chem 110 mcgill sos
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Chem 110 McGill SOS. Midterm 2. Topics on Midterm II. Chapter 10+11+12 Remember, only 15% of the questions are going to be from midterm 1! Don’t overdo it!. Chapter 10 – Bonding I. Covalent bonding Polarity & dipole moment Lewis structures Formal charge Resonance VSEPR theory - PowerPoint PPT PresentationTRANSCRIPT
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Chem 110McGill SOS
• Midterm 2
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Topics on Midterm II Chapter 10+11+12
Remember, only 15% of the questions are going to be from midterm 1! Don’t overdo it!
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Chapter 10 – Bonding I
Covalent bonding Polarity & dipole moment Lewis structures Formal charge Resonance VSEPR theory Bond Energies
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3 types of bonding
• Covalent = sharing of electrons
• Ionic = transfer of electrons
• Metallic = delocalization of electrons
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Covalent bonding
• pure covalent = same element, equal sharing
• polar covalent = different elements (non-metals), unequal sharing
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Chapter 10 – Bonding I
Covalent bonding Polarity & dipole moment Lewis structures Formal charge Resonance VSEPR theory Bond Energies
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Polar covalent = think Dipole moment
Arrow always points from the partial positive charge (δ+) to partial negative charge (δ-)
electronegativity differences determine the degree of polarity
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Electronegativity a measure of the electron-
attracting power of a bonded atom type of bond determined by
difference in EN remember, 1.7 is arbitrary!
Ionic -greater than 1.7 Polar Covalent - between 0.5-1.7 Pure Covalent - less than 0.5
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Chapter 10 – Bonding I
Covalent bonding Polarity & dipole moment Lewis structures Formal charge Resonance VSEPR theory Bond Energies
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Lewis theory Valence electrons are most
important in chemical bonding Why do atoms bond? To acquire
a stable octet. Remember, eight is yummy!
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Lewis theory -terminology
Octet Rule: Each atom must have 8 valence shell electrons in a lewis structure
Bonding pair vs. lone pair Single, double and triple covalent
bonds
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Lewis Symbols
Chemical symbol in the middle
Dots represent valence electrons
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Lewis Structures - Ionic Compounds
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Lewis Structures - Covalent Compounds
Strategy 1. Count total # valence electrons Add/subtract additional charges
+ve charge, remove e- -ve charge, add e-
2. Identify the central and the terminal atoms (any carbon or hydrogen?)
3. Join atoms through SINGLE covalent bonds. This is your skeletal structure
4. Achieve octet around each atom For remaining valence e-, add to central
atom. Add double and triple bonds when necessary to complete an octet.
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Lewis Structures - Tips
Hydrogen is always a terminal atom Carbon is always a central atom Central atom (if not carbon) will be
the least electronegative element Go for a compact, symmetrical
structure
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Example: HCN
Step 1: Count total # valence e-
H=1, C=4, N=5
So…1+4+5=10 e-
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Example: HCN Step 2: Identify the central and
the terminal atom (any carbon or hydrogen?)
Carbon=central atom Hydrogen= terminal atom therefore Nitrogen= also
terminal atom
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Example: HCN Step 3: Join atoms through SINGLE
covalent bonds. This is your skeletal structure.
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Example: HCN Step 4: Complete the octet of terminal
atoms. Count the number of valence
electrons remaining. If you have any remaining, use them on the central atom. If your central atom does not have a complete octet, start switching single covalent bonds to double (or triple if necessary)
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Chapter 10 – Bonding I
Covalent bonding Polarity & dipole moment Lewis structures Formal charge Resonance VSEPR theory Bond Energies
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Formal charges
• helps us draw Lewis structures
• tells us where the e- are located
• F.C. = # valence e- - # lone pair e- - 1/2 # bonding e-
• NOTE: sum of all F.C in molecule must be equal to overall charge!
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Formal Charge - Question Types
2 types of questions: 1. what is the formal charge on
each atom?Use FC equation
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Example: CHO2(-) Draw structure out, if necessary Use F.C. = # valence e- - # lone pair e-
- 1/2 # bonding e-
F.C. (H) =1-0-0.5(2)=1-0-1=0
F.C. (C) =4-0-0.5(8)=4-0-4=0
F.C. (O)-SB =6-6-0.5(2)=6-6-1=-1 F.C. (O)-DB =6-4-0.5(4)=6-4-2=0 overall F.C. = -1 = charge on
molecule
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Formal Charge - Question types
II. Use formal charges to see if the lewis structure is plausible or not?
Rules: Sum of the formal charges must equal 0
exceptions: the molecule is an ion (like in the last example)
Keep structures as small as possible Negative formal charges are usually on the
most electronegative atoms Positive formal charges are usually on the
least electronegative atoms
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Example: NH2CN NH2CN has two possible Lewis
structures How do we know which one is
more plausible?
