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OVERVIEW Valence electrons Ions (cations/anions) Types of bonds
Ionic, covalent Properties of bonds
Bond strength, hardness/texture, melting point/boiling point, conductivity
Metallic Bonds Lattice Energy Electronegativity & Bonds Lewis Dot Structures
Octet Rule, covalent compounds, exceptions, ions, resonance, formal charge, isomers
Bond order, energy, length
Ionic Lewis Structures VSEPR Theory
Molecular Geometry Intermolecular Forces
Hydrogen Ion-dipole Dipole-dipole London dispersion
Valence Bond Theory Hybridization Sigma/pi bonds
HOW ARE COMPOUNDS FORMED?
Electrons are lost, gained, or shared.
Valence electrons: the electrons available to be lost, gained, or shared Number of valence electrons is electrons in s and
p orbitals of highest energy level (outermost shell)
Main groups on table correspond to number of valence electrons.
IONS
Atom(s) that has a negative (-) or positive (+) charge
Cation: positively charged ion
Anion: negatively charged ion
Na Na+ + e-
e- + Cl Cl-
IONIC BOND Electrostatic attraction between cation and anion
(opposite charges attract) For example, an uncharged chlorine atom can pull
one electron from an uncharged sodium atom, yielding Cl−and Na+.
Cation becomes smaller and anion becomes larger Typically between a metal and nonmetal
IONIC BONDS
Form crystal lattice Solid crystal at ordinary
temperatures Organized in a characteristic
pattern of alternating positive and negative ions
All salts are ionic compounds and form crystals
Ionic compounds do not exist as single molecules
COVALENT BOND
Bond involves sharing of electrons
Forms between atoms of two nonmetallic elements
Forms molecules
Nonpolar covalent bond: electrons are shared equally resulting in an even distribution of the negative charge (no partial negative charge)
Polar covalent bond: one atom in the bond attracts electrons more than the other atom making the electron negative charge shift to that atom giving it a partial negative charge
BOND STRENGTH
Strong bonds Crystal lattice
Repeating symmetrical pattern
Weaker bonds Covalent Network Molecules
Strong forces within molecule but weak between molecules
IONIC COVALENT
MELTING & BOILING POINT
High melting/boiling points Sturdy crystal lattice
Low volatility
Low melting/boiling points Weak forces between
molecules Most are gases High volatility
IONIC COVALENT
TEXTURE & HARDNESS
Brittle Sturdy but collapses
easily if disrupted Hard structure
Soft compounds Most are gases
IONIC COVALENT
CONDUCTIVITY – IONIC
Solids – not conductors Ions can’t move
Molten state – conductors Solution – conductors (ions
separate)
Not good conductors
IONIC COVALENT
METALLIC BOND
Results from attraction between metal cations and the surrounding “sea of electrons”
Vacant p and d orbitals in metal's outer energy levels overlap, and allow outer electrons to move freely throughout the metal
Valence electrons do not belong to any one atom
METALLIC BONDS Very strong bonds
Due to “sea of electrons” Highest melting/boiling points
Electrons result in strong forces holding together Low volatility Very hard
METALLIC BONDS
Malleable/ductile When struck, one plane of metal atoms can slide past
another without breaking Great conductors (heat and electricity)
Freedom of electrons carries current Shiny and have luster
Electrons absorb light, get excited, fall, re-radiate the light
LATTICE ENERGY
The energy required to separate ions of an ionic solid
Magnitude of lattice energy depends on Charges of ions Sizes of ions Arrangement in ions
in the solid
Lattice energyCompound kJ/molLiCl 834NaCl 769KCl 701NaBr 732Na2O 2481Na2S 2192MgCl2 2326MgO 3795
Lattice energyCompound kJ/molLiCl 834NaCl 769KCl 701NaBr 732Na2O 2481Na2S 2192MgCl2 2326MgO 3795
For given arrangement of ions, lattice energy
increases as the charges on the ions increase and as their radii decrease
ELECTRONEGATIVITY
Atom’s ability to attract electrons Difference in electronegativity values
between two elements can determine type of bond Ionic Bonding = Over 1.7 Polar Covalent = 1.7 - 0.4 Non-polar Covalent = 0-0.3
Examples KCl = 2.2 (ionic) CuS = 0.6 (polar covalent) O2 = 0.0 (nonpolar covalent)
LEWIS DOT NOTATIONS
Valence electrons are represented by dots drawn around the symbol of an element
1 valence e-
X2 valence e-
X3 valence e-
X4 valence e-
X5 valence e-
X6 valence e-
X7 valence e-
X8 valence e-
X
OCTET RULE
Compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has 8 electrons in its outer shell.
