Chapter 1 Aisyah SK
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Transcript of Chapter 1 Aisyah SK
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CHAPTER 1 STRUCTURE AND BONDING
AISYAH SALIHAH KAMAROZAMAN
013-2119614
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The Ancient Greeks used to believe that everything was
made up of very small particles. I did some experiments in
1808 that proved this and called these particlesATOMS:
Dalton
shellproton
N
++N
--
neutronelectron
1.1 Atomic Structure
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1.1 Atomic Structure Atom: Nucleus – neutrons (neutral) + protons (+ve), the mass of the
atom. Electron - Negatively charged, negligible mass
Atom is neutral, No. of protons = no. of electrons
Diameter of an atom is about 2 10-10 m (200 picometers
(pm)) [the unit angstrom (Å) is 10-10 m = 100 pm]
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Atomic number (Z) - number of protons in the atom's nucleus
Mass number (A) - number of protons plus neutrons
(protons + neutrons) mass number
Isotopes - atoms of the same element but different numbers of neutrons different mass numbers
Atomic mass (atomic weight) - weighted average mass in atomic mass units (amu) of an element’s naturally occurring isotopes
A
Z X(protons only) atomic number
1 321H
1H1H
hydrogen deuterium tritium
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PRINCIPAL QUANTUM NUMBER, n
An orbital describes a region in space where an electron
may be found
n can be any integer 1
The larger the value of n, the more energy the orbital has
The larger the value of n, the larger the orbital
Orbitals with the same value of n are in the same principal
energy level (or same shell)
1.2 Atomic Structure: Orbitals
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n = 1, 2, 3, 4, ….
n=1 n=2 n=3
distance of e- from the nucleus
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ANGULAR MOMENTUM QUANTUM NUMBER, l
• The angular momentum quantum number determines the
shape of the orbital
• l can have integer values from 0 to (n – 1)
• Orbitals with the same values of n and l are said to be in the
same sublevel (or same subshell)
• For a given value of n, l = 0, 1, 2, 3, … n-1
n = 1, l = 0
n = 2, l = 0 or 1
n = 3, l = 0, 1, or 2
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4 kinds of orbitals: s, p, d, and f
Orbitals are organised into different layers – electron
shells
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l = 0 (s orbitals)
1 orbital diagram
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l = 1 (p orbitals)
3 orbital diagram
3 p orbitals within a given shell are oriented in perpendicular directions: px, py, and pz
Lobes of a p orbital are separated by region of zero electron density, a node
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l = 2 (d orbitals)
5 orbital diagram
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f orbitals
7 orbital diagram
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MAGNETIC QUANTUM NUMBER, ml
The magnetic quantum number is an integer that
specifies the orientation of the orbital
the direction in space the orbital is aligned relative to the
other orbitals
Values of ml are integers from −l to +l including zero
For a given value of l, ml = -l, …., 0, …. +l
if l = 1 (p orbital), ml = -1, 0, or 1
if l = 2 (d orbital), ml = -2, -1, 0, 1, or 2
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SPIN QUANTUM NUMBER ms
ms = +½ or -½
ms = -½ms = +½
• Describes the spin of a specific
electron and properties of
electron
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Orbitals and Shells Orbitals are grouped in shells of increasing size and energy Different shells contain different numbers and kinds of orbitals Each orbital can be occupied by two electrons 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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1.3 Atomic Structure: Electron
Configurations
Ground-state electron configuration (lowest energy arrangement) – listing of orbitals occupied by the electrons.
Rules 1:
Lowest-energy orbitalsfirst: 1s 2s 2p 3s 3p 4s 3d (Aufbau(“build-up”) principle)
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Rules 2:
Electrons act as they are
spinning around axis,
2 spin orientations: or
Only 2 electrons occupy an
orbital
Pauli exclusion principle
Rules 3:
If 2 or more empty orbitals
of equal energy are
available, 1 electron
occupies with parallel
spins until orbitals are
half-full
Hund’s rule
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1.4 Development of Chemical Bonding
Theory
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 (example: Na+Cl-)
Organic compounds have covalent bonds from
sharing electrons (G. N. Lewis, 1916)
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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
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Lewis structures (electron dot) show valence electrons of an atom as dots Hydrogen has one dot, representing its 1s electron
Carbon has four dots (2s2 2p2)
Kekule structures (line-bond structures) have a line drawn between two atoms indicating a 2 electron covalent bond.
Method of indicating covalent bonds
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Note that:
Stable molecule = noble gas configuration
completed shell, octet (eight dots) for main-group atoms, two for hydrogen
The number of covalent bonds an atom form depends on how many
additional valence electrons to reach noble-gas configuration.
Carbon has four valence electrons (2s2 2p2), forming four bonds (CH4).
Nitrogen has five valence electrons (2s2 2p3) but forms only three bonds (NH3).
Oxygen has six valence electrons (2s2 2p4) but forms two bonds (H2O)
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Nonbonding electron
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 electron
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Drawing Structures
Several shorthand methods have been developed to write
structures.