Use formal charges!
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Chapter 10 – Bonding I
Covalent bonding Polarity & dipole moment Lewis structures Formal charge Resonance VSEPR theory Bond Energies
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Resonance forms Two or more plausible lewis structures
can be written for the same molecule. Note: They will have the same
skeletal structure. It’s only the electron distribution that changes!
Each lewis structure contributes to the resonance hybrid.
increasing resonance = more stable
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Exceptions to the Octet Rule
Odd-electron species total # of valence electrons is odd use formal charges to determine where to
place the single electron Incomplete octets in order to avoid breaking a formal charge rule
(i.e. negative formal charges are usually on the most electronegative atoms), you break the octet rule
usually only seen with compounds containing Be, B, Al
Expanded valence shells seen with elements containing d-obitals like S,
P + bonded to a highly electronegative atom like O,F,Cl
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Chapter 10 – Bonding I
Covalent bonding Polarity & dipole moment Lewis structures Formal charge Resonance VSEPR theory Bond Energies
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VSEPR Theory concerns shapes of molecules Electron pairs (either lone pair or
bond pairs) assume an orientation about the atom that will minimize repulsionstherefore we want to maximize the
distance between electron pairs order of repulsion interactions:
Lone pair-lone pair = most repulsive lone pair-bond pair bond pair-bond pair = least repulsive
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VSEPR Theory - Tips
Molecular shape depends on: number of electron pairs type of electron pair (lone pair vs
bonding pair) When determining the molecular
geometry, you must first determine the electron group geometry!
Treat multiple bonds (double, triple) as if it were a single bond
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KNOW THIS: Molecular Geometries
2 electron groups▪ 0 lone pairs=linear
3 electron groups▪ 0 lone pairs= trigonal-planar▪ 1 lone pair= bent
4 electron groups▪ 0 lone pairs= tetrahedral▪ 1 lone pair= trigonal-pyramidal▪ 2 lone pairs=bent
5 electron groups▪ 0 lone pairs=trigonal-bipyramidal▪ 1 lone pair=seesaw▪ 2 lone pairs=T-shaped▪ 3 lone pairs=linear
6 electron groups: ▪ 0 lone pair=octahedral▪ 1 lone pair=square-pyramidal▪ 2 lone pair=square-planar
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VSEPR Theory Step 1: Count the # of electron
groups (the number of bonded atoms+the number of lone pairs)
Step 2: Determine electron group geometry 2 electron groups: linear 3 electron groups: trigonal-planar 4 electron groups: tetrahedral 5 electron groups: trigonal-
bipyramidal 6 electron groups: octahedral
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VSEPR Theory Step 3: Count the number of lone
pairs. If there are no lone pairs, the electron group geometry IS the molecular geometry.
Step 4: Determine molecular group geometry. each additional lone pair
changes the molecular geometry
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Example: Sulfur Tetrafluride
5 electron groups Electron geometry is
trigonal bipyramidal However, because 1
of our electron groups is a lone pair, the molecular geometry will be seesaw
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Chapter 10 – Bonding I
Covalent bonding Polarity & dipole moment Lewis structures Formal charge Resonance VSEPR theory Bond Energies
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Bond engergies, order, and bond
length Bond energy is the energy required to break the bond between atoms single covalent bond= weakest Triple covalent bond = strongest
Bond order Single covalent bond=1 Double covalent bond=2 Triple covalent bond=3
Bond length is the distance between the centers of the two atoms joined by a covalent bond
NOTE: As bond order increases, bond energy increases and bond length decreases.
