Chapters 9 & 10: Chemical Bonding

Chapters 9 & 10: Chemical Bonding

Chemical Bonds Review:
Atoms naturally form chemical bonds by linking with other atoms to minimize energy Most chemical properties are determined by an atom’s valence electrons N-3 1s2 2s2 2p6 Al+3 1s2 2s2 2p6 3s0 3p0 Atoms seek to obtain a configuration like that of a Noble Gas (octet) through bonding. Electrons are attracted to the nucleus of other atoms by electrostatic forces.

Types of Bonds Ionic Bond (metals & non-metals)
e-’s are transferred from atom to atom Held by electrostatic attraction Covalent Bond (b/w non-metal atoms) Valence electrons are shared Metallic Bond (between metal atoms) electrons belong to no specific atom e-’s flow freely through the substance TedEd: How Atom’s Bond

: . : F : Electron Dot Notation
Dots that represent the valence electrons are placed around the element symbol Maximum of 8 Except H and He only has 2 Must obey Hund’s Rule Spread out before pairing Not used for transition metals Fluorine has 7 valence e- F 1s2 2s2 2p5 : . : F :

Lewis Dot Symbols for the Representative Elements & Noble Gases
5

The Ionic Bond Ionic bond: the electrostatic force that holds ions together in an ionic compound. Li+ F - Li + F 1s2 1s22s22p6 1s22s1 1s22s22p5 [He] [Ne] Li Li+ + e- + e- + F - = F - Li+ + Li+ LiF

Electrostatic (Lattice) Energy
Lattice energy (U) is the energy needed to completely separate a solid ionic compound into gaseous ions. Q+ is the charge on the cation E = k Q+Q- r Q- is the charge on the anion r is the distance between the ions E is the potential energy “Coulomb’s law” Compound Lattice Energy (kJ/mol) NaCl 788 Q: +1,-1 Higher ~ more stable MgF2 MgO 2957 Q: +2,-1 Q: +2,-2 Lattice energy increases as charges increase or ion size decreases. 3938 LiF LiCl 1036 853 r F- < r Cl-

The solubility is directly related to Lattice energy because dissolving ionic compounds involves breaking the ions apart Larger charge, Stronger bond = Less Soluble

Lattice Energy Problems
Predict which of the following compounds will have the highest lattice energy? NaCl MgS MgCl2 Na2S +2 +1 -2 -1 With all period 3 atoms, the greatest ion charges hold together strongest. 2. Predict which of the following compounds will have the highest lattice energy? KBr NaCl LiF RbI +1 -1 Charges are all equal, strongest ionic bond comes from smallest ions (Li and F atoms are in lowest period and thus smaller atoms)

Properties of Ionic Compounds
Generally water soluble (polar dissolves polar) Dense, brittle, and hard solids High melting points Poor conductors in solid form, but good conductors when dissolved in water or molten

Metallic compounds Not limited to octet rule (d orbitals)
Transition metals relatively stable when neutral Electron sea theory – low ionization energies, e- can easily move from atom to atom Explains the hardness of metals – difficult to separate Held together by delocalized electrons Act like “glue”

Metallic Properties Metallic Properties– typically same as pure metals
Conductive – because electrons can easily travel from one metal atom to the next. Luster (shine) Malleable Substitution alloy (e.g. bronze: Cu & Sn) Interstitial alloy (e.g. Steel: Fe & C)

The Covalent Bond A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. 7e- 7e- 8e- 8e- F F + F Shared electrons are counted for both to fulfill each atom’s octet A pair of shared electrons is represented as a line between atoms lone pairs lone pairs lone pairs lone pairs F single covalent bond (2 e-) Valence electrons not involved in bonding are called Lone Pair electrons Lewis structure of F2

Crash Course: Bonding Models and Lewis Structures:
Lewis structure of water single covalent bonds 2e- 8e- 2e- H + O + H O H or Double bond – two atoms share two pairs of electrons 8e- O C or O C double bonds Triple bond – two atoms share three pairs of electrons 8e- N 8e- or N triple bond Crash Course: Bonding Models and Lewis Structures: 15

