Chemical Bonding and Lewis Structures
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Transcript of Chemical Bonding and Lewis Structures
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Chemical Bonding and Lewis StructuresChemical
Bonding and Lewis Structures
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Chemical Bonding Chemical Bonds are the forces that hold
atoms together. Atoms form bonds in order to attain a minimal
energy state. Bond formation is an exothermic process (just
as bond breaking is endothermic) The type and strength of bond that forms
between reacting particles dictates the physical and chemical characteristics of the molecule or ion in question.
Bond energy is the energy needed to break a bond and is an indication of bond strength.
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The Relationship Between Electronegativity and Bond Type
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Ionic Bonds
Ionic bonds occur between ions due to electrostatic attraction between positive cations and negative anions.
Form between atoms with large differences in electronegativity (>1.7), usually between a metal and a nonmetal.
Relative strength can be determined using Coulomb’s Law E = k(q1)(q2)
r2
K= coulomb’s constant; q1=charge on one ion; q2= charge on other ion; r=distance between ions
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Ionic Bond, continued
The strength of an ionic bond is directly proportional to the magnitude of the charges involved and inversely proportional to the square of the distance between them.
Ionic solids tend to have high melting points
Often soluble in polar solvents, such as water.
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Ionic Solid
UNIT CELL Crystal lattice – at the lattice sites in the unit cell
are positive and negative ions held together by Coulombic attraction force between positive and negative ions.
CONDUCTIVITY Solid has no conductivity due to zero empty
valence orbitals. The ions in the crystal are isoelectronic with Group 18.
Molten and aqueous solutions will conduct due to mobile ions.
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Covalent Bonds
Nonpolar covalent bonds form between atoms of nonmetals with nearly identical electronegativities while polar covalent bonds for between nonmetals with dissimilar electronegativities (0.4-1.7).
Covalent bonds within molecules are strong, but the binding forces between molecules are relatively weak.
Molecular solids usually have low melting points.
Usually soluble in nonpolar solvents (carbon tetrachloride)
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Which of the following bonds would be the least polar yet still be considered polar covalent?
Mg-O C-O O-O Si-O N-O
React 5
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Metallic Bonding
Occurs in metallic solids Metal atoms usually have large, positively
charged nuclei and few valence electrons. Nuclei are positioned in a regular geometric
array (lattice) by electrostatic repulsion. Valence electrons are attracted equally by all
nuclei. Leads to the “sea of electrons” model with
nuclei bobbing in a “sea of electrons” Useful for explaining the physical characteristics of
metals (ie, conductivity) Metals have a wide range of melting points.
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Metallic Solid
CRYSTAL LATTICEAt lattice sites in unit cell are positive ions
held together by mobile valence electrons traveling through empty valence orbitals.
POSITIVE CHARGE DENSITYSmaller ionic radius = higher melting point
CONDUCTIVITYSolid: excellent due to empty valence
orbitalsLiquid: good due to mobile ions.
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Intermolecular Forces
The group of weaker attractive forces between atoms or molecules. NOT BONDS!!!!!
Van der Waals forces London dispersion forces-all atoms and molecules; caused by
temporary dipoles created as electrons move about the nucleus. Strength depends on the number of electrons moving, so molecules
with larger masses (more electrons) have greater London forces. Dipole-dipole force- an attraction between opposite polar ends
of adjacent molecules. Hydrogen bonding- occurs when hydrogen bonds with a very
electronegative anoin (F, O, N) resulting in a very polar molecule. Strongly positive and negative ends of the molecule have stronger
interactions than either london or dipole-dipole forces.
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The Effect of an Electric Field on Hydrogen Fluoride Molecules
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Lewis Structure
Lewis Structures – shows how the valence electrons are arranged among the atoms of a molecule
There are rules for Lewis Structures that are based on the formation of a stable compound
Atoms want to achieve a noble gas configuration
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Octet & Duet Rules
Octet Rule – atoms want to have 8 valence electrons
Duet Rule – H is the exception. It wants to be like He & is stable with only 2 valence electrons
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Steps for drawing Lewis Structures
1. Sketch a simple structure with a central atom and all attached atoms
2. Add up all of the valence electrons for each individual atom
If you are drawing a Lewis structure for a negative ion add that many electrons to create the charge
If you are drawing a Lewis structure for a positive ion subtract that many electrons to create the charge
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1. Subtract 2 electrons for each bond drawn2. Complete the octet on the central atom &
subtract those electrons3. Complete the octet on the surrounding atoms
& subtract those electrons4. Get your final number
If 0 you are done! If + add that many electrons to the central atom If - need to form multiple bonds to take away
that many electrons
Steps for drawing Lewis Structures
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Examples
CCl4 Sketch a simple structure with a central
atom and all attached atoms Cl │Cl – C – Cl │ Cl
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Examples
Add up all of the valence electrons for each individual atom
4 + 4(7) = 32 Subtract 2 electrons for each bond drawn 32-8 = 24 Complete the octet on the central atom &
subtract those electrons Done
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Examples
Complete the octet on the surrounding atoms & subtract those electrons
24 – 24 = 0 Final number = 0…DONE! Final structure is… __ │Cl │ __ │ __│Cl – C – Cl │ │ │Cl │
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Examples
HF
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Examples
NH3
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Examples
NO+
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Exceptions to the octet rule
Sometimes the central atom violates the octet rule and has more or less than 8 valence electrons
Keep using the same rules to draw Lewis Structures
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Examples
SF4
ICl3
XeF4
ICl4
-
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Resonance
When more than one Lewis Structure can be written for a particular molecule
Resonance structure – all possible Lewis structures that could be formed
The actual structure is the average of all of the structures
You MUST show all structures!
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Examples
SO3
NO2-
NO3-
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Covalent Network Crystal
CRYSTAL LATTICEAt the lattice site in the unit cell there are
atoms held together by strong covalent bonds.
Si, SiO2, C (diamond), and C (graphite)This crystal has the highest melting point.
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SUMMARY
Covalent network solids have the highest melting points because of the strongest forces holding the crystal together (covalent bonds).
Metallic solids generally have the next highest melting points, but the larger the ionic radius, the lower the melting point.
Ionic crystals next. Molecular crystals last.
Consider IMF to rank these. The stronger the IMF, the higher the melting point.