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Chapter 11- Bonding II
Valence Bond Theory Hybridization M.O. Theory Metallic Bonding
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Valence Bond Theory
In a molecule, electron probability is highest where the orbitals of the atoms overlap to form covalent bondsoverlap allows electrons greater
freedom across both orbitals thus greater orbital overlap =
more stable (think about single vs double vs triple bonds)
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Valence Bond Theory - types of
hybrid bonds σ bond: most of e- density around bond axis formed by overlap of
s-s, s-p, or p-p head-to-head overlap
π bond: e- density above and below the bond axis formed by overlap of
p-p side-to-side overlap
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Valence Bond Theory - multiple
bonds Single covalent bond= 1 sigma bond
Double covalent bond=1 sigma bond (σ)+ 1 pi bond (π)
Triple covalent bond=1 sigma bond (σ) + 2 pi bonds (π)
NOTE: add π bonds only after σ bonds! π bonds only for double or triple bonds
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Practice QuestionThe resonance structure of PO4(3-) where the
central atom carries a formal charge of zero issurrounded by:
A. 4 sigma bonds B. 4 sigma and 1 pi bonds C. 4 sigma and 2 pi bonds D. 4 sigma and 3 pi bonds E. 4 sigma and 4 pi bonds
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Practice question - steps
• draw Lewis structure
• remember to complete octets and use F.C. to determine best structure
• count single and multiple bonds
• single bond = 1 sigma
• multiple bonds = 1 sigma + additional pi
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Chapter 11- Bonding II
Valence Bond Theory Hybridization M.O. Theory Metallic Bonding
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Hybridization Bonded atoms do not have the same
orbitals as non-bonded, isolated atoms. They have hybrid orbitals.VSEPR -> geometryValence bond theory -> bonding
Number of hybrid orbitals= Number of electron groups
NOTE: lone pairs are also hybridized! Double and triple bonds are not!
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Hybrid OrbitalsNumber of Electron Groups
Electron Group Geometry
Molecular Group Geometry
Hybrid Orbitals
2 Linear Linear sp+2p
3 Trigonal planar Trigonal PlanarBent
sp^2+p
4 Tetrahedral TetrahedralTrigonal pyramidalBent
sp^3
5 Trigonal-Bipyramidal
Trigonal BipyramidalSeasawT-shapedLinear
sp^3d
6 Octahedral OctahedralSquare PyramidalSquare Planar
sp^3d^2
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Hybridization Question: What molecular shapes
can a sp^3d^2 hybridized molecule have?
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Hybridization Answer: Octahedral, Square
Pyramidal and Square Planar
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Hybridization - tips You need the electron group
geometry to determine the hybrid orbitals
to determine the electron group geometry, you need to draw a lewis structure
Lewis structure-> VSEPR->Hybridization
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Example CO2
Steps: Draw lewis structure Determine electron group geometry think: hybridization using chart you
will have memorized!
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Chapter 11- Bonding II
Valence Bond Theory Hybridization M.O. Theory Metallic Bonding
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Molecular Orbital Theory
Compensates for Valence Bond Theory in explaining magnetism
Used for diatomic molecules Excellent for predicting
magnetic properties (diamagnetic and paramagnetic)
M.O. theory -> think: molecular orbitals
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Molecular orbitals
Two types of molecular orbitals arise from the combination of atomic orbitals: Bonding Molecular Orbitals Antibonding Molecular Orbitals
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• Bonding M.O = increase e- density
• maximum e- density between 2 nuclei
• Anti-bonding M.O = decrease e- density
• no e- density between 2 nuclei
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Key principles of the M.O. Theory
# of atomic orbitals combined = # of molecular orbitals formed
Antibonding M.O have higher energy levels than bonding M.O
In ground state electron configuration, electrons enter lowest energy M.O available Use Hund’s Rule + Pauli’s exclusion
principle: They fill up M.O orbitals singly before pairing up
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Stabilization energy• differs from molecule to molecule
• presence bonding e- = stabilize molecule
• presence antibonding e- = destabilize molecule
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Bond Order Bond Order= 1/2 (# bonding e- -
# antiboding e- ) NOTE: bond order always > 0! For first period elements
(hydrogen and helium), we only use σ1s M.O and σ*1s M.O
For second period elements (lithium, beryllium…), the M.O. listed above are already filled and we concern ourselves with the following M.O.s: σ2p, σ*2p, π2p(x2), π*2p(x2)
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Chapter 11- Bonding II
Valence Bond Theory Hybridization M.O. Theory Metallic Bonding
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Metallic Bonding - Electron Sea Model
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Properties of metals
• luster of reflectivity
• high electrical conductivity
• high thermal conductivity
• malleability & ductility
• electronic emission
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Band Theory In conductors, the valence
band and the conduction band overlap one another. No energy gap!
In semiconductors, the valence band and the conduction band are seperated by a small energy gap
In insulators, the valence band and the conduction band are seperated by a large energy gap!
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Semiconductors Intrinsic semiconductors have a
energy gap with a FIXED size
Extrinsic semiconductors have a energy gap whose size can be controlled via doping
Doping =adding of impurities. It increases conductivity.
