Ionic Bonding, LDS, Polarity, Geometry, Formal Charges, Hybridization AP Chemistry Lovrencic
Formal Charge
Formal charge is a fictitious charge assigned to each atom in a Lewis structure that helps us to distinguish among competing Lewis structures. In a Lewis structure, we calculate an atom's formal charge, which indicates the charge it would have if all bonding electrons were shared equally between the bonded atom. FC = # valence e− − [nonbonding e− + ½ bonding e−] Sum of all the formal charges in a molecule = 0. -In an ion, total equals the charge.
sp2 Hybridization
Hybrid orbitals will overlap on axis with orbitals from other atoms. Unhybridized p orbital will overlap sideways, or side by side, with an unhybridized p orbital of another atom.
Lewis Theory and Ionic Bonding
Lewis symbols can be used to represent the transfer of electrons from a metal atom to a nonmetal atom, resulting in ions that are attracted to each other and, therefore, bond.
Exceptions to the Octet Rule
Odd number electron species (e.g., NO) -Will have one unpaired electron -Free-radical -Very reactive Incomplete octets -B, Be, Al, H Expanded octets -Elements with empty d orbitals can have more than eight electrons.
Valence Bond Theory: Hybridization
One of the issues that arises is that the number of partially filled or empty atomic orbitals did not predict the number of bonds or orientation of bonds. -C = 2s22px12py12pz0 would predict two or three bonds that are 90° apart, rather than four bonds that are 109.5° apart. To adjust for these inconsistencies, it was postulated that the valence atomic orbitals could hybridize before bonding took place. -One hybridization of C is to mix all the 2s and 2p orbitals to get four orbitals that point at the corners of a tetrahedron.
Representing Three-Dimensional Shapes on Paper
One of the problems with drawing molecules is trying to show their dimensionality. By convention, the central atom is put in the plane of the paper. Put as many other atoms as possible in the same plane and indicate with a straight line. For atoms in front of the plane, use a solid wedge. For atoms behind the plane, use a hashed wedge. Straight line Bond in plane of paper Hatched wedge Bond going into the paper Solid wedge Bond coming out of the paper
Lewis theory
One of the simplest bonding theories
Predictions of Molecular Formulas by Lewis Theory
Oxygen is more stable when it is singly bonded to two other atoms.
Electronegativity Difference (∆EN) medium Polar covalent HCl
Polar Covalent Bond Electrons shared unequally 0.4 to 2.0
Molecular Polarity Affects Solubility in Water
Polar molecules are attracted to other polar molecules. Because water is a polar molecule, other polar molecules dissolve well in water. -And ionic compounds as well Some molecules have both polar and nonpolar parts.
Valence Shell Electron Pair Repulsion Theory (VSEPR)
Properties of molecular substances depend on the structure of the molecule. Valence shell electron pair repulsion (VSEPR) theory is a simple model that allows us to account for molecular shape. Electron groups are defined as lone pairs, single bonds, double bonds, and triple bonds. VSEPR is based on the idea that electron groups repel one another through coulombic forces.
Electronegativity Difference (∆EN) small Covalent Cl₂
Pure (non polar) Covalent Bond Electrons shared equally 0.0 to 0.4
Hybridization
Some atoms hybridize their orbitals to maximize bonding. -More bonds = more full orbitals = more stability Hybridizing is mixing different types of orbitals in the valence shell to make a new set of degenerate orbitals. -sp, sp2, sp3, sp3d, sp3d2 The same type of atom can have different types of hybridization, depending on the total number of electron regions. -C = sp, sp2, sp3
A negative energy change (exothermic)
indicates a stable anion is formed. The larger the negative number, the more stable the anion. Small negative energies indicate a less stable anion.
A positive energy change (endothermic)
indicates the anion is unstable.
Orbital Diagrams of Bonding
"Overlap" between a hybrid orbital on one atom with a hybrid or nonhybridized orbital on another atom results in a σ bond. "Overlap" between unhybridized p orbitals on bonded atoms results in a π bond. Hybrid orbitals overlap to form a σ bond. Unhybridized p orbitals overlap to form a π bond.