1)Condensed structures don’t have C-H or C-C single bonds
shown. ( ) used to indicate more than one identical group
attached to same previous central atom
2)Skeletal structures don’t show carbon and hydrogen atom or
C-H bonds unless they are part of a functional group
Skeletal structure
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Skeletal structure
3 General Rules:
1) Carbon atoms aren’t usually shown. Instead a
carbon atom is assumed to be at each intersection
of two lines (bonds) and at the end of each line.
2) Hydrogen atoms bonded to carbon aren’t shown.
3) Atoms other than carbon and hydrogen are shown.
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1.5 The Nature of Chemical Bonds:
Valence Bond Theory
2 models to describe covalent bonding:
Valence bond theory (hybridization)
Molecular orbital theory
2 types of bond : sigma () bond and pi () bond
A sigma () bond : the interacting atomic orbitals point along
the axis connecting the two bonding nuclei (end-to-end overlapping)
A pi () bond : the bonding atomic orbitals are parallel to each
other and perpendicular to the axis connecting the two bonding
nuclei (side-way overlapping)
between unhybridized parallel p orbitals
The interaction between parallel orbitals is not as strong as
between orbitals that point at each other; therefore bonds are
stronger than bonds
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Types of Bonds
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1.5 The Nature of Chemical Bonds:
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
H-H bond is cylindrically symmetrical, bond formed by the
head-on overlap of 2 atomic orbitals are called sigma ()
bonds
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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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# of Lone Pairs +
# of Bonded AtomsHybridization
2
3
4
5
6
sp
sp2
sp3
sp3d
sp3d2
How do I predict the hybridization of the central atom?
1. Draw the Lewis structure of the molecule.
2. Count the number of lone pairs AND the number of atoms
bonded to the central atom
Geometry
Linear
Trigonal planar
Tetrahedral
Trigonal bipyramidal
Octahedral
1.6 Hybridization
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1.6 sp3 Orbitals Carbon has 4 valence electrons (2s2 2p2)
In CH4, all C–H bonds are identical (tetrahedral)
sp3 hybrid orbitals: s orbital and three p orbitalscombine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3), Pauling (1931)
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sp3 Hybridization of a Carbon Atom
1s2 2s2 sp2Ground state of C
px py pzs
2s12p3Promotion of electron
s px py pz
sp3 hybridized state sp3
sp3
tetrahedralFour sp3 orbitals of equal length, energy and in shape
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Structure of methane
sp3 orbitals on C overlap with 1s orbitals on 4 H atoms to form four identical C-H bonds
Each C–H bond has a strength of 436 (438) kJ/mol and length of 109 pm
Bond angle – bonds specific geometry, methane: each H–C–H is 109.5°, the tetrahedral angle.
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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 orbitalsto form six C–H bonds
C–H bond strength in ethane 423 kJ/mol C–C bond is 154 pm long and strength is 376 kJ/mol All bond angles of ethane are tetrahedral value of 109.5°
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1.7 sp2 Orbitals sp2 hybrid orbitals: 2s orbital combines with two 2p
orbitals, giving 3 orbitals (spp = sp2). This results in a double bond.
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 format 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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sp2 Hybridization of a Carbon Atom
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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 134 pm (C–C 154 pm)
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Bonding in Ethylene
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1.8 sp Orbitals C-C a triple bond sharing six electrons
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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sp Hybridization of a Carbon Atom
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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 and C-H bonds
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1.9 Hybridization of Nitrogen and
Oxygen Elements other than C can have hybridized orbitals
H–N–H bond angle in ammonia (NH3) 107.3°
C-N-H bond angle is 110.3 °
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 and CH3.
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1.10 The nature of chemical
bonds: Molecular Orbital Theory
A molecular orbital (MO): describe a region of space where electrons are most likely to be found (specific energy and general shape) in a molecule
Additive combination (bonding) MO is lower in energy
Subtractive combination (antibonding) MO is higher energy
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Molecular Orbitals in Ethylene
The bonding MO is from combining p orbital lobes
with the same algebraic sign
The antibonding MO is from combining lobes with
opposite signs
Only bonding MO is occupied
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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
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Valence bond theory - electron sharing occurs by overlap
of two atomic orbitals
Molecular orbital (MO) theory, - bonds result from
combination of atomic orbitals to give molecular orbitals,
which belong to the entire molecule
Sigma () bonds - Circular cross-section and are formed by
head-on interaction
Pi () bonds – “dumbbell” shape from sideways interaction of
p orbitals
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Summary
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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Exercises
1) Give the ground state
following elements:
electron configuration for each of the
a)
c)
What
Potassium
Aluminium
is the hybridization
b)
d)
of
Arsenic
Germanium
each carbon atom in acetonitrile,2)
C2H3N?
Fill in any nonbonding valence
the following structures?
3) electrons that are missingfrom
-NH2 O
O OS CH3H C S3 CH3 CH3
Acetate ion
Dimethyldisulfide Acetamide
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Exercises
4) Convert the following molecular formulas into line-bond
structures
a) C2H6O
c) C2H4O
that are consistent with valence rules:
b)C3H7Br
d) C3H9N
5) What kind of
atom in the
a) Propane
hybridization do you expect for each carbon
following molecules?
b) 2-methylpropane
c) Acetic acidc) 1- butane-3-yne