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Chemical Bonding

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Julia Saunders

on 11 December 2013

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Transcript of Chemical Bonding

Chemical Bonding
atoms are found in groups, held together by electrostatic forces of attraction
isolated atoms are rarely found in nature (except noble gases)
Molecular Architecture
VSEPR Theory

V - valence
S - skill
E - electron
P - pair
R - repulsion
Valence Bond Theory
developed by Linus Pauling
a 1/2-filled orbital in one atom can overlap w/ 1/2-filled orbital of a different atom to form a new bonding orbital
the new bonding orbital contains a pair of electrons of opposite spin
each bonding orbital can have no more than 2 electrons
the bonding atoms arrange themselves in space to achieve a maximum overlap to produce a bonding orbital of lowest energy
energy is released to get lowest energy and maximum stability
hybridization is a concept that was required to explain the equal bonds of carbon-based molecules and the existence of double and triple bonds
explains how CH4 molecules have identical orbitals with the same characteristics (they repel each other equally w/ 109.5 degree bond angles
this occurs w/ bonding not isolated atoms
Lewis Theory
atoms and ions are stable if they have a noble-gas-like electron structure
electrons are most stable when they are paired (w/ some exceptions)
atoms form chemical bonds to achieve a stable octet of electrons
a stable octet may be formed when electrons are exchanged between metal and non-metal atoms or by non-metals sharing electrons
Bonding
Covalent Bonding
atoms in molecular substances might be held together by an electrostatic force of attraction between atomic nuclei
electron pairs are shared by the atoms
Lewis (1875-1946) proposed this type of bonding
molecules are groups of electrically neutral atoms that are held together tightly enough to be considered one particle
octet rule: when atoms combine they try to form as many octet structures as possible (a full valence shell with 8 electrons
chemical bonds are electrostatic forces of attraction which hold atoms, ions or molecules together
Ionic Bonding
involves the electrostatic forces of attraction between oppositely charged ions
metals will donate a valence electron to a non-metal, forming an ionic bond
valence electrons are in the outermost energy level of an atom
metals have a lower ionization energy (energy needed to remove an electron)
non-metals have higher electron affinity (ability to attract electrons and the energy released to make a negative ion)
ionic substances conduct electricity in aqueous solution
Ex. [Na] + [Cl]
NaCl
+
_
electron transfer forms an ionic bond
Covalent Bond
when orbitals overlap and electrons are found in the same space sharing atomic nuclei
energy is released and the molicule becomes more stable and less reactive
potential energy decreases when the atoms bond
lowest energy = most stability
if atoms are forced together too closely they will repel
Polar Bonds vs. Polar Molecules
Polar Bonds
Linus Pauling combined the concepts of bond energies and the valence bond theory to create the idea of "electro-negativity"
the values of electro-negativity (EN) were arbitrarily assigned to make Fluorine (F) was most electro-negative w/ a value of 4.0
EN increases moving upward and right along the periodic table of elements
the polarity of a chemical bond and the electro-negativity difference (E.N.D.) are related since a greater E.N.D. results in a more polar bond
the polarity of a molecule is equal to the geometric sum of the polarities of the individual bonds
ex.
the bond between atoms in a hydrogen chloride (H - Cl) molecule is polar since the electro-negativities (EN) are different
Hydrogen (H) has an EN value of 2.1 and chlorine (Cl) has an EN value of 3.0
the electrons are not shared equally and a molecular dipole is created
the H atom has a partially positive charge and the Cl atom has a partially negative charge
O=C=O Carbon dioxide (CO2) contains polar bonds but the molecule is non-polar because the net polarity is zero
the electro-negativities of the Oxygen (O) atoms are equal and the linear shape of the molecule ensures that one side is not partially negative or partially positive
General Rule
a molecule is non-polar if the atoms around the central atom are identical and symmetrically arranged
free electron pairs that create bent molecules are able to destroy symmetry; therefore the molecule is polar
water molecules (H2O) are polar so they bond w/ other H2O molecules since oxygen is slightly negative and hydrogen is slightly positive
these properties allow water to stay a liquid through extensive temperature changes
Polar Molecules
a bond dipole is the electro-negativity difference (E.N.D.) of 2 bonded atoms represented by a lower (slightly positive) to a higher (slightly negative) EN
carbon dioxide (CO2) is a non-polar molecule w/ polar bonds ; the net polarity is zero since it is a symmetrical molecule
water (H20) is a polar molecule because the bond dipoles do not cancel each other out
to determine molecular polarity, consider bond dipoles and the E.N.D. of bonds
Electro-negativity Difference
E.N.D. > 1.7 results in an Ionic Bond
E.N.D. 1.7 - 0.1 results in a Polar Covalent Bond
E.N.D = 0 results in a Non-Polar Covalent Bond
Intermolecular Forces
forces between individual molecules in the liquid or solid states of a covalent (molecular) compound; Van der Waal's Forces
London Forces
instantaneous dipole interactions
they exist due to the creation of momentary charge imbalances
the forces are very weak
the strength of the force is related directly to the number of electrons in the molecule
more electrons results in stronger forces