Hydrogen is exception and only needs 2 electrons for a complete shell
8Electrons!
LEWIS DOT STRUCTURES
Show electron distribution in compounds Lone pair: nonbonding pair of valence electrons Bond pair: valence electrons shared between
two elements Shared electrons pairs represented by two
dots (:) or by a single line ( - )
HCllone pair
shared orbond pair
•••• ••••
STEPS FOR BUILDING A DOT STRUCTURESTEPS FOR BUILDING A DOT STRUCTURE
Ammonia, NH3
1. Add up the total number of valence electrons that can be used.
H = 1 and N = 5
Total = 1 + 1 + 1 + 5 = 8 electrons
2. Decide on the central atom;
If carbon it is always central atom
If no carbon, choose the least electronegative atom
Never use hydrogen!
Therefore, N is central on this one
3. Draw the basic skeletal structure of the molecule.
H H
H
N
Building a Dot Structure
H H
H
N4. Draw in the electrons around the central atom. Then place the remaining electrons to form complete octets and lone pairs as necessary.
N has 3 BOND PAIRS and 1 LONE PAIR.
..
.. .. ..
5. Check to make sure there are 8 electrons around each atom except H.
Building a Dot Structure
6. Count electrons! Make sure that the number of electrons in your structure equals the total number from the beginning.
H HHN.... .. ..
H HHN.... .. ..Total Electrons = 8
EXCEPTIONS TO THE OCTET RULE
Element is 3rd period or higher, is the central atom, AND is bonded to electronegative atoms (such as O, F, Cl, Br) may have more than 8 electrons
Electrons use empty valence d orbitals
Be is stable with 4 electrons
B is stable with 6 electrons
LEWIS DOT STRUCTURES FOR IONS
Add or remove electrons based on charge of ion If the ion has a negative (-) charge add electrons
to the Lewis structure. If the ion has a positive (+) charge, then
subtract electrons from the Lewis structure.
Try CO3-2 O
O
C O
-2
SPECIES WITH AN ODD TOTAL NUMBER OF ELECTRONS A very few species exist where the total
number of valence electrons is an odd number
This must mean that there is an unpaired electron which is usually very reactive.
Radical Species that has one or more unpaired electrons They are believed to play significant roles in
aging and cancer. Example – NO
It has a total of 11 valence electrons: 6 from oxygen and 5 from nitrogen.
:N::O:
:.
RESONANCE STRUCTURES Occurs when more than one valid Lewis
structure can be written for a particular molecule.
The resonance structures are the same except for the placement of the electrons (meaning the bonds).
RESONANCE BONDS
Resonance bonds are shorter and stronger than single bonds.
Resonance bonds are longer and weaker than double bonds.
H
H
H
H
H
H
H
H
H
H
H
H
FORMAL CHARGES A bookkeeping system for electrons.