To calculate this potential energy, you need to consider the following interactions:
-Nucleus-to-nucleus repulsions -Electron-to-electron repulsions -Nucleus-to-electron attractions
Born-Haber Cycle
-Use Hess's law to add up enthalpy changes of other reactions to determine the lattice energy. -ΔH°f(salt) = ΔH°f(metal atoms, g) + ΔH°f(nonmetal atoms, g) + ΔH°f(cations, g) + ΔH°f(anions, g) + ΔH° (crystal lattice) -ΔH° (crystal lattice) = lattice energy -For metal atom(g) → cation(g), ΔH°f = first ionization energy •Don't forget to add together all the ionization energies to get to the desired cation. -M2+ = 1st IE + 2nd IE -For nonmetal atoms (g) → anions (g), ΔH°f = electron affinity
Predicting Molecular Geometry
1) Draw the Lewis structure. 2) Determine the number of electron groups around the central atom. 3) Classify each electron group as a bonding or lone pair, and then count each type. -Remember, multiple bonds count as one group. 4) Use Table 10.1 to determine the shape and bond angles.
Valence Bond Theory: Main Concepts
1) The valence electrons of the atoms in a molecule reside in quantum-mechanical atomic orbitals. The orbitals can be the standard s, p, d, and f orbitals, or they may be hybrid combinations of these. 2) A chemical bond results when these atomic orbitals interact and there is a total of two electrons in the new molecular orbital. -The electrons must be spin paired. 3) The shape of the molecule is determined by the geometry of the interacting orbitals.
Writing Lewis Structures of Molecules
1) Write the correct skeletal structure for the molecule. -Hydrogen atoms are always terminal. The more electronegative atoms are placed in terminal positions. 2) Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule. 3) Distribute the electrons among the atoms, giving octets (or duets in the case of hydrogen) to as many atoms as possible. 4) If any atoms lack an octet, form double or triple bonds as necessary to give them octets.
Predicting Polarity of Molecules
1. Draw the Lewis structure, and determine the molecular geometry. 2. Determine whether the bonds in the molecule are polar. a) If there are no polar bonds, the molecule is nonpolar. 3. Determine whether the polar bonds add together to give a net dipole moment.
Predicting Hybridization and Bonding Scheme
1. Start by drawing the Lewis structure. 2. Use VSEPR theory to predict the electron group geometry around each central atom. 3. Use Table 10.3 to select the hybridization scheme that matches the electron group geometry. 4. Sketch the atomic and hybrid orbitals on the atoms in the molecule, showing overlap of the appropriate orbitals. 5. Label the bonds as σ or π.
Types of Bonds:
A sigma (σ) bond results when the interacting atomic orbitals point along the axis connecting the two bonding nuclei. -Either standard atomic orbitals or hybrids •s to s, p to p, hybrid to hybrid, s to hybrid, etc. A pi (π) bond results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei. -Between unhybridized parallel p orbitals The interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore, σ bonds are stronger than π bonds.
Why Do Atoms Bond?
Chemical bonds form because they lower the potential energy between the charged particles that compose atoms. A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms.
Bonding with Valence Bond Theory
According to valence bond theory, bonding takes place between atoms when their atomic or hybrid orbitals interact. -"Overlap" To interact, the orbitals must either -be aligned along the axis between the atoms, or -be parallel to each other and perpendicular to the interatomic axis.
Orbital Interaction
As two atoms approach each other, the half-filled valence atomic orbitals on each atom would interact to form molecular orbitals. -Molecular orbitals are regions of high probability of finding the shared electrons in the molecule. The molecular orbitals would be more stable than the separate atomic orbitals because they would contain paired electrons shared by both atoms. -The potential energy is lowered when the molecular orbitals contain a total of two paired electrons compared to separate, one-electron atomic orbitals.
sp3d Hybridization
Atom with five electron groups around it -Trigonal bipyramid electron geometry -Seesaw, T-shape, linear -120° and 90° bond angles Use empty d orbitals from valence shell
sp3 Hybridization
Atom with four electron groups around it -Tetrahedral geometry -109.5° angles between hybrid orbitals Atom uses hybrid orbitals for all bonds and lone pairs.
sp3d2
Atom with six electron groups around it -Octahedral electron geometry -Square pyramid, Square planar -90° bond angles Use empty d orbitals from valence shell to form hybrid
sp Hybridization
Atom with two electron groups; for C2H2 -Linear shape -180° bond angle Atom uses hybrid orbitals for σ bonds or lone pairs and uses nonhybridized p orbitals for π bonds Usually will for two σ bonds and two π bonds
Lewis Bonding Theory
Atoms bond because bonding results in a more stable electron configuration. -More stable = lower potential energy Atoms bond together by either transferring or sharing electrons. Usually, this results in all atoms obtaining an outer shell with eight electrons. -Octet rule -There are some exceptions to this rule: The key to remember is to try to get an electron configuration like a noble gas.