Properties of Molecules
Low melting points (dependent on size) Poor conductors of heat and electricity Fairly easy to separate in mixtures Exceptions to properties Network Covalent Substances 3D crystal (similar to ionic bond) Diamond

Isomers possess the same chemical formula,
Isomers possess the same chemical formula, but a different arrangement of atoms. Distinct chemicals with different properties C3H6O has 11 possible isomers Acetone Allyl Alcohol Methyl vinyl ether Propanal Cyclopropanol Oxetane

Length/Strength of Covalent Bonds
(1 pm = m) 198 pm : I2 945 kJ/mol Very inert 436 kJ/mol 298 kJ/mol 151 kJ/mol Bond Length a) Bond Length ↑ as Atomic radius ↑ b) Single bond > Double Bond > Triple Bond Bond Strength a) As bonds get shorter, bond strength ↑. C-C > Si–Si b) As bond order ↑, bond strength ↑. C=C > C–C

Carbon forms stable bonds
The CAS (Chemical Abstracts Service) Registry contains more than 100 million unique chemical substances on file. Over 80% of them are Organic substances (Carbon compounds). Small compounds (Acetone, ibuprofen, octane, Vitamin A, ATP) Biomolecules (Proteins, DNA, Carbohydrates, Lipids) → oil/gas Plastics (Rubber, PVC, Teflon, Vinyl, Propylene, Styrofoam) Why so many? Carbon is versatile 4 x single bonds 1 x double and 2 single bonds 2 x double bonds 1 x triple and 1 single bond Carbon forms stable bonds Smaller atomic size than Silicon Shorter bonds = Stronger bonds Can bond with a variety of atoms

Covalent Bonding Atoms will normally have a number of bonds equal to the number of valence electrons needed. Bonds Formed 4A: 4 x C single bonds 1 x C double and 2 single bonds 1 x C triple and 1 single bond 2 x double bonds 5A: 3 x N single bonds 1 x N double and 1 single bond 1 x N triple bond 6A: 2 x O single bonds; 1 x O double bond O C : : 7A: 1 x F single bond :

Drawing Lewis Structures
Place element with largest number of unpaired electrons in the center Place other elements around the outside Add Lewis electron dots Match up the unpaired electrons Create Single Bonds If necessary, create double or triple bonds so every atom has 8 valence e- Hydrogen will only have 2 NH3 H N H H

Lewis Structure Practice
Br2 CO2 H2CO O C O H C O H Extra help: Bozeman Science: Lewis Diagrams

Problem: Lewis Structure
Draw two plausible Lewis Structures based on the following arrangement of C3H6O. Note: Carbon always has 4 bonds Oxygen 2 bonds, 2 lone pairs Hydrogen 1 bond

Draw a possible Lewis structure:
many isomers can exist for some CS2 Si2Cl6 C4H9NO N3OF3 C3H4 Benzene (C6H6; forms a ring)

Lewis Structure Practice
Bell Ringer Lewis Structure Practice Draw two isomers of C4H8N2S

Formal charges in Polyatomic ions: Possess both covalent bonding and a net charge
NH4+ CO3-2 OH- NO2-1 We add a dot (electron) for every - and remove for + -1 O H O –H × + + Oxygen can have only 1 bond (instead of 2) with a -1 charge Nitrogen can have 4 bonds with a +1 charge -1 -1 O N O

Two possible skeletal structures of formaldehyde (CH2O)
Which is correct? H C O H C O An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. 1 2 total number of bonding electrons ( ) Formal charge on an atom = Valence e- in free atom - # nonbonding electrons - The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. Method 2: Quicker way to find formal charge: 1. Note how many (#) electrons an atom needs to gain an octet 2. If it has the same # bonds as the #, it’s 0 If it has 1 less bond than the #, it’s -1 ; (2 less = -2) If it has 1 more bond than the #, it’s +1 ; (2 more = +2)