ex. oxygen gas (O2) O=O one side may be slightly negative and the other slightly positive for an INSTANT
Dipole-dipole Forces
permanent dipole interactions
molecules w/ permanent dipoles have 1 slightly positive and 1 slightly negative end
the strength of the force is determined by the polarity of the molecules
Hydrogen Bonding
the attraction of Hydrogen (H) atoms to bond w/ nitrogen (N), oxygen (O) or fluorine (F) atoms
the H atom is attracted to the lone pair of electrons from these atoms as they exist in adjacent molecules
inter-molecular forces are between molecules
intra-molecular forces are between atoms w/ in a molecule
intermolecular forces effect physical properties of substances
hydro-philic substances dissolve in water and hydophobic substances do not
melting and boiling points are also determined as well as surface tension
Bond Models
Molecular Crystals
crystals that are soft and have low melting points due to weak intermolecular forces
individual particles are molecules (groups of covalently bonded atoms
the substance is neutral therefore not electrically conductive
giant polymers (plastics and hydrocarbons (wax) is excluded
ex. Iodine
Network (Covalent) Crystals
these crystals are macromolecules
an interlocking structure makes the covalent network bonding between atoms strong
the substances are very hard crystals w/ a high melting point
electrons are held w/ in atoms or w/ in the covalent bonds
the electrons cannot move through the network so these substances are non-conductive
ex. carbon (C) and silicone dioxide (SiO2)
Metallic Crystals
non-directional bonds are caused by electrostatic attractions of positive centres surrounded by a negative field of electrons
they are malleable since the planes of atoms can slide over each other while remaining bonded
the attraction between the cations (positively charged particles) and the negative electron field is strong so the crystalline structure can be hard and solid
the melting point and hardness range from low to high
this field of electrons allows valence electrons from the cations to move freely throughout the structure so electricity can be conducted as electrons move in and out of the substance
valence electrons absorb and re-emit energy from all wavelengths of visible and near visible light so the crystals are shiny
Ionic Crystals
conducts electricity in liquid and solution due to dissociation of ions
strong ionic bonds results in the hardness and high melting point
directional bonds result in the brittleness; as ions shift they repel each other
crystal lattice structure creates full ionic charges which are stronger than hydrogen bonding
ex. sodium chloride (NaCl)
ex. aluminum (Al) and copper (Cu)
Orbital Hybridization
all electron charge clouds (orbitals) of the outer energy level of an atom hybridize (change or blend) to become identical to each other during the bonding process
we imagine orbitals w/ identical characteristics positioned as far away from one another as possible
the electrons are promoted to a higher energy level then they hybridize
this only occurs in 1 energy level; whichever level happens to be the valence shell
Practical Link
Biological catalysts are enzymes
they catalyze and control biochemical reactions
enzymes have a specific shape which is why molecular structure is so important
the reactant is called a substrate
specific reactants fit into a complementary enzyme
this "lock and key " method allows a reaction to occur
Co-ordinate Covalent
Bonding

1 atom supplies both bonding electrons
ex. ammonia bonding w/ a hydrogen ion to create the ammonium ion
Resonance Structures
developed to explain experimentally determined bond lengths and strengths
expressing 1 1/3 bonds as if is spreads out over the entire structure
the molecule needs a double bond w/ 1 or more single bonds
usually occurs w/ oxygen and carbon or oxygen and sulfur
hydrogen cannot exhibit resonance since it makes only single bonds
H

H N : + H (ion)

H

H

H N H

H
+
: O :

S O

: O :

. .
. .
. .
. .
S+
S-
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