Bond Rotation
Because the orbitals that form the σ bond point along the internuclear axis, rotation around that bond does not require breaking the interaction between the orbitals. But, the orbitals that form the π bond interact above and below the internuclear axis, so rotation around the axis requires the breaking of the interaction between the orbitals.
Evaluating Resonance Structures
Better structures have fewer formal charges. Better structures have smaller formal charges. Better structures have the negative formal charge on the more electronegative atom.
Electron affinity (E.A.) Trends
Broadly speaking, the trend is toward more negative electron affinities going from left to right in a period.
Nonmetal and nonmetal
Covalent Electrons shared Ice, H₂O(s), H₂O molecules
Polar Covalent Bonding
Covalent bonding between unlike atoms results in unequal sharing of the electrons. -One atom pulls the electrons in the bond closer to its side. -One end of the bond has larger electron density than the other. The result is a polar covalent bond. Bond polarity -The end with the larger electron density gets a partial negative charge. -The end that is electron deficient gets a partial positive charge.
Bond Dipole Moments
Dipole moment, μ, is a measure of bond polarity. -A dipole is a material with a + and − end. -It is directly proportional to the size of the partial charges and directly proportional to the distance between them. • μ = (q)(r) • Not Coulomb's law • Measured in Debyes, D Generally, the more electrons two atoms share and the larger the atoms are, the larger the dipole moment.
VSEPR Theory
Electron groups around the central atom will be most stable when they are as far apart as possible. We call this VSEPR theory. -Because electrons are negatively charged, they should be most stable when they are separated as much as possible. The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule.
Covalent Bonding: Bonding and Lone Pair Electrons
Electrons that are shared by atoms are called bonding pairs. Electrons that are not shared by atoms but belong to a particular atom are called lone pairs. -Also known as nonbonding pairs
Crystal Lattice
Electrostatic attraction is nondirectional! -No direct anion-cation pair Therefore, there is no ionic molecule. -The chemical formula is an empirical formula, simply giving the ratio of ions based on charge balance.
Bonding Theories
Explain how and why atoms attach together to form molecules Explain why some combinations of atoms are stable and others are not Why is water H2O, not HO or H3O? Can be used to predict the shapes of molecules Can be used to predict the chemical and physical properties of compounds
Electronegativity Difference and Bond Type
If the difference in electronegativity between bonded atoms is 0, the bond is pure covalent. -Equal sharing If the difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent. If the difference in electronegativity between bonded atoms is 0.4 to 1.9, the bond is polar covalent. If the difference in electronegativity between bonded atoms is larger than or equal to 2.0, the bond is "100%" ionic.
Lewis Structures of Atoms
In a Lewis structure, we represent the valence electrons of main-group elements as dots surrounding the symbol for the element. -Also known as electron dot structures We use the symbol of the element to represent the nucleus and inner electrons.
Trends in Bond Lengths
In general, the more electrons two atoms share, the shorter the covalent bond. -Must be comparing bonds between like atoms -C≡C (120 pm) < C═C (134 pm) < C—C (154 pm) -C≡N (116 pm) < C═N (128 pm) < C—N (147 pm) Generally, bond length decreases from left to right across period. -C—C (154 pm) > C—N (147 pm) > C—O (143 pm) Generally, bond length increases down the column. -F—F (144 pm) < Cl—Cl (198 pm) < Br—Br (228 pm) In general, as bonds get longer, they also get weaker.
