Chemical Bonding...The force of attraction of an element’s nucleus for electrons is called...
Transcript of Chemical Bonding...The force of attraction of an element’s nucleus for electrons is called...
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Chemical Bonding
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Compounds are formed from chemically bound atoms or ions
Substances become more stable through chemical bonding, where
2 or more atoms are joined together by a simultaneous
attraction.
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Valence electrons are electrons in the highest occupied energy level of an atom ( the last shell).
Bonding involves only valence electrons
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i. Ionic Bonding
ii. Covalent Bonding
iii.Metallic Bonding
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Ionic Bonds occur when the more electronegative element “steals” the electron pair away from the other atom.
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The atom that has stolen the electron pair becomes a negativeion (anion) while the “victim” becomes a positive ion (cation).
The two atoms are held together by their opposite charges.
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Na Cl
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Can you predict which atoms will gain electrons and which will looseelectrons by looking at the trend in electronegativity?
Increase in Electronegativity
Inc
reas
e in
Ele
ctro
nega
tivi
ty
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When you consider that for an ionic bond to form there must be agreat deal of difference in electronegativity between the atoms, can you predict what two types of atoms allow this to occur?
Metals Non-Metals
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Ionic Properties
Why do most ionic compounds have similar properties?
We can hypothesis that it is due to the bonds formed between the
ions, holding them firmly in a rigid structure
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Forming Ions
Metals tend to lose electrons and
become positive (cations)
Non-metals tend to gain electrons and
become negative (anions)
An ionic compound is formed when a
metal bonds to a non-metal. The total
charge of any compound is zero.
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Ex:
Na and Cl which one will lose
electrons which one will gain
electrons
Write out Tin’s electron
configuration what will it do??
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Lattice energy
measure of the energy released when
ions are combined to make a compound
Directly related to Coulombs Law, the
potential energy (E) between two ions is
directly proportional to the product of
their charges and inversely proportional
to the distance of separation between
them.
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The anions and cations in an ionic compound are locked in a regular neutrally charged structure, held by the balance of attractive bonds and electrical repulsion.
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The component ions in such crystals are arranged in repeating three-dimensional
(3-D) patterns.
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Alkali metals combine with halogens in 1:1 ratios since alkali metals need to lose 1 e1- and halogens need to
gain 1e1-.
Alkaline earth metals combine with halogens in 1:2 ratios since alkaline earth metals need to lose 2 e1- and
halogens need to gain 1e1-.
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LEWIS DOT STRUCTURES: Elements
Board Practice
Elements #1-20
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Lewis Structures can be used to illustrate the formation of ionic
bonds.
Be 2 F+ [Be]2+F F1- 1-
Write an equation with electron dot diagrams to illustrate the
formation of aluminum chloride.
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Al 3 Cl+ [Al]3+Cl Cl
1- 1-
Cl1-
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Lewis Structures
Duet Rule = applies to H and He and states these two atoms are stable with 2 electrons in their outer shell
Octet Rule= elements are most stable with 8 electrons in their outer shell
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1. Most are crystalline in structure
2. High melting/boiling points
3. Electrically neutral
4. Can conduct electricity when
melted or in aqueous solution
5. Hard/ Brittle
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Electrostatic attraction force
between the cation and free
electrons.
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Any successful bonding model for
metals must account for the typical
physical properties of metals:
malleability, ductility, and efficient and
uniform conduction of heat and
electricity in all directions.
Most metals are durable and have high
melting points.
These facts indicate that the bonding in
most metals are strong and non-
directional.
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Metal atoms are arranged in very
compact and orderly patterns.
i) Body-centered cubic
ii) Face-centered cubic
iii) Hexagonal close-packed
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1. Can conduct electricity (free electrons)
2. Malleable (put into shape)
3. Ductile ( made into wires)
4. Good conductors of heat
5. Metals are usually shiny
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Covalent Bonding
Br + Br Br Br
O + O O O
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Covalent Bonding
Electrons are shared by
nuclei Electrons can be shared equally (
non-polar covalent) or unequally
(polar covalent)
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Types of Bonds1) Single bond – 1 pair of e- are shared
- weakest bond
-longest bond length
2) Double bond – 2 pairs of e- are shared
3) Triple bond- 3 pairs of e- are shared
- strongest bond
- shortest bond length
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If the sharing is equal, this is called
Non – polar COVALENT BONDING
H HElectron pair
H H
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when the sharing
is unequal, one
atom becomes
slightly positive
the other slightly
negative creating
a polar covalent
bond
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The force of attraction of an element’s nucleus for electrons is called electronegativity. Atoms of different elements have different electronegativities. The higher the electronegativity, the stronger the attraction for electron pairs.
HFDifference
in En?
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The bonding electrons are on the average closer to the fluorine than to the hydrogen atom.
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The movement of the negatively charged electrons away from hydrogen toward fluorine, due to a difference in electronegativity, builds up a partial negative charge on the fluorine and a partial positive charge on the hydrogen.
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This is not a complete transfer of an electron from hydrogen to fluorine; it is merely a drifting of electrons toward fluorine.
H
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HWhen a charge separation of this
type is present, the molecule possesses an electric dipole, and
the bond is called a POLARCOVALENT BOND , or simply
a POLAR BOND.
