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The Basic – Bonding and Molecular Structure. Organic Chemistry and Life.

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Presentation on theme: "The Basic – Bonding and Molecular Structure. Organic Chemistry and Life."— Presentation transcript:

1 The Basic – Bonding and Molecular Structure

2 Organic Chemistry and Life

3

4 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

5 Atomic Orbitals

6 S - orbital p- orbital

7 Electrons Configuration Show

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

9 1.2 The structural Theory of Organic Chemistry Atoms in organic compounds can form a fixed number of bonds using their valence electrons

10 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

11 1.3 Isomers: The Importance of Structural Formulas Constitutional isomers – non identical compounds with same molecular formula Do not necessary share similar properties

12 1.4 Ionic Bonds Occurs in ionic compound Results from transferring electron Created a strong attraction among the closely pack compound

13 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

14 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

15 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+

16 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

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

18 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

19 Examples Write the Lewis structure of CH 3 F, ClO 3 -, F 2

20 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

21 Examples Determine the formal charge for each atom in the following molecules NH 4 + NO 2 - CO 3 2-

22 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 ( )

23 Examples

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

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

26 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

27 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

28 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

29 Wave function

30 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

31 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 sp 3 sp 2 sp

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

33 sp 3 molecular orbitals sp 3 orbitals responsible for creating all “single bonds” of all organic molecules  alkanes

34 Examples

35 sp 2 molecular orbitals All sp 2 molecular orbitals responsible for creating all double bonds in organic molecules  alkenes

36 Examples

37 1.13B – Cis –Trans Isomerism Which of the following alkene can exist as cis-trans isomers? Write their structure

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

39 Examples

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

41 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

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

43 Molecular orbitals

44 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

45 Shape of Molecules

46 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

47 Examples Use VSEPR theory to predict the geometry of each of the following molecules and ions SiF 4 BeF 2

48 1.17 Representation of Structural formulas Structual formula for propyl alcohol

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

50 Dash Formula - Isomerism

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

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

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


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