pharmamatrixworkshop2010 - university of alberta
TRANSCRIPT
PharmaMatrix Workshop 2010 University of Alberta
July 12, 2010
Jonathan Y. Mane and Melissa Gajewski
Ma3er
Pure substance
Element Composed of atoms
Cannot be chemically decomposed
Compound Combina?on of >1 elements represented by chemical formula
Molecular/Covalent Electrons are shared between atoms (e.g. H2O)
Ionic Electrons not shared (e.g. NaCl)
Mixture
Homogeneous Uniform
throughout
Heterogeneous
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English name Symbol Non-‐English name English name Symbol Non-‐English name
An?mony Sb s?bium Potassium K kalium
Copper Cu cuprum Silver Ag argentum
Gold Au aurum Sodium Na natrium
Iron Fe ferrum Tin Sn stannum
Lead Pb plumbum Tungsten W wolfram
Mercury Hg hydragyrum
Element (Naming and interna?onal symbols)
First 1-‐2 dis?nguishing le3ers in the name
Not all names derived from English language
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The Periodic Table of Elements Alkali metals Alkaline earth metals
Halogens
Noble gases
d-‐transi?on metals
f-‐transi?on metals
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Elements and their common ions
caAons (lose e-‐) anions (gain e-‐)
hydride H-‐
fluoride F-‐
chloride Cl-‐
bromide Br-‐
oxide O2-‐
sulfide S2-‐
nitride N3-‐
Elemental anions
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Compound (= combina?on of >1 element represented by chemical formula)
Exercise • Write the formula of the compound formed between
barium and bromine? • Write the formula of the compound formed between
the ammonium ion and the carbonate ion? • Break the formula Mg3(PO4)2 into its ions?
Chemical/molecular formula = specifies the number of atoms of each element in one unit/molecule of the substance
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ammonuim NH4+
nitrate NO3-‐
sulfate SO42-‐
carbonate CO32-‐
phosphate PO43-‐
cynanide CN-‐
hydroxide OH-‐
peroxide O22-‐
Common polyatomic ions
Molecular/covalent compound Electrons are shared between atoms (e.g. H2O)
Ionic compound Electrons are not shared but completely transferred (ca?on + anion : Na+ + Cl-‐ NaCl)
Covalent bonding – sharing of electrons between atoms
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Core electrons = electrons other than the valence
= electrons in the outermost main energy level; involved in chemical reac?on
Covalent bonding – sharing of electrons between atoms
Lewis electron dot structure (G.N. Lewis, 1875-‐1946) – a symbolism that shows the number of valence electrons of atoms
Elements Metal ions (for ionic bonding)
Na+ Mg2+ Al3+
Anions
Bond and # of pairs of shared electrons Single bond = one ( 2 dots or single line)
Double bond = two ( 4 dots or 2 lines) Triple bond = three ( 6 dots or 3 lines)
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Examples
Covalent bonding – Lewis structure Rules for wriAng Lewis structures: 1. Count the number of outermost electrons on each atom of the formula before bonding and obtain the
total number of electrons. 2. Arrange the atoms by designa?ng a central atom and the rest surrounding it. (Hydrogen cannot be a
central atom; Oxygen never bond to another oxygen except in O2(g), ozone (O3), and peroxide (O22-‐).
3. Distribute the electrons by placing 8 electrons (dots) around all atoms, one pair on each side (making only single bonds for now), except for any hydrogen (max. 2 dots).
4. Step 1 #e-‐ = Step 3 #e-‐, celebrate your structure is CORRECT!. 5. Step 1 #e-‐ ≠ Step 3 #e-‐, create double or triple bonds between the central atom and other atoms. 6. If the central atom is phosphorus, or atoms of elements to the right and/or below phosphorus in PT, the
central atom may accommodate more than 8 electrons, if necessary, to make the counts in Steps 1 and 4 match.