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HPolar covalent bond (polar bond)
covalent bond joins two atoms of
different elements and the bonding
electrons are shared unequally
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Non-polar covalent bond
bonding electrons are shared
equally
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1. Soft , brittle solids
2. Low boiling/melting points
3. Tend to be more flammable
4. Do not conduct electricity
5. Usually non-soluble in water
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Lewis Structures for
Molecular Compounds
N N
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Bonding capacity is the number of covalent bonds (shared electron pairs) that an atom can form.
Covalent molecules often consist of atoms of different elements, with
different bonding capacities.
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Exceptions to the Octet Rule
1) B and Be usually have less than 8
electrons
2) Elements in the 3rd energy level and
above can have more than 8 electrons
in their outer shell
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1. Determine the central atom.
The least electronegative element
The atom with the smallest number of
valence electrons
The “oddball” element
It is NOT Hydrogen
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2. Count the number of valence electrons
in each atom.
Ex: NO3- N = 5 O = 6
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3. Include the number of electrons lost or
gained due to charge.
Ex: NO3- N = 5 O = 6 -ve = 1
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4. Total the number of valence electrons.
Ex: NO3-
N 5
3O 18
-ve 1
24
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5. Draw the skeletal structure with the
central atom in the middle.
6. Draw the first bonds connecting all
atoms to the central atom.
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7. Subtract the number of electrons
involved in these bonds (two electrons for
each bond) from the total number of
electrons.
24 – 6 = 18
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8. Finish adding electrons around the
atoms up to the number of electrons left.
Work from the outside in.
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9. Make sure all atoms satisfy the octet
rule. If they don’t move electrons around to
form multiple bonds (double or triple) so
that everything has a full octet
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10. If it is an ion, put [square brackets]
around the diagram and put the charge
outside.
[ ]-
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Resonance
= occurs when more than
one valid Lewis structure
can be written for a
particular molecule
Ex. NO3-1
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Formal Charge
charge assigned to an atom in a molecule,
assuming that electrons in all chemical
bonds are shared equally between atoms,
regardless of relative electronegativity,
helps determine the most correct structure
of a molecule when there is more than one
option.
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Formal Charge
FC= # of valence electrons – (# of covalent bonds) –
(# of electrons in lone pairs)
The total formal charge should add up to
the charge on the molecule
The most correct structure will be the one
with the lowest formal charge on each
element in the molecule
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Draw Lewis Structures for the following:
H2O, NF3, Cl2, SnCl2, PCl5, SO3,
BeCl2, C2H6, C2H2, ClF3, CHCl3, ICl,
O2, N2, SF6, CO2, BF3, C2H4, O3, IF7
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Compounds are arranged
in many different shapes
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The VSEPR Theory states that
because electron pairs repel,
molecular shape adjusts so the
valence-electron pairs are as far
apart as possible.
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VSEPR model
Valence Shell Electron Pair
Repulsion
Used to predict the geometry
of molecules
The structure will minimize
electron pair repulsions
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Types of Molecular Shapes
There are several classifications of
molecular geometry based on the number
of electron domains:
2 - Linear
3 - Trigonal planar
4 - Tetrahedral
5 - Trigonal bipyramidal
6 - Octahedral
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Intermolecular forces play a key
role in determining the physical
and chemical properties of
covalent compounds.
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Van Der Waals consists of 2 possible types of forces:
1. London Dispersion Forces
2. Dipole-Dipole Forces
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-This is the only type of force present in non-polar covalent molecules.
-It is the weakest of the intermolecular interactions caused by the motion of the electrons.
-The strength of dispersion forces generally increases as the number of electrons in a molecules increases because the molecule is more polarizable.
Ex. Halogen diatomic molecules.
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- This occurs when polar covalent bonds are attracted to one another.
- Electrostatic attractions occur between oppositely charged regions. (partially (–) and partially (+)).
- Dipole interactions are similar to but much weaker than ionic bonds.
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Dipolar molecules
Have a center of positive charge and a center of negative charge.
aka: dipole moment
Ex. HF
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Dipole moment in NH3
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Dipole cancels out in CO2
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This is found in polar covalent molecules that
have hydrogen that is bonded to a very
electrostatic element (N, F, O) and attracted to a
(N, O, F) in another adjacent molecule.
Hydrogen bonds are the strongest of the
intermolecular forces.
Hydrogen > dipole-dipole > London Dispersion
Bonds interactions forces
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Hydrogen bonds are extremely important in
determining the properties of water and biological
molecules such as proteins.
The water molecule has a bent shape (105°) and
is considered to be polar and the universal solvent.
The attraction in water results from the
intermolecular hydrogen bonds.
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Surface tension: the inward force, or pull that tends to minimize the surface area of a liquid
- this surface tension tends to hold a drop of liquid in a spherical shape
The higher the surface tension, the more nearly spherical is the drop of that particular.
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Because of hydrogen bonding, water absorbs
a large amount of heat as it evaporates or
vaporizes.
The hydrogen bonds must be broken before
water changes from the liquid to vapor state.
Vapor Pressure the force exerted due to the
gas above the liquid
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Boiling Point: occurs when the
temperature at which the vapor pressure
of the liquid is just equal to the external
pressure.
Boiling leads to evaporation of a liquid.
In the case of water, hydrogen bonds
break in order for the liquid to vaporize.