For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. Lewis structures with large formal charges are less plausible than those with small formal charges. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. (b) is more likely structure Because no formal charge (b) is most likely structure for N2O

( ) - - Potential CH2O structure #1 -1 +1 H C O 1 2 Formal charge =
valence electrons nonbonding electrons - - bonding electrons Formal charge = Method 2 C needs 4 e- from 4 bonds, has one less than needed so: -1 Formal charge on C = 4 - 2 - ½ x 6 = -1 O needs 2 e- from 2 bonds, has one more than needed so: +1 Formal charge on O = 6 - 2 - ½ x 6 = +1

( ) - - Potential CH2O structure #2
H C O ( ) 1 2 valence electrons nonbonding electrons - - bonding electrons Formal charge = Method 2 C needs 4 e- and has 4 bonds: 0 Formal charge on C = 4 - 0 - ½ x 8 = 0 Formal charge on O = 6 - 4 - ½ x 4 = 0 O needs 2 e- and has 2 bonds: 0 Based on the two possible structures, this one is favored because it has smaller formal charges

Lewis Structure with Formal charges
Bell Ringer Lewis Structure with Formal charges Draw Aspartic acid C4H6NO4- H N C C + - C - C -

A resonance structure is one of multiple structures for a single molecule that cannot be represented accurately by only one Lewis structure. O + - O + - Ozone Inability to accurately represent bonding with Lewis Theory Unlike isomers, the atom connections are the same, but e- placement changes Does not switch back and forth, but is a hybrid between the two structures. Carbonate has 3 resonance structures: C—O bonds are 143 pm C=O bonds are 121 pm However, all bonds in CO3-2 are 131 pm The -2 charge is evenly distributed across the 3 O’s

Draw all resonance structures for the Nitrate ion (NO3-1)
– O +

Exceptions to the Octet Rule
Incomplete Octet – Not enough e- to gain 8 through sharing. F B Boron has only 3 e- allowing 3 covalent bonds for 6 total e- Odd-electron Molecules – pairing leaves one unpaired electron Free Radicals: Strong non-specific oxidizers (likely to steal electrons). N O S F Expanded Octet – Able to form more bonds than normal allowing for more than 8 shared electrons. Only found with Period 3 elements and lower. Utilize 3d orbitals for sharing more electrons

: : : Draw the Lewis structure with charges of: Perchlorate ClO4-1
Problem: Lewis Structure with Charges Draw the Lewis structure with charges of: Elements in 3p subshell (and below) can have expanded octets by utilizing available 3d orbitals. Generally, they can form as many bonds equal to their total number valence electrons : Perchlorate ClO4-1 Phosphate PO4-3 : : Sulfate SO4-2

Valence shell electron pair repulsion (VSEPR)
Predicts the geometry of a molecule from the electrostatic repulsions between ALL electrons, both bonding and nonbonding pairs. Molecule shapes arise to allow all electrons to spread out the furthest Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear B

Beryllium Chloride Carbon Dioxide C
0 lone pairs on central atom Cl Be 2 atoms bonded to central atom 0 lone pairs on central atom Carbon Dioxide O C 2 atoms bonded to central atom Double & triple bonds still count as 1 pair between atoms

Arrangement of electron pairs
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear trigonal planar trigonal planar AB3 3 Nitrate (NO3-1) Boron Trifluoride Ethene: 2 central atoms

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear AB3 3 trigonal planar AB4 4 tetrahedral tetrahedral *Important for organic and biochemistry

Methane Dashed line projects away; Propane (C3H8):
Phosphate (PO4-3) Dashed line projects away; Bold line projects toward viewer Propane (C3H8): 3 overlapping tetrahedrons

VSEPR AB2 2 linear linear AB3 3 trigonal planar AB4 4 tetrahedral
Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear AB3 3 trigonal planar AB4 4 tetrahedral tetrahedral trigonal bipyramidal trigonal bipyramidal AB5 5 Phosphorous Pentachloride (PCl5)