O
H H C
Cl
Cl
ClCl
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Examples
water carbon tetrachloride nitrate
resonance structures
Covalent bonding – formal charge
The following must be saAsfied to get zero formal charge for these atoms
C = 4 bonds | N = 3 bonds | O = 2 bonds | Halogens = 1 bond
FC = group number – # e-‐ in lone pairs – 0.5( # e-‐ in bonding pairs)
O
H HC
Cl
Cl
ClCl
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How to calculate formal charge (FC):
Valence shell electron-‐pair repulsion theory A theory that helps in predic?ng geometric shapes of molecules based on the concept that the electron pairs, be it bonding or non-‐bonding pairs, will repel each other.
Prerequisite: Knowledge of drawing Lewis electron-‐dot structure
A = central atom X = number of sigma bonds between the
central atom and the other atom (mul?ple bond counts as one X)
E = number of lone electron pairs surrounding the central atom
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Exercise – Geometric shapes of molecules Draw the Lewis electron dot diagram and predict the shape of the following molecules using VSEPR theory:
Carbon dioxide CO2
Nitrite NO2-‐
Carbonate CO32-‐
Ammonia NH3
Phosphate PO43-‐
Sulfur tetrafluoride SF4
Xenon tetrafluoride XeF4
Iodine heptafluoride IF7
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Structure and QM model of the atom Essence of the QM model: • Energies are quan?zed • Electrons exist in principal energy levels, n (aka shells), in energy sublevels (aka subshells) within these principal levels, and in regions of space called orbitals within the sublevels • Each orbital can hold a maximum of 2 e-‐
• Electrons have a par?cular spin direc?on
Shell, n = 1, 2, 3, … Max number of e-‐ possible in level n = 2n2
n Max # of e-‐ # orbitals Orbital types
1 2 1 1s
2 8 4 1s, 3p
3 18 9 1s, 3p, 5d
4 32 16 1s, 3p, 5d, 7f
Number of orbitals in level n = n2 Subshell = s, p, d, f, g, etc.
Subshell Max # of e-‐
s 2
p 6
d 10
f 14
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Rela?onship between atomic orbitals and the periodic table
• Period = main energy level • Subshells s, p, d, f = block or group of elements
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Filling up orbitals and electronic configura?on Hund’s rule: All orbitals within a given sublevel must get one electron before any get two.
Electronic configuraFon
1s 2s 2px 2py 2pz
H 1s1
He 1s2
Li 1s22s1
Be 1s22s2
B 1s22s22p1
C 1s22s22p2
N 1s22s22p3
O 1s22s22p4
F 1s22s22p5
Ne 1s22s22p6
Na [Ne] 3s1
Paired e-‐
Unpaired e-‐
Core e-‐
Valence e-‐
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Molecular Orbitals
• Regions of space in which shared electrons reside in covalent bonding • Overlap of atomic orbitals (end-‐to-‐end and side-‐to-‐side) • The number of molecular orbitals is EQUAL to the number of atomic orbitals
End-‐to-‐end overlap Side-‐to-‐side
σ-‐bond • always present in any covalent bond • one of the bonds in double or triple bond (formed from two p orbitals)
π-‐bond • overlap between two p orbitals; • always accompanied by a σ-‐bond • only occur in double or triple bonds • a double bond = 1 π-‐bond • a triple bond = 2 π-‐bond
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Orbital hybridizaFon (the concept of mixing atomic orbitals to form new hybrid orbitals suitable for the qualita?ve descrip?on of atomic bonding proper?es)
sp3 hybridizaFon -‐ combinaAon of four orbitals (one s and three p-‐orbitals) forming a tetrahedral geometry
sp2 hybridizaFon -‐ one s and two p-‐orbitals from each C are used to construct a σ-‐bond network forming a trigonal geometry -‐ the remaining one p-‐orbital from each C are orthogonal to the σ-‐bond network
