Organic Chemistry Chapter 1 PDF

# ATOMIC STRUCTURE - Atomic # → (Z) = Group # - Electron → the only moving particle/no mass - Proton → +ve, 1 amu - Neutron → x, 1 amu - Isotopes → same element, different mass due to # of neutrons - Eg: C-12 - 6 protons & 6 neutrons - C-14 - 6 protons & 8 neutrons - Atomic No. → # of proton...

# ATOMIC STRUCTURE - Atomic # → (Z) = Group # - Electron → the only moving particle/no mass - Proton → +ve, 1 amu - Neutron → x, 1 amu - Isotopes → same element, different mass due to # of neutrons - Eg: C-12 - 6 protons & 6 neutrons - C-14 - 6 protons & 8 neutrons - Atomic No. → # of protons - Mass No. → total # of protons & neutrons - # of protons = ## of electrons - Must have same atomic # - Istopes: atoms with the same # of protons = same atomic na BUT diff # of neutrons = diff mass # - Eg: Li - Same # of protons - diff # of neutrons = mass # - atomic # = 8 - 3 = 5 neutrons ## Valence Electrons - Upto 7 shells - The outermost shell ### Octet Rule - Atoms want to gain or lose electrons until their valence shell is full - Most elements 8e = full shell - Full valence shell → stability ### Ionic Bonds - Formed by the transfer of 1 or more electrons to create ions - Eg: Na + Cl - Na needs to lose 1 valence electron to become 8e - Cl needs to gain 1 valence electron to become 8e ### Covalent Bonds - Result when atoms share electrons - Eg: N + 3H - N needs to gain 3 valence electrons to become 8e - H needs to gain 1 valence electron to become 2e ## Ionic Bonds - Ions: atoms lose or gain electrons ### Equations - Na → Na+1e or Cl + 1e → Cl-1 - Transfer of Electrons ### Ionic Bond - Opposing charges are attracted to eachother - Electrostatic Forces ### Dot & Cross Diagram - Eg: Dot & Cross Diagram for MgCl2 Atoms - Mg+2 - Cl - Cl ## Electronegativity - Increases across horizontal row of periodic table (left to right) - Increasing as we go up vertical - Eg: How well an atom can attract shared electrons - Nucleus exert force on shared electrons ## Noble Gases - Don't have en values ### Periodic Table - Eg: N → 5 valence electrons - Wants 3 more electrons - F → 7 valence electrons - Wants 1 more electron - Br → 7 valence electrons - Wants 1 more electron - F has more protons, so it pulls the electrons closer which makes it more electronegative - Eg: List from least to most electronegative - Si, Ar, Cs, Si, P, O - Noble Gas # Lewis Structures - Used to represent covalent bonding in molecules and ions - Atoms share valence electrons to get a full octet (except for H) - H & He only need 2 valence electrons for a full shell - One shared = single bond - H2 → H-H - O2 → O=O - C2H4 - H-C=C-H ## Steps - Count all the valence electrons - Determine the central atom - The element there is only one of - Draw single bonds to central atom - Put remaining valence electrons on atoms as lone pairs. - Turn lone pairs into double or triple bonds to give every atom an octet ### Lewis Diagram of H2O - H2 = 2 - O = 6 - 2 + 6 = 8 valence electrons - H - O - H ### Lewis Diagram for SO3 - S = 6 - O3 = 6 x 3 = 18 - 6 + 18 = 24 valence electrons - O - S - O - O - S only has 6 electrons and wants 8 electrons - O will share a lone pair to create a double bond ### Lewis Structure for CH4 - C = 4 - H4 = 1 x 4 = 4 - 4+4 = 8 valence electrons - H - | - H - C - H - | - H ### Lewis Structure for Cl2 - Cl2 = 7 x 2 = 14e - Cl - Cl ### Lewis Structure for C2H6 - C2 = 4 x 2 = 8 - H6 = 1 x 6 = 6 - 8 + 6 = 14 valence electrons - H - C - C -H - H - C - H ### Lewis Structure for SO4-2 - S = 6 x 1 = 6 - O4 = 6 x 4 = 24 - 6 + 24 = 30 + 2 - 32 - When there is a charge, put brackets with the charge outside. - [ O - S - O ] - O - O - Formal Charge = valence electrons - (bonding electrons + lone pairs) - S= 6 - (4+0); 6-4 = +2 (central atom should be zero) - S = 6 - (6+0) = 0 (GOOD) - O = 6-(1+6) = 6-7 = -1 (single O) - O = 6 - (2+4) = 0 (double O) - SO4-2 - Single O = -1 + -1 = -2 - Double O = 0 + 0 = 0 ### Lewis Structure for BrO3- - Br = 7 x 1 = 7 - O3 = 6 x 3 = 18 - 7 + 18 = 25 + 1 = 26e - [ O - Br - O ] - O - Br = 7 - (3+2) = 2 - Br = 7 - (4+2) = 1 → 1 - (5+2) = 0 (GOOD) ### Formal Charge: - F = Z - (1/2)S - U - F = Formal charge - Z = # of valence electrons - S = # of shared electrons - U = # of unshared electrons - Eg: H - C - H - 2 lone pairs - FCs 6-2-6/2 = +1 - FCO 6 - 4 - 4/2 = 0 - FCC = 6 - 6 - 2/2 = -1 - FCO = 6 - 2 - 8/2 = 0 - FCO = 6 - 4 - 4/2 = 0 - Br- - [ O - Br - O ] - O # Isomers - Same chemical formula BUT different structures (order of bonding) - **Constitutional Isomers** - Difference in bonding → leads to different physical properties and chemical properties - Eg. - H3C - CH3 - H3C - | - S - | - CH3 - CH3 - and - H3C - CH3 - | - S - | - CH3 - CH3 ## Structural Formulas 1. **Dash Formula** - Show the connectivity of atoms - Eg. CH4 - H -C-H - H 2. **Condensed Formula** - Eg: H H H - C C C = CH3CHOHCH3 - H H 3. **Bond-Line Formula** - Eg: CH3CHCICH2CH3 - CH3 - CH2 - | - CI - CH - | - CI - CH3 ## 3-D Formulas - X(-) A flat in plane - X(4) A projects out of plane/towards you - X(III) A projects behind the plane/away from you - Eg: Hoc - C-CH2CH3 - OH - H3C - | - C - | - H - CH2CH3 - OH - H3C - | - H - CH2CH3 - OH - H3C - | - CH2CH3 # Resonance Theory - When a molecule/ion can be represented by 2 or more Lewis structures. - Only differ in the positions of electrons. - Not realistic representations with their physical or chemical properties. - The actual molecule will be better represented by an average of these structures. - NOT REAL only exist on paper ## Difference between Resonance & Equilibrium - **Equilibrium** - Diff structures & moving atoms - **Resonance** - Atoms do NOT move - Structures exist only on paper - Allow us to describe molecules for those a single Lewis Structure is inadequate. - Write 2 or more Lewis Structures calling them resonance contributors. - Connect the structures with '&' & the real molecule is a hybrid of them - Eg: - HSC-CH-CH=CH2 - & - H3C-CH=CH-CH2 - H2C-CH2-CH=CH2 - Not a proper resonance structure of 1 & 2 because the hydrogen atom has been moved. - All structures must be proper Lewis Structures - Eg: H - H - | - C - | - O - H - H - Not proper because C has 5 bonds (should be 4) ## Resonance Stabilization - Resonance stabilizes a molecule when the structures are not equivalent. - Energy of resonance hybrid is lower than any contributing structure. - If resonance structures are equivalent → resonance stabilization is LARGE - Eg: NO3 - [ N = O ]- - O - [ O = N - O ]- - O ### Curved Arrows - Movement of 2 electrons - Head - Tail ## Rules for Resonance Structures 1. Only move electrons (-) → lone pairs/double or tripple bonds → so the structure doesn't change - Eg: [O(-):S(=O)] → [O=S(-):O] 2. Never resonate onto sp3 carbon atom - Eg: - H - C - | - [O(-)] - Too many electrons 3. Keep an eye on octets → full shell (8e) cannot take more - *Sulfur is an exception to octet rule - Eg: - [O(-):S(=O)] - O(-) - Carbon does not have a full octet ## Key Arrow Patterns - Start with most negative electron: - Lone pair cannot resonate further: - TT Bond: - Eg: H + - + H - Move electron towards charge: ## Practice Questions - a) NO- - N = 5 - O=18 - 18 + 5 + 1 = 24 - [ O(-):N(=O)] - O(-) - FCN = 5 - 0 - 8/2 = +1 - FCO = 6 - 4 - 4/2 = 0 - FCO -- 6 - 6 - 2/2 = - 1 - O = 0 - N = +1 - minor structure - N is not at full octet - Very unstable - b) CO2 - C = 4 - O=18 - 4 + 18 + 2 = 24 - [O=C=O] - FCO = 4 - 0 - 8/2 = 0 - FCO = 6 - 6 - 2/2 = -1 - FCO = 6 - 6 - 4/2 = -2 - FCO = 6 - 4 - 2/2 = +1 - [O(-)C(=O)] - O(-) - *Cl is an exception to the octet rule ## Electron Configuration - Electrons in 1s orbitals have the lowest energy because they are closest to the positive nucleus. - Electrons in 2s orbitals have the next lowest energy - Electrons of the three 2p orbitals have equal but higher energy than 2s - Degenerate orbitals: orbitals of equal energy (2p) - Eg: - 2s 2s 2s 2s 2s 2s - | - | - | - | - | - | - 2p 2p 2p 2p 2p 2p - | - | - | - | - | - | - 1s 1s 1s 1s 1s - | - | - | - | - | - | - Boron Carbon Nitrogen Oxygen Fluorine Neon - Ground state of a carbon atom (C) - 1s 2s 2px 2py 2pz - | | | | - | | | | ## Structure of Methane and Ethane: sp3 Hybridiziation - Structure of Methane (sp3) - 4sp3 orbitals - Tetrahedral at 109.5 degrees - Eg: H: C: H - H - Structure of Ethane (sp2) - Carbon with 3 sigma bonds + 1 pi bond - Planar at ~120 degrees ## Aufbau Principle - Orbitals are filled so that those of the lowest energy are filled first. ## Pauli Exclusion Principle - A maximum of two electrons may be placed in each orbital but only when the spins of the electrons are paired. ## Hund's Rule - When we come to orbitals of equal energy (degenerate orbitals) such as the three p orbitals, we add one electron to each with their spins unpaired until each of those orbitals contains one electron. (This allows the electrons, which repel each other, to be farther apart. ) Then we begin adding a second electron to each degenerate orbital so that the spins are paired. - Eg: - sp3 Hybridized carbon atom - sp2 Orbital - | - | - sp2 Orbital - | - | - p orbital - | - | - sp2 Orbital - | - | - *An sp3 hybridized carbon atom - *sp2 orbital - *sp2 orbital - *p orbital - *sp2 orbital - * Ground state - * 2p (11 1) - * 2s (1) - * 1s (1) ## Restricted Rotation & Double Bond - There is a large energy barrier to rotation associated with groups joined by a double bond. - Eg: - C=C - ~264 KJ/mol (strength of pi bond) - Rotation of groups joined by C-C single bonds is ~ 13-26 KJ/mol - Eg - C - C ## Cis - Trans Isomerism - Stereochemistry of double bonds - Trans → Opposite sides - Cis → Same side - Eg: - Trans - H - | - C=C - | - H - Cis - H - | - C=C - | - H - 1 sigma bond + 2 pi bonds ## Structure Of Ethyne: Sp Hybridization - H - C≡C - H - Acetylene - Linear structure of carbon with two sigma bonds and two pi bonds at ~180 degrees. - Sp2 hybridized carbon ## Hybridizing - Hybridizing 3 p orbitals with one s orbitals → yields 4 sp3 orbitals = tetrahedral - Hybridizing 2 p orbitals with one s orbital → yields 3 sp2 orbitals = triangular planar - Hybridizing 1 p orbital with one s orbital → yields 2 sp orbitals = linear molecule - Sigma bond (sigma) - A type of single bond - The electron density has circular symmetry viewed along bond axis - Pi bond (pi) - Part of double and triple C-C bonds - Electron densities of two adjacent parallel p orbitals overlap sideways, forming a bonding pi molecular orbital. ## Predicting Molecular Geometry: Valence Shell Electron Pair Repulsion Model (VSEPR) - Molecules in which the central atom is covalently bonded to two or more atoms. 1. Consider all the valence electron pairs of the central atom shared (bonding pairs) and unshared (nonbonding, or lone pairs) 2. Electron pairs repel each other. Electron pairs of a valence shell stay far apart. 3. Repulsion between bond-line pairs is greater than bonding pairs. 4. Describe the shape of the molecule with the positions of nuclei (atoms) and NOT by the positions of electron pairs - Eg - Methane - H - | - H - C - H - | - H - 4 sigma bonds - Ammonia - H - | - H - N - H - | - lone pair - 3 sigma bonds - Water - H - | - H - O - H - | - lone pair - 2 sigma bonds - Boron Trifluoride - F - | - F - B - F - | - F - 3 sigma bonds - Beryllium Hydride - H - Be - H - 2 sigma bonds - The nonbonding pair occupies more space than bonding pairs. - *4 electrons of each double bond act as a single unit and are separated from each other. - TABLE 1. SHAPES OF MOLECULES AND IONS FROM VSEPR THEORY - Electron Pairs at Central Atom - Bonding Nonbonding Total Hybridization Shape of Central Atom or Ion Examples - 2 0 2 sp linear BeH2 - 3 0 3 sp2 Trigonal planar BF3, CH3 - 4 0 4 sp3 Tetrahedral CH4, NH4 - 3 1 4 sp3 trigonal pyramidal NH3, CH3 - 2 2 4 sp3 Angular H2O