Metal and nonmetal
Ionic Electrons transferred Table Salt, NaCl(s), Na⁺ and Cl⁻ ions in a crystal lattice
Electronegativity Difference (∆EN) large Ionic NaCl
Ionic Bond Electrons transferred 2.0 to 3.3
Lewis Model
Lewis theory emphasizes valence electrons to explain bonding. Using Lewis theory, we can draw models, called Lewis structures. -Also known as electron dot structures Lewis structures allow us to predict many properties of molecules. -Molecular stability, shape, size, and polarity
Problems with Lewis Theory
Lewis theory generally predicts trends in properties but does not give good numerical predictions. -For example, bond strength and bond length Lewis theory gives good first approximations of the bond angles in molecules but usually cannot be used to get the actual angle. Lewis theory cannot write one correct structure for many molecules where resonance is important. Lewis theory often does not predict the correct magnetic behavior of molecules. -For example, O2 is paramagnetic, although the Lewis structure predicts it is diamagnetic.
Lewis Theory of Covalent Bonding
Lewis theory implies that another way atoms can achieve an octet of valence electrons is to share their valence electrons with other atoms. The shared electrons would then count toward each atom's octet. The sharing of valence electrons is called covalent bonding.
Covalent Bonding: Model versus Reality
Lewis theory implies that some combinations should be stable, whereas others should not. -Because the stable combinations result in "octets" Using these ideas from the Lewis theory allows us to predict the formulas of molecules of covalently bonded substances. Hydrogen and the halogens are all diatomic molecular elements, as predicted by Lewis theory. Oxygen generally forms either two single bonds or a double bond in its molecular compounds, as predicted by Lewis theory. -Though, as we'll see, there are some stable compounds in which oxygen has one single bond and another where it has a triple bond, but it still has an octet.
Resonance
Lewis theory localizes the electrons between the atoms that are bonding together. Extensions of Lewis theory suggest that there is some degree of delocalization of the electrons; we call this concept resonance. Delocalization of charge helps to stabilize the molecule. The organic molecule benzene has six σ bonds and a p orbital on each carbon atom. In reality the π electrons in benzene are not localized, but delocalized. The even distribution of the π electrons in benzene makes the molecule unusually stable.
Covalent Bonding: Model versus Reality 2
Lewis theory of covalent bonding implies that the attractions between atoms are directional. -The shared electrons are most stable between the bonding atoms. Therefore, Lewis theory predicts covalently bonded compounds will be found as individual molecules. -Rather than an array like ionic compounds Compounds of nonmetals are made of individual molecule units.
Covalent Bonding: Model versus Reality 5
Lewis theory predicts that neither molecular solids nor liquids should conduct electricity. -There are no charged particles around to allow the material to conduct. Molecular compounds do not conduct electricity in the solid or liquid state. Molecular acids conduct electricity when dissolved in water but not in the solid or liquid state, due to them being ionized by the water.
Covalent Bonding: Model versus Reality 4
Lewis theory predicts that the hardness and brittleness of molecular compounds should vary depending on the strength of intermolecular attractive forces. -The kind and strength of the intermolecular attractions vary based on many factors. Some molecular solids are brittle and hard, but many are soft and waxy.
Covalent Bonding: Model versus Reality 3
Lewis theory predicts that the melting and boiling points of molecular compounds should be relatively low. -This involves breaking the attractions between the molecules but not the bonds between the atoms. -The covalent bonds are strong, but the attractions between the molecules are generally weak. Molecular compounds have low melting points and boiling points. -Melting points generally < 300 °C -Molecular compounds are found in all three states at room temperature.
Covalent Bonding: Model versus Reality 7
Lewis theory predicts that the more electrons two atoms share, the shorter the bond should be. -When comparing bonds to like atoms Bond length is determined by measuring the distance between the nuclei of bonded atoms. In general, triple bonds are shorter than double bonds, and double bonds are shorter than single bonds.
Covalent Bonding: Model versus Reality 6
Lewis theory predicts that the more electrons two atoms share, the stronger the bond should be. Bond strength is measured by how much energy must be added into the bond to break it in half. In general, triple bonds are stronger than double bonds, and double bonds are stronger than single bonds. -However, Lewis theory would predict that double bonds are twice as strong as single bonds; the reality is that they are less than twice as strong.
Lewis Theory Predictions for Ionic Bonding
Lewis theory predicts the number of electrons a metal atom should lose or a nonmetal atom should gain in order to attain a stable electron arrangement. -The octet rule This allows us to predict the formulas of ionic compounds that result. It also allows us to predict the relative strengths of the resulting ionic bonds from Coulomb's law.
