university of isfahan department of chemistry fundamentals of organic chemistry for biology students...
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University of IsfahanUniversity of IsfahanDepartment of Department of
ChemistryChemistry
FUNDAMENTALS OF FUNDAMENTALS OF ORGANIC CHEMISTRYORGANIC CHEMISTRY
FOR BIOLOGY STUDENTSFOR BIOLOGY STUDENTS
Dr. Gholam Ali Koohmareh
اول 90-91نيمسال
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Text Book:
Fundamentals of Organic Chemistry
By: John. Mc. Murry
مباني شيمي آلي
نوشته جان مك موري
ترجمه دكتر عيسي ياوري
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مورد در چند نصايحيآلي شيمي مطالعه
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مفقوده ماده
در دانش در ساختار دانش ساختارآلي آلي شيمي شيمي
شبيه ساختارهااست متعادلي هرم
روي كهاست ايستاده نوك
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ميشود خراب ساختار !!!كل
باشد شده گم چيزي .....اگر
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كنيد مطالعه بايد مواد ساختار درك .براي
بار چندين نميشويد مسلط كه را هرمطلبيكنيد .تكرار
بحثهاي وارد و ميكنيم شروع ساده مباحث از ماميشويم تر .پيچيده
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از قبل و زودتر چه هر نميشويد متوجه را مطلبي اگربگيريد كمك مطالب شدن ....انباشته
شيمي گروه در من اطاق در اداري ساعات طول در يا كالس از بعد ميتوانيدكنيد مطرح را خود درسي مشكل ايميل طريق از يا و
Email: [email protected]
كنيد حل تمرين داشتيد وقت كه زماني .هر
است آن مستلزم اين و ميكند بيشتر را شما درك تمرينات انجامگيريد بكار ميگيريد ياد چه هر .كه
مفيدند بسيار فصل آخر .تمرينات
است اجباري ازمطالعه موثرتر بسيار فعال !يادگيري
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ميگيريم %90 بكار و ميگوييم كه را چيزي
ميكنيم %70 صحبت و ميگوييم كه را چيزي
ميشنويم %50 و ميبينيم كه را چيزيميكنيم حفظ ما
ميبينيم %30 كه را چيزي
ميشنويم %20 كه را چيزي
ميخوانيم %10 كه را چيزي
روانشناسي ..… مطالعه يك
واقعي ولي واقعي عجيب ولي عجيب
فعال
اجباري
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. پيوند و ساختار اول فصل
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Organic Chemistry “Organic” – until mid 1800’s referred to
compounds from living sources (mineral sources were “inorganic”)
Wöhler in 1828 showed that urea, an organic compound, could be made from a non-living source:
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Organic Chemistry
Today, organic compounds are those based on carbon structures and organic chemistry studies their structures and reactions Includes biological molecules, drugs,
solvents, dyes, food additives, pesticides, and others.
Does not include metal carbonate salts and other simple ionic compounds (inorganic)
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Organic Chemistry
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1.1 Atomic Structure Structure of an atom
Positively charged nucleus (very dense, protons and neutrons) and small (10-15 m)
Negatively charged electrons are in a cloud (10-10 m) around nucleus
Diameter is about 2 10-10 m (200 picometers (pm)) [the unit angstrom (Å) is 10-10 m = 100 pm]
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1.1 Atomic Structure
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Atomic Number & Mass Number The atomic number (Z) is the
number of protons in the atom's nucleus
The mass number (A) is the number of protons plus neutrons
All the atoms of a given element have the same atomic number
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Atomic Number and Atomic Mass All the atoms of a given element have
the same atomic number Isotopes are atoms of the same
element that have different numbers of neutrons and therefore different mass numbers
The atomic mass (atomic weight) of an element is the weighted average mass in atomic mass units (amu) of an element’s naturally occurring isotopes
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1.2 Atomic Structure: Orbitals Quantum mechanics: describes electron energies
and locations by a wave equation Wave function is a solution of the wave equation Each Wave Function describes an orbital,
A plot of 2 describes the probabililty of finding the electron
The electron cloud has no specific boundary so we show the most probable volume.
The boundary is called a “probablility surface”
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Shapes of Atomic Orbitals for Electrons For the existing elements, there are four different
kinds of orbitals for electrons based on those derived for a hydrogen atom.