VSEPR AB2 2 linear linear AB3 3 trigonal planar AB4 4 tetrahedral
Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear AB3 3 trigonal planar AB4 4 tetrahedral tetrahedral AB5 5 trigonal bipyramidal AB6 6 octahedral octahedral Sulfur Hexafluoride (SF6)

Central atom has no lone pairs
Know These Three

lone-pair vs. lone-pair
bonding-pair vs. bonding-pair repulsion lone-pair vs. lone-pair repulsion lone-pair vs. < Lone-pair electrons take up more space than bonded electrons Lone pairs effect the molecule shape, but are not spatially taking up space

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal planar trigonal planar AB3 3 trigonal planar AB2E 2 1 bent

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB4 4 tetrahedral tetrahedral trigonal pyramidal AB3E 3 1 tetrahedral

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB4 4 tetrahedral tetrahedral AB3E 3 1 tetrahedral trigonal pyramidal AB2E2 2 2 tetrahedral bent Results in water’s high polarity, Making it a great solvent

VSEPR trigonal bipyramidal trigonal bipyramidal AB5 5 trigonal
Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal bipyramidal trigonal bipyramidal AB5 5 trigonal bipyramidal distorted tetrahedron AB4E 4 1 50

VSEPR trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1
Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1 trigonal bipyramidal distorted tetrahedron trigonal bipyramidal AB3E2 3 2 T-shaped 51

VSEPR trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1
Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1 trigonal bipyramidal distorted tetrahedron AB3E2 3 2 trigonal bipyramidal T-shaped trigonal bipyramidal AB2E3 2 3 linear 52

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB6 6 octahedral square pyramidal AB5E 5 1 octahedral 53

Know These Three

Electron Affinity - measurable, Cl is highest
Electronegativity is the ability of an atom to attract the shared electrons in a chemical bond toward itself. Electron Affinity - measurable, Cl is highest X (g) + e X-(g) Electronegativity - relative, F is highest 56

The Electronegativities of Common Elements
Linus Pauling mathematically determined the measure of attraction between an atom's nucleus and its valence electrons by analyzing bond strength of molecules. Scale is relative from 0 to 4.0 57

Polar covalent bond or polar bond is a covalent
Polar covalent bond or polar bond is a covalent bond with unequally shared electrons. Greater electron density around one of the two bonded atoms. Low Electronegative atom bonded to Very Electronegative electron rich region electron poor region d- d+ F H F H d+/- indicates a “partial charge” Unequally shared electron density

Polar Covalent bonds: unequally shared electrons caused
Polar Covalent bonds: unequally shared electrons caused by atoms of differing electronegativity H―F , H―Cl , H―Br , H―I The further away on the periodic table, the more polar the bond. Only slightly polar Very polar Non-polar covalent: Equally shared electron density in bond because of similar or equal electronegativity N―Cl C―H H―H

Classification of bonds by difference in electronegativity
Bond Type < 0.5 Non-polar Covalent 0.5 < and < 2 Polar Covalent  2 Ionic Increasing difference in electronegativity Covalent share e- Polar Covalent partial transfer of e- Ionic transfer e- 60

Bond Type Problem Classify the following bonds as ionic, polar covalent, or covalent: Strategy We follow the 2.0 rule of electronegativity difference and look up the values the bond in HCl The bond in KF the CC bond in H3C-CH3 the bond in AlP The electronegativity difference between H and Cl is 0.9 Therefore, the bond between H and Cl is polar covalent. (b) The difference between K and F is 3.2, above the 2.0 mark; therefore, the bond between K and F is ionic. (c) The two C atoms are identical; therefore 0, the bond between them is Non-polar covalent. (d) Despite being a Metal and Non-metal, the difference is only 0.6 and is a polar covalent

Q is the electronegativity difference
Dipole Moments: Quantitative measure of polarity. Di-pole = "two opposite poles" m = Q x r Debye units (m); 1 D = 3.36 x C•m Q is the electronegativity difference r is the distance between poles Arrow points to electron rich side

F - 4.0 Electronegativity (H - 2.1) I - 2.5 1.92 D Dipole moment 0.38 D Electronegativity difference (Q) has greater affect on m than distance (r).