Derivatives of the Trigonal Bipyramidal Electron Geometry
Lone pairs on central atoms with five electron groups will occupy the equatorial positions because there is more room. The result is called the seesaw shape (aka distorted tetrahedron). When there are two lone pairs around the central atom, the result is T-shaped. When there are three lone pairs around the central atom, the result is a linear shape. The bond angles between equatorial positions are less than 120°. The bond angles between axial and equatorial positions are less than 90°. -Linear = 180° axial to axial.
Multiple Central Atoms
Many molecules have larger structures with many interior atoms. We can think of them as having multiple central atoms. When this occurs, we describe the shape around each central atom in sequence. -The shape around N is trigonal pyramidal. -The shape around left C is tetrahedral. -The shape around right C is trigonal planar. -The shape around right O is tetrahedral-bent.
Metal and metal
Metallic Electrons pooled Sodium metal, Na(s), Na⁺ and e⁻ sea where the electrons forget when Na atom they belong to, help carry electric currents easier in the sea of electrons
Stable Electron Arrangements and Ion Charge
Metals form cations by losing valence shell electrons. Nonmetals form anions by gaining valence electrons.
Dipole Moments of Several Molecules in the Gas Phase
Molecule ∆EN Dipole Moment (M) Cl₂ ______0 _____0 ClF ______1.0 ___0.88 HF ______1.9 ___1.82 LiF ______3.0 ___6.33
Bond Polarity
Most bonds have some degree of sharing and some degree of ion formation to them. Bonds are classified as covalent if the amount of electron transfer is insufficient for the material to display the classic properties of ionic compounds. If the sharing is unequal enough to produce a dipole in the bond, the bond is classified as polar covalent.
Born-Haber Cycle for NaCl
Na(s) → Na(g) +108 kJ ½ Cl₂ (g) → Cl(g) +½(244 kJ) Na(g) → Na+(g) +496 kJ Cl (g) → Cl−(g) −349 kJ Na+ (g) + Cl−(g) → NaCl(s) ΔH (NaCl lattice) Na(s) + ½ Cl2(g) → NaCl(s) −411 kJ ΔH°f(NaCl, s) = ΔH°f (Na atoms, g) + ΔH°f(Cl—Cl bond energy) + ΔH°f (IE1 of Na) + ΔH°f (EA of Cl) + ΔH°(NaCl lattice) NaCl lattice energy = (−411 kJ) − [(+108 kJ) + (+122 kJ) + (+496 kJ) + (−349 kJ) ] = −788 kJ
Covalent Bonds
Nonmetal atoms have relatively high ionization energies, so it is difficult to remove electrons from them. When nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons. -Potential energy is lowest when the electrons are between the nuclei. Shared electrons hold the atoms together by attracting nuclei of both atoms.
Determining Lattice Energy: The Born-Haber Cycle
The Born-Haber cycle is a hypothetical series of reactions that represents the formation of an ionic compound from its constituent elements. The reactions are chosen so that the change in enthalpy of each reaction is known except for the last one, which is the lattice energy.
Molecular Polarity
The H─Cl bond is polar. The bonding electrons are pulled toward the Cl end of the molecule. The net result is a polar molecule. The O─C bond is polar. The bonding electrons are pulled equally toward both O ends of the molecule. The net result is a nonpolar molecule. The H─O bond is polar. Both sets of bonding electrons are pulled toward the O end of the molecule. Because the molecule is bent, not linear, the net result is a polar molecule.
Electron Groups
The Lewis structure predicts the number of valence electron pairs around the central atom(s). Each lone pair of electrons constitutes one electron group on a central atom. Each bond constitutes one electron group on a central atom, regardless of whether it is single, double, or triple. There are three electron groups on N: -Three lone pair -One single bond -One double bond
Electronegativity
The ability of an atom to attract bonding electrons to itself Increases across period (left to right) and decreases down group (top to bottom) -Fluorine is the most electronegative element. -Francium is the least electronegative element. -Noble gas atoms are not assigned values. -Opposite of atomic size trend The larger the difference in electronegativity, the more polar the bond. -Negative end toward more electronegative atom
The Effect of Lone Pairs
The actual geometry of the molecule may be different from the electron geometry. Lone pair electrons typically exert slightly greater repulsion than bonding electrons, affecting the bond angles. A lone electron pair is more spread out in space than a bonding electron pair because a lone pair is attracted to only one nucleus while a bonding pair is attracted to two nuclei. In general, electron group repulsions vary as follows: -Lone pair-lone pair > lone pair-bonding pair > bonding pair-bonding pair
Effect of Lone Pairs
The bonding electrons are shared by two atoms, so some of the negative charge is removed from the central atom. The nonbonding electrons are localized on the central atom, so the area of negative charge takes more space.