Denoted s, p, d, and f s and p orbitals most important in organic
chemistry (carbon atoms have no d or f orbitals)
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Shapes of Atomic Orbitals for Electrons s orbitals: spherical, nucleus at center p orbitals: dumbbell-shaped, nucleus at
middle
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Orbitals and Shells
Orbitals are grouped in shells of increasing size and energy
As shells increase in size, they contain larger numbers and kinds of orbitals
Each orbital can be occupied by two electrons (Pauli Exclusion Principle)
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Orbitals and Shells First shell contains one s orbital, denoted 1s, holds
only two electrons Second shell contains one s orbital (2s) and three p
orbitals (2p): eight electrons Third shell contains an s orbital (3s), three p orbitals
(3p), and five d orbitals (3d): 18 electrons
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Orbitals and Shells
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p-Orbitals In each shell there
are three perpendicular p orbitals, px, py, and pz, of equal energy
Lobes of a p orbital are separated by region of zero electron density, a node
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1.3 Atomic Structure: Electron Configurations
Ground-state electron configuration of an atom lists orbitals occupied by its electrons.
1. Lowest-energy orbitals fill first: 1s 2s 2p 3s 3p 4s 3d (Aufbau (“build-up”) principle)
2. Electron spin can have only two orientations, up and down . Only two electrons can occupy an orbital, and they must be of opposite spin (Pauli exclusion principle) to have unique wave equations
3. If two or more empty orbitals of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron (Hund's rule).
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Electron Configurations
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Electron Configurations
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1.4 Development of Chemical Bonding Theory Kekulé and Couper independently observed that
carbon always has four bonds van't Hoff and Le Bel proposed that the four bonds
of carbon have specific spatial directions Atoms surround carbon as corners of a tetrahedron
Note that a wedge indicates a bond is coming forward
Note that a dashed line indicates a bond is behind the page
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Tetrahedral Carbon:
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1.5 The Nature of the Chemical Bond
Atoms form bonds because the compound that results is more stable than the separate atoms
Ionic bonds in salts form as a result of electron transfers
Organic compounds have covalent bonds from sharing electrons (G. N. Lewis, 1916)
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1.5 The Nature of the Chemical Bond
Lewis structures shown valence electrons of an atom as dots Hydrogen has one dot, representing
its 1s electron Carbon has four dots (2s2 2p2)
Stable molecule results when the outermost (valence) shell is completed: octet (eight dots) for main-group atoms (two for hydrogen)
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Number of Covalent Bonds to an Atom
Atoms with one, two, or three valence electrons form one, two, or three bonds
Atoms with four or more valence electrons form as many bonds as they need electrons to fill the s and p levels of their valence shells to reach a stable octet
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Number of Covalent Bonds to an Atom
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Valences of Carbon Carbon has four valence electrons (2s2
2p2), forming four bonds (CH4)
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Valences of Oxygen Oxygen has six valence electrons (2s2 2p4) but forms
two bonds (H2O)
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Valences of Nitrogen Nitrogen has five valence electrons (2s2
2p3) but forms only three bonds (NH3)
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Non-bonding electrons Valence electrons not used in bonding are called
nonbonding electrons, or lone-pair electrons Nitrogen atom in ammonia (NH3)Shares six
valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair
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Some simple molecules:
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1.6 Valence Bond Theory Covalent bond forms when two
atoms approach each other closely so that a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom
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1.6 Valence Bond Theory
Electrons are paired in the overlapping orbitals and are attracted to nuclei of both atoms
H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals
H-H bond is cylindrically symmetrical, sigma () bond
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Bond Energy Reaction 2 H· H2 releases 436 kJ/mol Product has 436 kJ/mol less energy than two
atoms: H–H has bond strength of 436 kJ/mol. (1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ)
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Bond Length
Distance between nuclei that leads to maximum stability
If too close, they repel because both are positively charged
If too far apart, bonding is weak
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What is a Bond? An Energy Minimum
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1.7 Hybridization: sp3 Orbitals and the Structure of Methane Carbon has 4 valence electrons (2s2 2p2) In CH4, all C–H bonds are identical (tetrahedral) sp3 hybrid orbitals: s orbital and three p orbitals
combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3), Pauling (1931)
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sp3 Hybridization
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Tetrahedral Structure of Methane sp3 orbitals on C overlap with 1s orbitals on 4 H
atom to form four identical C-H bonds Each C–H bond has a strength of 438 kJ/mol and
length of 110 pm Bond angle: each H–C–H is 109.5°, the
tetrahedral angle.
