Transcript Gen. Chem. Review notes

G&G Fig. 1.8 pg.9
The inorganic Precursors:
(18-64 daltons)
Carbon Dioxide, Water, Ammonia
Nitrogen (N2), Nitrate (NO3-)
Metabolites:
(50-250 daltons)
Pyruvate, Citrate, Succinate,
Glyceraldehyde-3-phosphate,
Fructose-1,6-bisphosphate,
3-Phosphoglyceric acid
Molecular Organization
in the Cell
is a Hierarchy
The Cell
Building Blocks:
(100-350 daltons)
Amino acids, Nucleotides,
Monosaccharides, Fatty acids,
Glycerol
Macromolecules:
(103-109 daltons)
Proteins, Nucleic Acids,
Polysaccharides, Lipids
Supramolecular
Complexes:
(106-109 daltons)
Ribosomes, Cytoskeleton,
Multi-enzyme complexes
Organelles:
Nucleus, Mitochondria,
Chloroplasts, Endoplasmic
reticulum, Golgi Apparatus,
Vacuole
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General Chemistry Review
Atomic Theory and Structure (3.1-3.4)
All matter is composed of atoms
atoms of elements differ from each other
compounds consist of atoms combined in whole number ratios
chemical reactions do not change atoms - just the way they are combined
 Composition of Atoms
Subatomic particles
protons - carry positive charge - inside nucleus
electrons - carry negative charge - surrounds the nucleus
neutrons - uncharged - inside nucleus
atoms are held together by attractive forces between electrons and
protons
#p=#e - atom electrically neutral
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 Elements and atomic number
Atomic number (Z)
The number of protons in an atom
all atoms of same element have same Z
in neutral atom, also equals # electrons
When does Z not equal the # electrons also?
in an ion
Mass number (A)
#p + #n
In phosphorus, Z=15. How many protons, electrons, and neutrons are there in
phosphorus atoms that have a mass number of A=31?
#p=#e=15, #n=31-15=16
An atom contains 28 protons and has A=60. Give the number of electrons and
neutrons in the atom, and identify the element.
Z=28 so #p=#e=28, A=60 so #n = 32. Ni has Z=28
Isotopes - atoms of same element - have same Z but different A - different # n
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Periodic Table of Elements
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Atomic Structure (3.6)
 Shell - same as the row number (period number)
the farther the shell is from the nucleus, the larger it is, the more
subshells it holds, the more electrons it can hold, and the higher the
energies of those electrons.
 Subshell - 4 different types of subshells:
subshells consist of different # of orbitals
each orbital can hold a maximum of 2 electrons
therefore, each subshell can hold a different maximum
# orbitals
max # electrons
s
1
2
p
3
6
d
5
10
f
7
14
Distance from nucleus
“general” order of subshell energy
The further an electron is from the nucleus, the higher its energy!
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 Electron distribution in atoms
Shell number
1
2
3
4
subshell designation
s
s, p
s, p, d
s, p, d, f
# orbitals
1
1 3
1 3
5
1 3
5
max # electrons
2
2
2
10
2
10 14
total electron capacity:
2
6
8
6
18
6
7
32
How many electrons are present in an atom that has its first and second shells
filled and has 4 electrons in its third shell? Name the element.
1st shell = 2 e
2nd shell = 8 e
3rd shell = 4
total e = 14 so atom has Z = 14 which is Si
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Electron Configuration - An electron’s address (3.7-3.8)
 Aufbau Principle - electrons fill lower energy orbitals
 Pauli Exclusion Principle - electrons in same orbitals have different spins
 Hund’s Rule - degenerate orbitals are filled equally
Orbitals of equal energy. Must be
same shell and subshell.
2p
2p
2p
2s
2s
2s
1s
1s
1s
H
He
Li

1s1
7
2p
2p
2p
2p
2s
2s
2s
2s
1s
1s
1s
1s
He
Be
B
C
2p
2p
2p
2p
2s
2s
2s
2s
1s
1s
1s
1s
N
O
F
Ne
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Tis Noble to Have a Complete Octet

Valence Shell
Outermost shell of electrons - responsible for reactivity of the atom