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Semiconductors
Two types of extrinsic semiconductors: N-type▪ Migration of electrons from donor atoms (i.e.
phosphorus)▪ absorbs blue light ▪ appears yellow
P-type▪ Migration of positive holes from acceptor atoms (i.e.
boron)▪ absorbs red and green light ▪ appears blue
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Chapter 12 – Intermolecular Forces
London dispersion forces Dipole-dipole Interactions Hydrogen Bonding Ionic and Network Covalent Solids
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Intermolecular forces
forces between molecules intermolecular forces keep molecules close together
explains physical properties of solids, liquids, and gases
THINK: attraction between molecules
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Intermolecular forces - physical
properties• when comparing substances of similar nature (elements in same group):
• melting point
• boiling point
• increases with increasing molecular size (weight)
• also affected by molecular shape
• THINK: hydrocarbons
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Chapter 12 – Intermolecular Forces
London dispersion forces Dipole-dipole Interactions Hydrogen Bonding Ionic and Network Covalent Solids
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London dispersion forces
dominant force in nonpolar molecules due to small, instantaneous dipole moments
THINK: electrons are always in motion London forces are the weakest of all
intermolecular interactions present in any molecule increases with molecular size
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Example Which of the following has a higher boiling
point?
n-pentane
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Chapter 12 – Intermolecular Forces
London dispersion forces Dipole-dipole Interactions Hydrogen Bonding Ionic and Network Covalent Solids
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Dipole-dipole Interactions
Dipole-dipole forces occur in polar molecules (unequal sharing of electrons) due to dipole moments
For example, HCl is polarchlorine is partially (-), and hydrogen is
partially (+)
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!!!
• NOTE: dipole-dipole interactions do not replace dispersion forces
• they occur in addition to them!
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Example 1Which of the following would
have the higher boiling point?a)CH3CNb)CH3CH2CH3
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Example 2
• Predict the trend in boiling points of HCl, HBr and HI
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Logical reasoning...?
• determine polarity by comparing electronegativity of Cl, Br, and I
• therefore order of increasing b.p: HI < HBr < HCl
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Trick question!!
• for large (heavy M.W) molecules, london dispersion forces are stronger than dipole-dipole forces
• therefore we determine increasing b.p by looking at M.W and not polarity!
• correct answer: HCl < HBr < HI
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Chapter 12 – Intermolecular Forces
London dispersion forces Dipole-dipole Interactions Hydrogen Bonding Ionic and Network Covalent Solids
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Hydrogen Bonding
occurs with small molecules with Hydrogen bonded to N, O, or F
have much higher melting and boiling points than expected from M.W
why? Hydrogen is attracted to the lone pairs on O, N, and F
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Chapter 12 – Intermolecular Forces
London dispersion forces Dipole-dipole Interactions Hydrogen Bonding Ionic and Network Covalent Solids
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Network covalent solids
• THINK: diamonds are forever
• each molecule is held together by strong, covalent bonds
• strong directionality = strong material
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Ionic solids
• made up of oppositely charged ions that interact by Coulombic forces between them
• think: salt! Made up of sodium and chloride ions
• crystalline ionic solids: highly organized fashion of ions
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Terminology• Unit cell -> lattice -> crystal
• most ionic solids are formed of cubic unit cells
• 6 faces
• 8 corners
• 12 edges
• 3 types of cubic unit cells:
• simple cubic
• body centered
• face centered
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Simple cubic unit cell• ions of the same type present at
corners of the cube
• how many ions present per unit cell?
• each corner is shared with 3 other corners (1 -> 1/2 -> 1/4 -> 1/8)
• therefore ion shared in corner = 1/8 inside each unit cell
• total in simple cubic unit cell:
• 8 corners x 1/8 ion = 1 ion per unit cell
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Body centered cubic unit cell
• ions of the same type present inside the body and at the corners of the cube
• how many ions present per unit cell?
• for corners: 8 corners x 1/8 ion = 1 ion per unit cell
• PLUS 1 ion inside not shared
• total = 2 ions / body centered cubic unit cell
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Face centered cubic unit cell
• ions of the same type present at corners and faces of the cube
• how many ions present per unit cell?
• corners: 8 corners x 1/8 ion = 1 ion
• faces: each face is shared with 1 other face
• 6 faces x 1/2 ion = 3 ions
• total = 1 ion for corners + 3 ions for faces = 4 ions per face centered cubic unit cell
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Example: NaCl
• what type of unit cell is a crystalline ionic solid of NaCl?
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NaCl - steps
• look at a single unit cell
• find ions of the same type
• simple, body-centered, or face-centered?
• larger ion occupies the corners, smaller ion comes in to fill the spaces
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Coordination number
• How many ions surround those of another type?
• in NaCl, each Cl- is surrounded by 6 Na- ions, and vice versa
• therefore coordianation number = 6
• ratio of cation : anion = 1:1
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