Determining the Number of Valence Electrons in an Atom
The column number on the periodic table will tell you how many valence electrons a main group atom has. -Transition elements all have two valence electrons. Why?
Bond Lengths
The distance between the nuclei of bonded atoms is called the bond length. Because the actual bond length depends on the other atoms around the bond, we often use the average bond length. -Averaged for similar bonds from many compounds Bond order is the number of chemical bonds between a pair of atoms. In molecules that have resonance or nonclassical bonding, bond order may not be an integer.
Ionic Bonding and the Crystal Lattice
The extra energy that is released comes from the formation of a structure, called a crystal lattice, in which every cation is surrounded by anions, and vice versa. The crystal lattice is held together by the electrostatic attraction of the cations for all the surrounding anions. The crystal lattice maximizes the attractions between cations and anions, leading to the most stable arrangement.
Lattice Energy
The extra stability that accompanies the formation of the crystal lattice is measured as the lattice energy. The lattice energy is the energy released when the solid crystal forms from separate ions in the gas state. -Always exothermic -Hard to measure directly, but can be calculated from knowledge of other processes Lattice energy depends directly on the size of charges and inversely on distance between ions.
Trends in Lattice Energy: Ion Size
The force of attraction between charged particles is inversely proportional to the distance between them. Larger ions mean the center of positive charge (nucleus of the cation) is farther away from the negative charge (electrons of the anion). -Less exothermic lattice energy with larger ionic radius -More exothermic lattice energy with increasing magnitude of ionic charge
Trends in Lattice Energy: Ion Charge
The force of attraction between oppositely charged particles is directly proportional to the product of the charges. Larger charge means the ions are more strongly attracted. -Larger charge = stronger attraction -Stronger attraction = larger lattice energy Of the two factors, ion charge is generally more important.
Hybrid Orbitals
The number of standard atomic orbitals combined = the number of hybrid orbitals formed. -Combining a 2s with a 2p gives two 2sp hybrid orbitals. -H cannot hybridize! •Its valence shell has only one orbital. The number and type of standard atomic orbitals combined determines the shape of the hybrid orbitals. The particular kind of hybridization that occurs is the one that yields the lowest overall energy for the molecule.
Percent Ionic Character
The percent ionic character is the percentage of a bond's measured dipole moment compared to what it would be if the electrons were completely transferred. The percent ionic character indicates the degree to which the electron is transferred.
Metallic Bonds
The relatively low ionization energy of metals allows them to lose electrons easily. The simplest theory of metallic bonding involves the metal atoms releasing their valence electrons to be shared as a pool by all the atoms/ions in the metal. -An organization of metal cation islands in a sea of electrons -Electrons delocalized throughout the metal structure Bonding results from attraction of cation for the delocalized electrons.
Electron Group Geometry
There are five basic arrangements of electron groups around a central atom. -Based on a maximum of six bonding electron groups •Though there may be more than six on very large atoms, it is very rare. Each of these five basic arrangements results in five different basic electron geometries. -In order for the molecular shape and bond angles to be a "perfect" geometric figure, all the electron groups must be bonds, and all the bonds must be equivalent. For molecules that exhibit resonance, it doesn't matter which resonance form you use since the electron geometry will be the same.
Limitations of Valence Bond (VB) Theory
VB theory predicts many properties better than Lewis theory. -Bonding schemes, bond strengths, bond lengths, bond rigidity VB theory presumes the electrons are localized in orbitals on the atoms in the molecule; it doesn't account for delocalization. There are still many properties of molecules it doesn't predict perfectly. -Magnetic behavior of O2
Valence Bond Theory
Valence Bond theory (VB) approaches chemical bonding based on an extension of the quantum-mechanical model. When orbitals on atoms interact, they make a bond. These orbitals are hybridized atomic orbitals, a kind of blend or combination of two or more standard atomic orbitals. When two atoms approach each other, the electrons and nucleus of one atom interact with the electrons and nucleus of the other atom. If the energy of the system is lowered because of the interactions, a chemical bond forms. A chemical bond results from the overlap of two half-filled orbitals with spin-pairing of the two valence electrons (or less commonly the overlap of a completely filled orbital with an empty orbital). The geometry of the overlapping orbitals determines the shape of the molecule.