Does not give actual charge of atoms Helps decide which Lewis structure is more preferred
Used to show the approximate distribution of electron density in a molecule or polyatomic ion
Assign an atom electrons: Each atom gets half of the bonding electrons it has
(single, double, and triple bonds) Each atom gets all unshared (nonbonding) electrons
that are found on it
FORMAL CHARGES
To solve for formal charge:
(Total valence electrons) – (electrons assigned to atom)
To decide which structure is preferred:
Choose the Lewis structure in which the atoms bear formal charges closest to zero
Choose the Lewis structure in which any negative charges reside on the more electronegative atoms
FORMAL CHARGES
Example: CO2
For each oxygen4 electrons from unshared electrons2 from the bonds (1/2 of 4)6 total
Formal charge = 6 - 6 = 0
For carbon 0 unshared electrons
4 from the bonds (1/2 of 8)4 total
Formal charge = 4 - 4 = 0
O=C=O
ISOMERSIsomers – compounds whose molecules have
the same overall molecules but different structures
C C O
H
H
H
H
H
H C CO
H
H
H
H
H
H
C CO
H
H
H
H
H
HC C O
H
H
H
H
H
H
BOND ORDER
Refers to the average number of bonds that an atom makes in all of its bonds to other atoms Draw Lewis structures and determine bonds Single bond counts as 1 bond, double counts as
2, and triple counts as 3
Bond Order =
Examples: CH3Cl = 4/4 = 1 CO2 = 4/2 = 2
CO3-2 = 4/3 = 4/3 ClO4
-1 = 4/4 = 1
number of bondsnumber of atoms bonded
BOND ENERGIES Bond energies and lengths differ between single,
double, and triple bonds Bonds between elements become shorter and
stronger as multiplicity increases. The greater the bond energy, the shorter the
bond length
Bond Bond Length Bond energy type order pm kJ/mol
C C 1 154 347C C 2 134 615C C 3 120 812
BOND LENGTH AND ENERGYBond Length
(pm)Energy
(kJ/mol)
C - C 154 346
C=C 134 612
CC 120 835
C - N 147 305
C=N 132 615
CN 116 887
C - O 143 358
C=O 120 799
CO 113 1072
N - N 145 180
N=N 125 418
NN 110 942
IONIC COMPOUNDS & LEWIS STRUCTURES
Na + Cl Na + Cl
The electron from Na is given to the Cl. Now both satisfy the octet rule.
Below, two electrons are given to S, one from each K.
+
VSEPR model
Most important factor in determining geometry is relative repulsion between electron pairs.
Electrons arranged so that pairs are as far apart from each other as possible.
Occurs in 3-dimensional space
Molecule adopts the shape that minimizes the electron
pair repulsions.
Valence Shell Electron Pair Repulsion theory
MOLECULAR GEOMETRY
Molecules have specific shapes.
Determined by the number of electron pairs around the central species
• Bonded and unshared pairs count (unshared pairs take up slightly more space)
• Multiple bonds are treated as a single bond for geometry.
• Geometry affects factors like polarity and solubility.
1 atom bonded to another atom
Electron Domains
Basic Geometr
y
0 lone pair
1 lone pair
2 lone pairs
3 lone pairs
4 lone pairs
1 Linear180°
1 ELECTRON DOMAIN
2 atoms, or lone electron pairs, or a combination of the two, bonded to a central atom.
Electron Domains
Basic Geometry
0 lone pair 1 lone pair 2 lone pairs 3 lone pairs
2 Linear180°
Linear180°
2 ELECTRON DOMAINS
3 atoms, or lone electron pairs, or a combination of the two, bonded to a central atom.
Electron Domains
Basic Geometry
0 lone pair 1 lone pair 2 lone pairs 3 lone pairs
3trigonal planar120°
bent / angular<120°
3 ELECTRON DOMAINS
4 ELECTRON DOMAINS 4 atoms, or lone electron pairs, or a
combination of the two, bonded to a central atom.
Electron Domains
Basic Geometry
0 lone pair 1 lone pair 2 lone pairs
4Tetrahedral
109.5°Pyramidal<109.5°
bent / angular<109.5°
5 ELECTRON DOMAINS 5 atoms, or lone electron pairs, or a
combination of the two, bonded to a central atom.
