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