Valence Electrons and Bonding
Valence electrons are held most loosely. Chemical bonding involves the transfer or sharing of electrons between two or more atoms. Because of the two previously listed facts, valence electrons are most important in bonding. Lewis theory focuses on the behavior of the valence electrons.
Types of Bonds
We can classify bonds based on the kinds of atoms that are bonded together. Ionic Covalent Metallic
Lewis Structures of Atoms Steps
We use dots around the symbol to represent valence electrons. -Pair the first two dots for the s orbital electrons. -Put one dot on each open side for the first three p electrons. -Then, pair the rest of the dots for the remaining p electrons.
Ionic Bonds
When a metal atom loses electrons it becomes a cation. -Metals have low ionization energy, making it relatively easy to remove electrons from them. When a nonmetal atom gains electrons it becomes an anion. -Nonmetals have high electron affinities, making it advantageous to add electrons to these atoms. The oppositely charged ions are then attracted to each other, resulting in an ionic bond.
Five Electron Groups: Trigonal Bipyramidal Electron Geometry
When there are five electron groups around the central atom, they will occupy positions in the shape of two tetrahedra that are base to base with the central atom in the center of the shared bases. This results in the electron groups taking a trigonal bipyramidal geometry. The positions above and below the central atom are called the axial positions. The positions in the same base plane as the central atom are called the equatorial positions. The bond angle between equatorial positions is 120°. The bond angle between axial and equatorial positions is 90°.
Pyramidal and Bent Molecular Geometries: Derivatives of Tetrahedral Electron Geometry
When there are four electron groups around the central atom, and one is a lone pair, the result is called a pyramidal shape, because it is a triangular-base pyramid with the central atom at the apex. When there are four electron groups around the central atom, and two are lone pairs, the result is called a tetrahedral-bent shape.
Four Electron Groups: Tetrahedral Electron Geometry
When there are four electron groups around the central atom, they will occupy positions in the shape of a tetrahedron around the central atom. This results in the electron groups taking a tetrahedral geometry. The bond angle is 109.5°.
Derivatives of the Octahedral Geometry
When there are lone pairs around a central atom with six electron groups, each even number lone pair will take a position opposite the previous lone pair. When one of the six electron groups is a lone pair, the result is called a square pyramid shape. The bond angles between axial and equatorial positions are less than 90°. When two of the six electron groups are lone pairs, the result is called a square planar shape. The bond angles between equatorial positions are 90°.
Octahedral Electron Geometry
When there are six electron groups around the central atom, they will occupy positions in the shape of two square-base pyramids that are base to base with the central atom in the center of the shared bases. This results in the electron groups taking an octahedral geometry. -It is called octahedral because the geometric figure has eight sides. All positions are equivalent. The bond angle is 90°.
Three Electron Groups: Trigonal Planar Electron Geometry
When there are three electron groups around the central atom, they will occupy positions in the shape of a triangle around the central atom. This results in the electron groups taking a trigonal planar geometry. The bond angle is 120°.
Two Electron Groups: Linear Electron Geometry
When there are two electron groups around the central atom, they will occupy positions on opposite sides of the central atom. This results in the electron groups taking a linear geometry. The bond angle is 180°.
Resonance Structures
When there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures. The actual molecule is a combination of the resonance forms—a resonance hybrid. -The molecule does not resonate between the two forms, though we often draw it that way. Look for multiple bonds or lone pairs.
Single Covalent Bonds
When two atoms share one pair of electrons, it is called a single covalent bond. -Two electrons One atom may use more than one single bond to fulfill its octet. -Can bond to diff atoms -H only duet
Triple Covalent Bond
When two atoms share three pairs of electrons the result is called a triple covalent bond. -Six electrons
Double Covalent Bond
When two atoms share two pairs of electrons the result is called a double covalent bond. -Four electrons