Electron Domains
Basic Geometry
0 lone pair 1 lone pair 2 lone pairs 3 lone pairs
5trigonal
bipyramidal180°, 120°,
90°
seesaw/ sawhorse t-shape
linear
6 ELECTRON DOMAINS 6 atoms, or lone electron pairs, or a
combination of the two, bonded to a central atom.
Electron Domains
Basic Geometry
0 lone pair 1 lone pair 2 lone pairs 3 lone pairs
6Octahedral90°, 180°
square pyramid
square planar
INTERMOLECULAR FORCES
Forces of attraction between molecules Forces within molecules are intramolecular
Why would boiling point be a good indicator of intermolecular force strength?
HYDROGEN BONDING H bonded to N, O, F
Example: Water Strongest intermolecular force
Large electronegativity difference Size of H atom allows it to get close to unshared
pair of electrons in adjacent molecule
POLARITY
A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.
H F+ -
Slight positive sideSmaller electronegativity
Slight negativeLarger electronegativity
DIPOLES
Dipole – created by equal but opposite charges that are separated by a short distance Molecules with dipoles are polar because of
uneven charge distribution Direction is from its positive to negative pole Molecules can have multiple dipoles Dipoles can cancel each other out if in equal and
opposite directions
ION-DIPOLE FORCE
Exists between ion and the partial charge on the end of a polar molecule Cations attracted to negative end of
dipole Anions attracted to positive end of
dipole Magnitude of attraction increases
as charge of ion increases as magnitude of dipole increases
Stronger than dipole-dipole forces
DIPOLE-DIPOLE FORCES
Not as strong as ion-dipole forces Strength of attraction increases with
increasing polarity
Example: HCl
LONDON DISPERSION FORCES
The constant motion of electrons sometimes creates temporary dipoles for an instant Momentary uneven charge creates a dipole Can induce a dipole in another molecule
LONDON DISPERSION FORCES
Weakest intermolecular force Sometimes called dipole induced dipole
All molecules have LDF Dependent on motion of electrons so more
electrons means more chances for LDF LDF strength increases
with increasing molar mass
Example: O2
VALENCE-BOND THEORY Covalent bonding involves sharing of
electrons Electrons exist in orbitals
Orbitals overlap so that electrons can form pairs to make a bond
As orbitals overlap, they mix and form new hybrid orbitals Mixing of the orbitals is called hybridization
HYBRIDIZATION = the blending of orbitals
Poodle
+
+ Labrador
=
=
=
=
+
+s orbital p orbital
Labradoodle
sp orbital
EXAMPLE: CH4
Orbital notation for carbon:
Carbon only has 2 electrons available to bond, but it has to make 4 bonds
Carbon promotes one of its electrons to the 2p orbital so that each electron can pair up with the 1s electron from each of the four carbons
1s orbitals of four hydrogen atoms
Promoted electrons in carbon allow 4 bonds
EXAMPLE: CH4
The overlap of carbon’s 1 electron in the s orbital and 3 electrons in the p orbital creates four hybrid orbitals This hybridization is called sp3
The new hybrid orbitals have more energy than an s orbital but less than a p orbital
HYBRID ORBITALS
To determine hybridization Draw Lewis structure Determine electron domains for target atom Hybrid orbitals correlate to number of domains
Orbital Name Orbitals Combined Electron domains
sp 1 s / 1 p orbital 2
sp2 1 s / 2 p orbitals 3
sp3 1 s / 3 p orbitals 4
sp3d 1 s / 3 p / 1 d orbital
5
sp3d2 1 s / 3 p / 2 d orbitals
6
SIGMA AND PI BONDS
Sigma () bonds exist in the region directly between two bonded atoms.
Exist on the line (internuclear axis) between two atoms
Single bonds
Sigma bonds
s s
s p
p p