemical compounds that contain carbon Inorganic compounds: those came from non-living sources Occur as a salts

Atomic Orbitals

Atomic Orbitals S - orbital p- orbital

Electrons Configuration Show

Orbital Diagram and Electron Configuration Electrons configuration: H, He, C, Mg Aufbau Principle: fill lowest orbital first to full capacity, then next

1.2 The structural Theory of Organic Chemistry Atoms in organic compounds can form a fixed number of bonds using their valence electrons

1.2 The structural Theory of Organic Chemistry A carbon atom can use one or more of its valence electrons to form bonds to other carbon atoms

1.3 Isomers: The Importance of Structural Formulas Constitutional isomers – non identical compounds with same molecular formula Do not necessary share similar properties

1.4 Ionic Bonds Occurs in ionic compound Results from transferring electron Created a strong attraction among the closely pack compound

Covalent Bonding Formation of a covalent Bond Two atoms come close together, and electrostatic interactions begin to develop Two nuclei repel each other; electrons repel each other Each nucleus attracts to electrons; electrons attract both nuclei Attractive forces > repulsive forces; then covalent bond is formed

Electronegativity Electronegativity (EN): the ability of an atom in a molecule to attract the shared electron in a bond Metallic elements – low electronegativities Halogens and other elements in upper right-hand corner of periodic table – high electronegativity

Polarity Polar covalent bonds – the bonding electrons are attracted somewhat more strongly by one atom in a bond Electrons are not completely transferred More electronegative atom: δ- . (δ represents the partial negative charge formed) Less electronegative atom: δ+

Lewis Structures represents how an atom’s valence electrons are distributed in a molecule Show the bonding involves (the maximum bonds can be made) Try to achieve the noble gas configuration

Rules Duet Rule: sharing of 2 electrons E.g H2 H : H Octet Rule: sharing of 8 electrons Carbon, oxygen, nitrogen and fluorine always obey this rule in a stable molecule E.g F2, O2 Bonding pair: two of which are shared with other atoms Lone pair or nonbonding pair: those that are not used for bonding

1 Lewis Structures of Molecules with Multiple Bonds Use 6N + 2 Rule N = number of atoms other than Hydrogen If Total valence – (6N + 2) = 2  1 double bond Total valance e- - (6N + 2) = 4  two double bonds or 1 triple bond

Examples Write the Lewis structure of CH3F, ClO3-, F2

1.7 Formal Charges Difference between the number of outer-shell electrons “owned” by a neutral free atom and the same atom in a compound

Examples Determine the formal charge for each atom in the following molecules NH4+ NO2- CO32-

Resonance Whenever a molecule or ion can be represented by two or more Lewis structures that differ only in the position of the electrons None of these resonance structures will be a correct representation for the molecule or ion The actual molecule or ion will be better represented by a hybrid or hypothetical structures Represented by a double headed arrows ( )

Examples

Resonance - stabilization The more covalent bonds a structure has, the more stable it is

Resonance-stabilization Structure in which all the atoms have a complete valence shell of electrons are especially

Resonance stabilization Charge separation decrease stabilization Resonance contributors with negative charge on highly electronegative atoms are stable ones with negative charge on less or nonelectronegative atoms

1.9 Quantum Mechanisms and Atomic Structure Schröndinger’s quantum mechanical model of atomic structure is frame in the form of a wave equation; describe the motion of ordinary waves in fluids. i. Wave functions or orbitals (Greek, psi , the mathematical tool that quantum mechanic uses to describe any physical system ii. 2 gives the probability of finding an electron within a given region in space iii. Contains information about an electron’s position in 3-D space defines a volume of space around the nucleus where there is a high probability of finding an electron say nothing about the electron’s path or movement

11.2 Electromagnetic Radiation Radiation energy – has wavelike properties Frequency (υ, Greek nu) – the number of peaks (maxima) that pass by a fixed point per unit time (s-1 or Hz) Wavelength (λ, Greek lambda) – the length from one wave maximum to the next Amplitude – the height measured from the middle point between peak and trough (maximum and minimum) Intensity of radiant energy is proportional to amplitude

Wave function

1.10 Atomic orbital Heisenberg Uncertainty Principle – both the position (Δx) and the momentum (Δmv) of an electron cannot be known beyond a certain level of precision 1. (Δx) (Δmv) > h 4π 2. Cannot know both the position and the momentum of an electron with a high degree of certainty

Molecular Orbitals Two types of atomic of atomic orbitals are combined as they come close to each other Hybridization: blending combination of atomic orbitals to form new orbital Carbon has three possible molecular orbitals sp3 sp2 sp

Orbitals repsonsible for creating the covalent bonds 2 special names for covalent bonds of organic molecules Sigma (σ) bond Pi (π) bond Created when “head on” overlap occurs of orbitals Created when “side on” overlap occurs of orbitals

sp3 molecular orbitals sp3 orbitals responsible for creating all “single bonds” of all organic molecules  alkanes

Examples

sp2 molecular orbitals All sp2 molecular orbitals responsible for creating all double bonds in organic molecules  alkenes

Examples

1.13B – Cis –Trans Isomerism Which of the following alkene can exist as cis-trans isomers? Write their structure

sp molecular orbitals All sp orbitals responsible for creating all triple bonds of organic molecules  alkynes

Examples

Examples Draw a bonding picture for the following molecule, showing all π, σ – bonds using σ-framework and π-framework

Molecular Orbitals Two types: Bonding molecular orbitals Contains both electrons in the lowest energry state or ground state Formed by intereaction of orbitals with same phase signs Increases the propability

Molecular orbitals Antimolecular orbitals Contains no electrons in the ground state Formed by intereaction of orbitals with opposite phase signs Result with nodes

Molecular orbitals

Shape of Molecules VSEPR Theory Valence shell electron pair repulsion Bond angles and geometry Steric number = # bond to - # lone pairs central atom to central atom Rules: 1- Carbon will always be the central atom 2 – Double bond; triple bonds will count as 1 bond

Shape of Molecules

Molecular shapes VSEPR method can be used to predict the shapes of molecules containing multiple bonds Assume that all electrons of a multiple bond act as one unit

Examples Use VSEPR theory to predict the geometry of each of the following molecules and ions SiF4 BeF2

1.17 Representation of Structural formulas Structual formula for propyl alcohol

Dash Structure Atoms are joined by single bonds can rotate relatively freely with respect to one another

Dash Formula - Isomerism

Condensed Structural Formulas All hydrogen atoms are written immediately after the carbon that they’re attached

Bond-Line Formulas Hydrogen and carbon atoms will not appear in the formula Each end of the line represents carbon atom

Examples For each of the following, write a bond line formula