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Tetrahedral Structure of Methane
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1.8 Hybridization: sp3 Orbitals and the Structure of Ethane
Two C’s bond to each other by overlap of an sp3 orbital from each
Three sp3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds
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1.8 Hybridization: sp3 Orbitals and the Structure of Ethane
C–H bond strength in ethane 420 kJ/mol
C–C bond is 154 pm long and strength is 376 kJ/mol
All bond angles of ethane are tetrahedral
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1.9 Hybridization: sp2 Orbitals and the Structure of Ethylene
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1.9 Hybridization: sp2 Orbitals and the Structure of Ethylene
sp2 hybrid orbitals: 2s orbital combines with two 2p orbitals, giving 3 orbitals (spp = sp2)
sp2 orbitals are in a plane with120° angles
Remaining p orbital is perpendicular to the plane
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Bonds From sp2 Hybrid Orbitals Two sp2-hybridized orbitals overlap to form a bond p orbitals overlap side-to-side to formation a pi () bond sp2–sp2 bond and 2p–2p bond result in sharing four
electrons and formation of C-C double bond Electrons in the bond are centered between nuclei Electrons in the bond occupy regions are on either side
of a line between nuclei
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Bonds From sp2 Hybrid Orbitals
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Structure of Ethylene H atoms form bonds with four sp2 orbitals H–C–H and H–C–C bond angles of about 120° C–C double bond in ethylene shorter and
stronger than single bond in ethane Ethylene C=C bond length 133 pm (C–C 154
pm)
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1.10 Hybridization: sp
Orbitals and the Structure of Acetylene
C-C a triple bond sharing six electrons
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1.10 Hybridization: sp
Orbitals and the Structure of Acetylene Carbon 2s orbital hybridizes with a single p
orbital giving two sp hybrids (two p orbitals remain unchanged)
sp orbitals are linear, 180° apart on x-axis Two p orbitals are perpendicular on the y-axis
and the z-axis
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Orbitals of Acetylene Two sp hybrid orbitals from each C form sp–sp bond pz orbitals from each C form a pz–pz bond by sideways
overlap and py orbitals overlap similarly
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Bonding in Acetylene Sharing of six electrons forms C C Two sp orbitals form bonds with hydrogens
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Comparison of C-C & C-H bonds:
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Comparison of C-H bonds:
Molecule BondEnergy (kcal)
Length (pm)
Ethane Csp3-H 100 110
Ethylene Csp2-H 106 108
Acetylene Csp-H 132 106
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Comparison of C-C bonds:
Molecule BondEnergy (kcal)
Length (pm)
Ethane Csp3-Csp3 90 154
Ethylene Csp2-Csp2 146 133
Acetylene Csp-Csp 200 120
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1.11 Hybridization of Nitrogen and Oxygen
Elements other than C can have hybridized orbitals
H–N–H bond angle in ammonia (NH3) 107.3°
N’s orbitals (sppp) hybridize to form four sp3 orbitals
One sp3 orbital is occupied by two nonbonding electrons, and three sp3 orbitals have one electron each, forming bonds to H
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Hybridization of Oxygen in Water The oxygen atom is sp3-hybridized Oxygen has six valence-shell electrons but
forms only two covalent bonds, leaving two lone pairs
The H–O–H bond angle is 104.5°
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1.12 Acids and Bases: The Brønsted–Lowry Definition
Brønsted–Lowry theory defines acids and bases by their role in reactions that transfer protons (H+) between donors and acceptors.
“proton” is a synonym for H+ - loss of an electron from H leaving the bare nucleus—a proton. Protons are always covalently bonded to another atom.
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Brønsted Acids and Bases “Brønsted-Lowry” is usually shortened
to “Brønsted” A Brønsted acid is a substance that
donates a hydrogen ion, or “proton” (H+): a proton donor
A Brønsted base is a substance that accepts the H+: a proton acceptor
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The Reaction of HCl with H2O
When HCl gas dissolves in water, a Brønsted acid–base reaction occurs
HCl donates a proton to water molecule, yielding hydronium ion (H3O+) and Cl
The reverse is also a Brønsted acid–base reaction of the conjugate acid and conjugate base
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The Reaction of HCl with H2O
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Quantitative Measures of Acid Strength
The equilibrium constant (Ke) for the reaction of an acid (HA) with water to form hydronium ion and the conjugate base (A-) is a measure related to the strength of the acid
Stronger acids have larger Ke
Note that brackets [ ] indicate concentration, moles per liter, M.
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Ka – the Acidity Constant The concentration of water as a solvent does not change
significantly when it is protonated in dilute solution. The acidity constant, Ka for HA equals Ke times 55.6 M
(leaving [water] out of the expression) Ka ranges from 1015 for the strongest acids to very small
values (10-60) for the weakest
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Ka – the Acidity Constant
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1.13 Acid and Base Strength
The ability of a Brønsted acid to donate a proton to is sometimes referred to as the strength of the acid.
