Organic Chemistry CHE111 Lecture 2 PDF

Faculty of Science Organic Chemistry CHE111 Lecture 2: Bonding and Molecular Structure Prepared by Assoc. Prof. Ahmed Ragab  Chemical Bonds: The Octet Rule ❑ Atoms without the electronic configuration of a noble gas generally...

Faculty of Science Organic Chemistry CHE111 Lecture 2: Bonding and Molecular Structure Prepared by Assoc. Prof. Ahmed Ragab  Chemical Bonds: The Octet Rule ❑ Atoms without the electronic configuration of a noble gas generally react to produce such a configuration because these configurations are known to be highly stable. For all of the noble gases except helium, this means achieving an octet of electrons in the valence shell. Octet Rule: The tendency for an atom to achieve a configuration where its valence shell contains eight electrons. Chemical Bonds: The Octet Rule The Octet Rule  Limitation ❑ Doesn't allow for H, He or Li [stable with 2 e- in their outer shells] - Duet Rule state (hydrogen and helium may have no more than two electrons in their valence shells. ❑ Transition elements - 18 electron rule ❑ BF3 which only has 6 e in its outer shell  Chemical Bonds: The Octet Rule Ionic bonds (Electron transfer) ▪ An atom may lose or gain electrons to acquire a filled valence shell. ▪ An anion: is an atom or group of atoms bearing a negative charge. ▪ A cation: is an atom or group of atoms bearing a positive charge. ▪ Ionic Bond results from the electrostatic attraction of an anion and cation. ▪ It is formed between atoms having large difference in electronegativity (ΔEN>2) [usually a metal and a non-metal].  Chemical Bonds: The Octet Rule ❑ Ionic bonds (Electron transfer) ▪ Electronegativity (EN) is a measure of the ability of an atom to attract electrons. ▪ EN increases as we go across a horizontal period of the periodic table from left to right and it increases as we go up a vertical group.  Chemical Bonds: The Octet Rule Covalent bonds (Electron sharing) ▪ Covalent bonds form by sharing of electrons between two atoms of similar electronegativities (EN < 2) to achieve the configuration of a noble gas. Covalent bonds Non-polar covalent bond Polar Covalent Bond electronegativity is < 0.5 EN is between 0.5 and 1.9  Chemical Bonds: The Octet Rule Non-polar covalent bond ❑ arises between atoms whose difference in electronegativity is <.0.5 ❑ For Example, chlorine is in group 7A, its atoms have 7 valence electrons. Two Cl atoms can share electrons (one electron from each) to form Cl2 molecule. Electron-dot Formula Dash Formula: each dash represents a pair of electrons shared by 2 atoms  Chemical Bonds: The Octet Rule ▪ A carbon atom (group 4A) with 4 valence electrons can share each of these electrons with 4 hydrogen atoms to form a molecule of methane, CH4. ▪ A Nitrogen atom (group 5A) has 5 valence electrons, to verify the octet rule, each N atom share 3 electrons to form N 2.  Chemical Bonds: The Octet Rule ▪ Polar Covalent Bond: ▪ ▪ arises between atoms whose difference in electronegativity is between 0.5 and 1.9. Summary Bond Polarity  Lewis model of bonding Lewis Structure of an Atom ❑ A Lewis Structure shows the symbol of an element surrounded by a number of dots equal to the number of valence electrons. 6 Valence 7 Valence 5 Valence electrons electrons electrons Lewis symbols  Lewis model of bonding ❑ To draw a Lewis structure of a molecule: 1. Lewis structures show the connections between atoms in a molecule or ion using only the valence electrons of the atoms involved. 2. If the structure we are drawing is a negative ion (an anion), we add one electron for each negative charge to the original count of valence electrons. 3. If the structure is a positive ion (a cation), we subtract one electron for each positive charge. 4. In drawing Lewis structures, we try to give each atom the electron configuration of a noble gas. To do so, we draw structures where atoms share electrons to form covalent bonds or transfer electrons to form ions.  Lewis model of bonding ❑ Notes: a. Hydrogen forms one covalent bond by sharing its electron with an electron of another atom so that it can have two valence electrons, the same number as in the noble gas helium. b. Carbon forms four covalent bonds by sharing its four valence electrons with four valence electrons from other atoms, so that it can have eight electrons (the same as the electron configuration of neon, satisfying the octet rule). c. To achieve an octet of valence electrons, elements such as nitrogen, oxygen, and the halogens typically share only some of their valence electrons through covalent bonding, leaving others as unshared electron pairs (Lone pair of electrons). [Nitrogen typically shares three electrons, oxygen two, and the halogens one].  Lewis model of bonding Example 1: Draw the Lewis Structure of CH3F. a. Find the total number of valence electrons of all atoms. C (4) + 3 H (3x1) + F (7) = 14 b. Use pairs of electrons to form bonds between all atoms. Represent bonds with lines. In this example, this require 4 pairs of electrons (8 of 14 valence electrons). c. Add the remaining electrons in pairs to verify the octet rule.  