used to construct the π-‐bond network
Con?nuous π-‐cloud above and below the σ-‐bond plane
sp hybridizaFon -‐ one s and one p-‐orbital from each C are used to construct a σ-‐bond network forming linear geometry
Con?nuous π-‐cloud sur-‐rounding the σ-‐bond plane
H
H
H
H
H H
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Examples using acyclic hydrocarbons
# of C Alkyl name Alkane
1 methyl methane
2 ethyl ethane
3 propyl propane
4 butyl butane
5 pentyl pentane
6 hexyl hexane
7 heptyl heptane
8 octyl octane
9 nonyl nonane
CH2H3C
CH3
Explicit C and H
Explicit C and H terminals H3C
CH3
Implicit C and H terminals
Alkenes (replace alkane suffix –ane with –ene)
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CH3CH2CH3
cis-‐2-‐butene trans-‐2-‐butene
propane
Examples using cyclic hydrocarbons
# of C Cycloalkane
3 cyclopropane
4 cyclobutane
5 cyclopentane
6 cyclohexane
… …
n n-‐cycloalkane
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Conforma?onal Isomerism
ConformaAonal isomers = molecules with the same structural formula and connec?vity but differs in their 3D structures due to rota?ons about one or more σ bonds. • gauche, an?, and eclipse conformers of butane • boat and chair conformers of cyclohexane
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Configura?onal Isomerism
cis-‐2-‐butene trans-‐2-‐butene
Cis-‐trans isomers = a form of stereoisomerism describing the orienta?on of groups (func?onal groups) within a molecule.
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E = entgegen means opposite(=trans) Z = means together zusammen (=cis)
3D Structures on paper and chirality
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C
Cl
BrH
F
C
Cl
BrH
F
Lines and wedges EnanAomers or opAcal isomers = mirror images of a molecule that cannot be superposed onto each other
R/S Configura?on • An important nomenclature system for deno?ng enan?omers • Labels a chiral center according to a system by which its subs?tuents are each assigned a priority based on atomic number • R = priority decreases in clockwise direc?on • S = priority decreases in counterclockwise (S)-‐alanine (R)-‐alanine
Chiral center = an atom that is bonded to four different atoms or groups of atoms in such a manner that it has a non-‐superimposable mirror image
Common Func?onal Groups (R, R’, R” = hydrocarbon radicals or H)
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R X
Alkyl halide
CH3CH2Br
bromoethane
R OH
CH3CH2OH
Alcohol
ethanol
R O R'
Ether
O CH3H3C
dimethyl ether Phenol
OH
Phenol
R C
O
H
Aldehyde
butanal
CH3CH2CH2 C H
O
C
O
R
R'
Ketone
H3C C CH3
O
propanone
R C
O
OH
Carboxylic acid
H3C C OH
O
propanoic acid
Ester
R C
O
O R'
H3C C O
O
CH3
methyl acetate
Common Func?onal Groups (R, R’, R” = hydrocarbon radicals or H)
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1o Amine
R NH2
CH3NH2
methylamine
2o Amine
R N H
R'
3o Amine
R N R"
R'
Nitrile
acetonitrile
R C N
CH3C N
Nitro
R N
O
O
nitromethane
CH3NO2
1o Amide 2o Amide 3o Amide
R C
O
NH2
R C
O
NH
R'
R C
O
N
R'
R"
C
O
NH2
acetamide
Common Func?onal Groups (R, R’, R” = hydrocarbon radicals or H)
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Sulfide
dimethyl sulfide
Sulfoxide Sulfone Thiol
methanethiol
R S R'
H3C S CH3
H3C S CH3
O
R S R'
O
dimethyl sulfoxide
R R'
O
O
S2+
H3C CH3
O
O
S2+
dimethyl sulfone
R S H
H3C S H
Reactant/s Product/s
Law of ConservaAon of Mass “in a chemical reac?on, mass can be neither created nor destroyed”
Balanced chemical equaAon
Δ 2H2(g) + O2(g) 2H2O(l)
elect 2H2O(l) 2H2() + O2()
Stoichiometric number, states, special symbols
3N2(g) + 5H2O(g) + 7CO(g) +7C(s) 2