The strength of the acid can only be measured with respect to the Brønsted base that receives the proton
Water is used as a common base for the purpose of creating a scale of Brønsted acid strength
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pKa – the Acid Strength Scale
pKa = -log Ka (in the same way that
pH = -log [H+] The free energy in an equilibrium is
related to –log of Keq (G = -RT log Keq) A larger value of pKa indicates a
stronger acid and is proportional to the energy difference between products and reactants
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pKa – the Acid Strength Scale
The pKa of water is 15.74
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pKa – the Acid Strength Scale
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1.14 Predicting Acid–Base Reactions from pKa Values
pKa values are related as logarithms to equilibrium constants
The difference in two pKa values is the log of the ratio of equilibrium constants, and can be used to calculate the extent of transfer
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Predicting Acid–Base Reactions from pKa Values
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Predicting Acid–Base Reactions from pKa Values
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Prob’s 2.14 & 2.15: Will these reactions take place as written?
pKa = 19
pKa = 36
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1.15 Organic Acids and Organic Bases
The reaction patterns of organic compounds often are acid-base combinations
The transfer of a proton from a strong Brønsted acid to a Brønsted base, for example, is a very fast process and will always occur along with other reactions
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Some organic acids:
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Organic Acids Those that lose a proton from
O–H, such as methanol and acetic acid
Those that lose a proton from C–H, usually from a carbon atom next to a C=O double bond (O=C–C–H)
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Organic Acids
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Conjugate bases:
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Organic Bases Have an atom with a lone pair
of electrons that can bond to H+
Nitrogen-containing compounds derived from ammonia are the most common organic bases
Oxygen-containing compounds can react as bases when with a strong acid or as acids with strong bases
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Organic Bases
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1.16 Acids and Bases: The Lewis Definition Lewis acids are electron pair acceptors; Lewis
bases are electron pair donors The Lewis definition leads to a general description of
many reaction patterns but there is no quantitatve scale of strengths as in the Brønsted definition of pKa
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Lewis Acids and the Curved Arrow Formalism
The Lewis definition of acidity includes metal cations, such as Mg2+
They accept a pair of electrons when they form a bond to a base
Group 3A elements, such as BF3 and AlCl3, are Lewis acids because they have unfilled valence orbitals and can accept electron pairs from Lewis bases
Transition-metal compounds, such as TiCl4, FeCl3, ZnCl2, and SnCl4, are Lewis acids
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Lewis Acids and the Curved Arrow Formalism
Organic compounds that undergo addition reactions with Lewis bases (discussed later) are called electrophiles and therefore Lewis Acids
The combination of a Lewis acid and a Lewis base can shown with a curved arrow from base to acid
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Illustration of Curved Arrows in Following Lewis Acid-Base
Reactions
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BF3 as a Lewis Acid:
BF3 O(CH3)2 F3B-O(CH3)2
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Lewis Bases Lewis bases can accept protons as
well as other Lewis acids, therefore the definition encompasses that for Brønsted bases
Most oxygen- and nitrogen-containing organic compounds are Lewis bases because they have lone pairs of electrons
Some compounds can act as either acids or bases, depending on the reaction
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Lewis Bases
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Prob. : Identify acids & bases
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Summary Organic chemistry – chemistry of carbon compounds Atom: positively charged nucleus surrounded by
negatively charged electrons Electronic structure of an atom described by wave
equation Electrons occupy orbitals around the nucleus. Different orbitals have different energy levels and different
shapes s orbitals are spherical, p orbitals are dumbbell-shaped
Covalent bonds - electron pair is shared between atoms
Valence bond theory - electron sharing occurs by overlap of two atomic orbitals
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Summary Sigma () bonds - Circular cross-section and are formed
by head-on interaction Pi () bonds – “dumbbell” shape from sideways interaction
of p orbitals Carbon uses hybrid orbitals to form bonds in organic
molecules. In single bonds with tetrahedral geometry, carbon has four sp3
hybrid orbitals In double bonds with planar geometry, carbon uses three
equivalent sp2 hybrid orbitals and one unhybridized p orbital Carbon uses two equivalent sp hybrid orbitals to form a
triple bond with linear geometry, with two unhybridized p orbitals
Atoms such as nitrogen and oxygen hybridize to form strong, oriented bonds
The nitrogen atom in ammonia and the oxygen atom in water are sp3-hybridized
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Summary
A Brønsted(–Lowry) acid donatea a proton
A Brønsted(–Lowry) base accepts a proton
The strength Brønsted acid is related to the -1 times the logarithm of the acidity constant, pKa. Weaker acids have higher pKa’s
A Lewis acid has an empty orbital that can accept an electron pair
A Lewis base can donate an unshared electron pair