Lewis model of bonding Example 2: Draw the Lewis Structure of CH2O. a. Find the total number of valence electrons of all atoms. C (4) + 2 H (2x1) + O (6) = 12 b. Use 3 pairs of electrons to form bonds between the carbon, 2 hydrogens and the oxygen atom. Why this arrangement of atoms? c. The H atom can have only two electrons associated with it, and so we will not add any more electrons to it. If we place the remaining six electrons around O to give it an octet, we do not achieve an octet on C: Not octet  Lewis model of bonding ❑ We therefore try a double bond between C and O using one of the unshared pairs we placed on O. ❑ The octet rule is satisfied for the C and O atoms, and the H atom has two electrons around it. This is a correct Lewis structure.  Formal charge ▪ Formal charge is the charge of an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. ▪ To calculate the formal charge of atoms in Lewis structures: 1. Write the correct Lewis Structure for the molecule or ion. 2. Assign to each atom its unshared (non-bonding) N B electrons [Lone pair of electrons] and half of its shared (bonding) B electrons. 3. Determine the number of Valence electrons V in the neutral Form Formal charge = V – ( N B + ½ B ) Charge on Molecule = Sum of Charges of Constituting atoms  Formal charge ▪ Formal Charges of Water and Hydronium ion.  Formal charge ❑ If we can draw several Lewis structures for a molecule, the concept of formal charge can help us decide which is the most important, which we shall call the dominant Lewis structure. ❑ Example 1: One Lewis structure for CO2, for instance, has two double bonds. However, we can also satisfy the octet rule by drawing a Lewis structure having one single bond and one triple bond. Calculating formal charges in these structures, we have Note that in both cases the formal charges add up to zero, as they must because CO2 is a neutral molecule. So, which is the dominant structure?  Formal charge How to Identify the Dominant Lewis Structure? 1. The dominant Lewis structure is generally the one in which the atoms bear formal charges closest to zero. 2. A Lewis structure in which any negative charges reside on the more electronegative atoms is generally more dominant than one that has negative charges on the less electronegative atoms. ❑ Thus, the first Lewis structure of CO 2 is the dominant one because the atoms carry no formal charges and so satisfy the first guideline. Dominant Lewis structure  Formal charge ❑ Example 2: Draw the Lewis Structure of ClO3-. a. Find the total number of valence electrons - of all atoms. Cl (7) + 3 O (3x6) + ē (1) = 26 b. Use 3 pairs of electrons to form bonds between the chlorine atom and the 3 oxygen atoms. c. Add the remaining 20 electrons in pairs to verify Octet rule.  Formal charge d. To determine the most stable ClO3– structure, we use the concept of formal charges. 17[Cl] 1s2 2s2 2p6 3s2 3p5 - Element V N B/2 FC Cl 7 2 6/2 +2 O 6 6 2/2 -1 e. Charges on these atoms as shown in the figure make it unstable. The higher the charges, the more unstable it is.  Formal charge f. To combat this, we can transfer the lone pairs in the oxygen atoms to form bonds with chlorine in the middle. This gives chlorine - more than 8 electrons in the middle. This is feasible due to chlorine’s empty 3d orbits, which make it an exception to the octet rule. 17[Cl] 1s2 2s2 2p6 3s2 3p5  Resonance theory ▪ In many molecules and ions (especially those containing π bonds), more than one equivalent Lewis structure can be drawn which represent the same molecule. ▪ For example, we can write three different but equivalent structures for CO3 2- (1–3) O O O C C C O O O O O O Have the same net formal charge 1 2 3  Resonance theory ❑ Structures 1 to 3, although not identical on paper, are equivalent; all of its carbon–oxygen bonds are of equal length. ❑ The bonds are all 1.28֯ A long, intermediate between C-O single bond (1.43֯ A) and C=O double bond (1.20֯ A).  “Curved Arrow Notation” ❑ It shows the direction of electron flow in a reaction mechanism. ❑ The tail of the arrow points from the source of an electron pair (2 ēs) and its head to the atom receiving the pair. ❑ It shows the flow of electrons from a site of higher electron density to a site of lower electron density.  Resonance theory ≡ Contributing Resonance Structures Hybrid structure ▪ An equilibrium is indicated by and resonance by. ▪ Resonance structures exist only on paper, do not represent real molecule. ▪ Lewis structures are called resonance structures or resonance contributors (ē's are moving so fast between those structures), the real molecule or ion is a hybrid of all the hypothetical resonance structures. ▪ Resonance represents the delocalization and stabilization of charge. Molecules with resonance are stabilized!  “Key Principle” ❑ Happy Molecule = Stable = Unreactive ❑ Unhappy Molecule = Unstable = Reactive  Quiz 1. Which element has the largest electronegativity ? a. K b. Br c. F d. C  Quiz 2. The type of chemical bonding between Na and O in CH3COONa is ………… a. Ionic bond b. Polar Covalent Bond c. Pure Covalent Bond d. Nonpolar Covalent Bond Quiz 3. Assign the proper formal charge for the Oxygen atom: a. 1 b. 0 c. -1 d. -2  Quiz 4. The charge of the nitrate ion is …….. a. +1 b. 0 c. -1 d. -2  Quiz 5. The total number of valence electrons of H 3 O + is ….. a. 8 b. 9 c. 10 d. 1