These are experimentally measured values
Increased dipole moment = Increased polarity

Highly symmetric molecules are often non-polar
Bond dipoles can cancel for Molecule dipole moments Molecules can have polar bonds, but be non-polar when opposite dipoles cancel out. 0.0 D : : : : Highly symmetric molecules are often non-polar

Place in order from largest to smallest molecule polarity
A > C > B = D

1) Determine the Molecular Geometry 2) Draw dipole moments 3) Determine if the molecule is polar or non-polar. Trigonal Planar Non-polar Tetrahedral Non-polar Trigonal Pyramidal Polar Bent Polar Linear Non-Polar

Predict the relative polarity of each molecule
: : F―F : : Br–Cl : : : : 2.98 D 0.0 D 0.52 D 1.85 D : : : 1.15 D 0.0 D 1.60 D 1.90 D : : 0.0 D : : 120° 120° : 120° 2.91 D 1.66 D 0.0 D

Behavior of Polar Molecules in electric field
field off field on Strong order by electrostatic interaction Weak order by electrostatic interaction Crash Course: Polar & Non-Polar Molecules;

Substances with similar polarity are likely to
Substances with similar polarity are likely to fully interact with each other (Solubility). “like dissolves like” Polar molecules and Ionic compounds are soluble in polar solvents C2H5OH in H2O NaCl in H2O Non-polar molecules are soluble in non-polar solvents CCl4 in C6H6 Polar & non-polar don’t mix: Olive oil in water “Universal solvent” DIY Hydrophobic Coating:

Chemistry In Action: Microwave Ovens
Polar molecules (H2O) absorb microwaves to oscillate producing molecular friction.

Does not take energy changes into account.
Lewis Theory illustrates pairing of electrons, but fails to explain several chemical properties: Does not take energy changes into account. Does not explain resonance Expanded Octets Bond length and Strength Valence Bond Theory (developed by Linus Pauling) Assumes electrons in a molecule occupy atomic orbitals Better accounts for covalent bond stability and molecular geometries Molecular Orbital Theory Formation of molecular orbitals from atomic orbitals Better explains charge distribution from resonance

Energy of Atoms by distance of separation of electron clouds
H2 Electron density distribution H2 most stable at 74 pm bond length Distance maximizes the attraction of one nucleus to other atom’s electrons, while minimizing the e-/e- repulsions of the two electron clouds.

But this doesn’t agree with structural data
Carbon possesses 2s2 and 2p2 valence electrons. The 2s2 electrons are paired With only 2 unpaired electrons, How then can it make 4 bonds to form CH4? We could theorize a promotion of one of the 2s electrons to allow for 4 bonds But this doesn’t agree with structural data

The p orbitals overlap perpendicular and we’d
The p orbitals overlap perpendicular and we’d expect 90° angles for each H-C-H bond. Instead, the measured bond angles are found to be ° for a tetrahedral molecule. Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals 4 x Hybrid orbitals have very different shape from original atomic orbitals.

Formation of sp3 Hybrid Orbitals
Always the same number of orbitals before and after

Formation of Covalent Bonds in CH4
Covalent bonds are formed by: Overlap of hybrid orbitals with atomic orbitals Overlap of hybrid orbitals with other hybrid orbitals

Formation of sp2 Hybrid Orbitals
Each Fluorine uses an unhybridized p orbital to overlap with an sp2 to bond in BF3

Be only has 2 e- in the 2s orbital in the ground state
Formation of sp Hybrid Orbitals helps explain the linear geometry of BeCl2 H—Be—H Be only has 2 e- in the 2s orbital in the ground state By promoting a 2s electron to the 2p orbital, we get the excited state

The 2s and 2p orbitals then mix to form two hybrid orbitals
The two Be−H bonds are formed by the overlap of the Be sp orbitals with the 1s orbitals of the H atoms.