The Octet Rule
Atoms exchange electrons or share electrons to complete the valence shell (8 e for
main group)
To form covalent
bonds
To form ionic
bonds

review biologically
important ions on
page 85
Ionic Bonds (4.1-4.5)
Typically formed between highly electronegative and electropositive atoms (metal to non metal)
Groups 1A and 2A atoms lose valence electron(s) to form cations
Group 7A atoms (the halogens) gain an electron to form anions
Ions are electrostatically attacted to one another and form an ionic bond
Li
+
F
Li
+
F
See “Biologically Important Ions” on Page 84
LiF
lithium fluoride
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Bonding is stronger when sharing occurs

Covalent Bonds (5.1-5.4)
Atoms share valence shell electrons to form covalent bonds
generally nonmetal to nonmetal
stronger than ionic bonds
The number of bonds an atom makes depends on the number of electrons it
needs to obtain an octet (or a duet in hydrogen’s case).
2p
2p
2p
2p
2s
2s
2s
2s
1s
1s
1s
1s
H
C
N
O
Atom
# bonds
# lone pairs
H
1
0
C
4
0
N
3
1
O
2
2
Stable with 6
Stable with 10
There are exceptions to the octet rule: B, P, and S
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Examples of molecules with covalent bonding:
H2O
NH3
CH3OH
CH4
What are the likely formulas for the following molecules? Why?
Draw their lewis dot formulas.
CH2Cl?
BH?
NI?
CH3CH?
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Drawing Lewis Structures (5.6)

Lewis structures show the arrangement of electrons
Step 1
Find total # valence electrons - add one for a
neg charge and subtract one for +
Step 2
Draw a line btw connected atoms to represent
two electrons
Work example 5.6
and problem 5.13
Step 3
Add lone pairs so each atom (except the central one and H)
has an octet
Step 4
Put all electrons left on the central atom
Step 5
If central atom still doesn’t’ have 8, make multiple
bonds
Step 6
Check chart on page 9 of notes
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Sometimes once isn’t enough

Multiple Covalent Bonds
>1 electron pairs must be shared to complete an octet for some molecules
single bond - one electron pair shared
double bond - two electron pairs shared
triple bond - three electron pairs shared
order of bond stability single < double < triple - more electrons being shared
Examples of molecules with multiple covalent bonding:
CO2
N2
CH3COOH
CH2O
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Converting between Condensed and Lewis Structures (5
Lewis structure
shows the connectivities of atoms with dashes and shows lone pairs
use ammonia and acetaldehyde (CH3CHO) as examples
Structural formula
shows the connectivities of atoms with dashes but lone pairs are not shown
Condensed formula (molecular formula)
shows number and kinds of atoms in a molecule
If the condensed formula is given, draw the structural formula and vice versa:
CH2CHCl
H
H
C
H
C
H
H
C
H
CH3COOH
H
C
C
N
H
H
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Shapes of Molecules (5.6)

Valence Shell Electron Pair Repulsion (VSEPR)
electrons in bonds and lone pairs try to get as far apart from each other as possible
giving molecules specific shapes
To predict shape:
Step 1
draw lewis structure
Step 2
Count # of electron clouds (bonds plus lone pairs)
Step 3
Predict shape based on table 5.1 pg 113
Examples:
15
16
bent
tetra
tetra
Trig planar
pyramidal
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Electronegativity (Electron Selfishness) (5.8-5.9)
ability of an atom to draw electrons towards itself
electronegativity

As you go across a row, electrons are added to the same shell but the
nuclear charge is increased. The greater the amount of positive charge in
the nucleus, the greater is its ability to draw electrons to itself. Therefore,
the greater the electronegativity.
The more
shells that are
occupied, the
more shielded
the nucleus
(the positive
charge) is. The
positive charge
in the nucleus
is the
attraction
for the
electrons.
Therefore, the
less shielded
atoms (like F),
can
attract
electrons more
and thus
have a greater
electronegativi
ty.
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
Polar Covalent Bonds
covalent bond between atoms with different electronegativities
unequal sharing of electrons
d-C Hd+
D=0.4
+
dC
Nd
D=0.5
Non-polar bond

+
dN
Hd
+
dC
Od-
dO
Hd+
D=1.4
D=1.0
D=0.9
d+
Polar bond
d-
Polar Molecules
Symmetrical molecules are non-polar even if they have polar covalent bonds
Cldd+
+
H C N
d
d
O C O
d-Cl Cd
O
Cl
Cl
ddH d- H
dNet polarity
d
+
d
d+
Molecules are symmetric
Equal bond dipoles in opposite directions =
no net polarity
Polar covalent bonds are very important in biochemistry. They create19a
charged environment that helps molecules identify and interact with each other.
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