Hybridization example
Describe the hybridization state of phosphorus in phosphorus pentabromide (PBr5).

The hybridization process can be imagined as follows
The hybridization process can be imagined as follows. The orbital diagram of the ground-state P atom is Promoting a 3s electron into a 3d orbital results in the following excited state: Mixing the one 3s, three 3p, and one 3d orbitals generates five sp3d hybrid orbitals: These hybrid orbitals overlap the 4p orbitals of Br to form 5 covalent P−Br bonds in PBr5

How to predict the hybridization of the central atom?
Draw the Lewis structure of the molecule. Count the number of lone pairs AND the number of atoms bonded to the central atom # of Lone Pairs + # of Bonded Atoms Hybridization Examples 2 sp BeCl2 3 sp2 BF3 4 sp3 CH4, NH3, H2O 5 sp3d PCl5 6 sp3d2 SF6

Crash Course: Orbital Hybridization

+ Bonding in Ethylene, (Double bond)
Pi bond (p) – electron density above and below plane of nuclei of the bonding atoms Sigma bond (s) – electron density between the 2 atoms

p bonds can lock local regions into a planar (flat) geometry
View of p Bonding in Ethylene, C2H4 p bonds can lock local regions into a planar (flat) geometry 87

Bonding in Acetylene, C2H2 (Triple bond)
+ +

How many pi (p) bonds in the following molecule?
Single covalent bonds contain only 1 sigma (s) bond Double bonds contain 1 x s and 1 x pi (p) bond Triple bonds contain 1 x s and 2 x pi (p) bonds

Experiments show O2 is paramagnetic
Valence Bond theory can’t explain magnetic properties of oxygen gas O No unpaired e- Should be diamagnetic by regards of Valence bond theory Molecular orbital (MO) theory – states bonds are formed from combination of atomic orbitals to form molecular orbitals where bonded electrons exist in.

MO theory states atomic orbitals form molecular orbitals
A bonding molecular orbital has lower energy and greater stability than the atomic orbitals from which it was formed. An antibonding molecular orbital has higher energy and lower stability than the atomic orbitals from which it was formed. Describes bonds in terms of wave and interference properties Energy levels of bonding and antibonding molecular orbitals in H2. 91

*Explains the nature of O2 being paramagnetic
Still obeys rules and principles of electron filling. *Explains the nature of O2 being paramagnetic

Delocalized molecular orbitals are not confined
Delocalized molecular orbitals are not confined between two adjacent bonding atoms, but actually extend over three or more atoms. MO theory better predicts the properties of resonance and conjugated systems Example: Benzene, C6H6 Delocalized p orbitals 93

Drawn to show delocalization of double bonds
Electron density above and below the plane of the benzene molecule. Delocalized Electrons are free to move about the ring Can be written either way as a Lewis structure Drawn to show delocalization of double bonds August Kekulé

Sample of Things to Study
Octet rule and energy of bonding Electron dot notation Lattice Energy (and things that influence it) Covalent vs. Ionic (electrostatic) vs. Metallic bonds Bond length and strength factors Drawing Lewis structures (including polyatomic ions with formal charges) Isomers Resonance Why Carbon is special (getting its own group) Octet exceptions VSPER – What it is, what it determines, how to use it) Electronegativity Polarity (what it is, how to tell if it is, and why it’s important) Bond and Molecule Dipole moments, Partial charge) Hybridization (Why C has 4 bonds; Why S can have 6 bonds) Sigma vs Pi bonds Lewis vs Valence Bond vs Molecular Orbital Theories of bonding

Chapters 9 & 10: Chemical Bonding

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Chapters 9 & 10: Chemical Bonding

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