CHEMICAL BONDING

The Chemical Bonding- attractive forces that hold atoms together in a compoundionic bonding- the transfer of one or more electrons from one atom or group of atoms to another.Ionic compoundsCovalent bonding- sharing of one or more electron pairs between atoms.Covalent compounds

IONIC BondingCation - an atom or group of atoms that has fewer electrons than protons- positively chargedAnion- is an atom or group of atoms that has more electrons that protons-negatively chargedMonoatomic ion consists of only one atom ex. Cl-, Mg2+Polyatomic ions contains more than one atom, NH4+, OH-, SO42-

---attraction of oppositely charged ions (Cations ad Anions) in large numbers to form a solid (ionic solid)

TERMSOxidized lose electrons to form cations (metals)Reduced- gain electrons to form anion (nonmetals)

Covalent Bonding

Valence Bond TheoryBasic PrincipleA covalent bond forms when the orbitals of two atoms overlap and the overlap region, which is between the nuclei, is occupied by a pair of electrons.The two wave functions are in phase so the amplitude increasesbetween the nuclei.

Two sine waves illustrate interference.

Two in phase sine waves add together with constructive interference; new amplitude is twice the original

Two out of phase sine waves add together with destructive interference; they cancel out

A set of overlapping orbitals has a maximum of two electrons that must have opposite spins.The greater the orbital overlap, the stronger (more stable) the bond.The valence atomic orbitals in a molecule are different from those in isolated atoms.There is a hybridization of atomic orbitals to form molecularorbitals.

Chemical Bonds and EnergyWhen two atoms approach one another, the negatively charged electron clouds are attracted to the other atoms positively charged nucleus.

The diagram represents electron density during bond formation.

Figure 11.1Orbital overlap and spin pairing in three diatomic molecules.Hydrogen, H2Hydrogen fluoride, HFFluorine, F2

Chemical Bonds and the Structure of MoleculesDuring ionic bond formation, the cations and anions achieve np6 electronic configurations (noble gas configuration).Metals lose electrons.Nonmetals gain electrons.

During covalent bond formation, electrons are shared between two atoms.Shared electrons are available to both bonding atoms.Sharing leads to 8 valence electrons around each atom.

Chemical Bonds and the Structure of MoleculesOCTET RULE - an atom will form covalent bonds to achieve a complement of eight valence electrons.

The valence shell electronic configuration is ns2np6 for a total of eight electrons.

For the n = 1 shell, hydrogen violates the octet rule and shares only 2 electrons

LEWIS DOT SYMBOLS for main group elements.Elements within a group have the same number of valence electrons and identical Lewis dot symbols.

The conceptual steps from molecular formula to the hybrid orbitals used in bonding.Molecular formulaLEWIS STRUCTUREGEOMETRY (Electronic/Ideal/ e- group arrangement and Molecular Shape

Hybrid orbitals

Writing Lewis StructuresLewis structures indicate how many bonds are formed and between which elements in a compound.

Step 1 - Count the total valence electrons in the molecule or ion.Sum the number of valence electrons for each element in a molecule.For ions, add or subtract valence electrons to account for the charge.

For the compound OF2, the number of valence electrons is 20.

Step 2 - Draw the skeletal structure of the molecule.The element written first in the formula is usually the central atom, unless it is hydrogen.Usually, the central atom is the LEAST ELECTRONEGATIVE.

Step 3 - Place single bonds between all connected atoms in the structure by drawing lines between them.A single line represents a bonding pair.Four electrons are placed in bonds.Sixteen electrons are left to place.

Step 4 - Place the remaining valence electrons not accounted for on individual atoms until the octet rule is satisfied. Place electrons as lone pairs whenever possible.Place electrons first on outer atoms, then on central atoms.Six electrons are placed as lone pairs on each F satisfies the octet rule for each F.The four remaining electrons are placed on the O to satisfy the octet rule for each O.

Step 5 - Create multiple bonds by shifting lone pairs into bonding positions as needed for any atoms that do not have a full octet of valence electrons.Correctly choosing which atoms to form multiple bonds between comes from experience.Multiple bonds are not required for OF2, as the octet rule is satisfied for each atom.

N - A = S ruleSimple mathematical relationship to help us write Lewis dot formulas.N = number of electrons needed to achieve a noble gas configuration.N usually has a value of 8 for representative elements. N has a value of 2 for H atoms.A = number of electrons available in valence shells of the atoms.A is equal to the periodic group number for each element. A is equal to 8 for the noble gases.S = number of electrons shared in bonds.A-S = number of electrons in unshared, lone, pairs.

For ions we must adjust the number of electrons available, A.Add one e- to A for each negative charge.Subtract one e- from A for each positive charge.The central atom in a molecule or polyatomic ion is determined by:The atom that requires the largest number of electrons to complete its octet goes in the center.For two atoms in the same periodic group, the less electronegative element goes in the center.

Formal ChargesFormal charges are analogous to oxidations numbers:They are not actual chargesThey keep track of electron ownership

Formal Charge = Valence e - no. of bonds no. of non-bonding e

EXAMPLEDraw the LEWIS Structure of the following molecules or ions:CO2 6. SOCl2 11. SO32- H2O7. SOF412. PCl5NH3 8. NH4+13. ClF5HF9. NO2-114. CO32-PCl3 10. SO3 15. C2H4

The conceptual steps from molecular formula to the hybrid orbitals used in bonding.Molecular formulaLewis structureGEOMETRY (Electronic/Ideal/e- group arrangement and Molecular ShapeHybrid orbitals

VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORYThe IDEAL/Electronic GEOMETRY of a molecule is determined by the way the electron pairs orient themselves in spaceThe orientation of electron pairs arises from electron repulsionsThe electron pairs spread out so as to minimize repulsion

The MOLECULAR GEOMETRY of a molecule is determined by the way the electron pairs orient themselves with respect to their relative strength to repel each other. The orientation of electron pairs arises from electron repulsionsThe electron pairs spread out so as to minimize repulsion

NOTE: DECREASING repulsion, DECREASING bond angleLP-LP > LP BP > BP BPLP lone (non-bonding) pair ; BP bonding pair

Copyright 2012 John Wiley & Sons, Inc.Klein, Organic Chemistry 1e 1-*

Methane has four equal bonds, so the bond angles are equal.HOW does the lone pair of ammonia affect its geometry?The bond angles in oxygen are even smaller. WHY?

Molecular Geometry

Figure 7.8 - Molecular Geometry Summary 2

The conceptual steps from molecular formula to the hybrid orbitals used in bonding.Molecular formulaLewis structureGEOMETRY (Electronic/Ideal/ e- group arrangement and Molecular ShapeHYBRID ORBITALS

Hybrid OrbitalsThe number of hybrid orbitals obtained equals the number of atomic orbitals mixed.The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed.Key PointsTypes of Hybrid Orbitals

Figure 11.2The sp hybrid orbitals in gaseous BeCl2.atomic orbitalshybrid orbitalsorbital box diagrams

Figure 11.3The sp2 hybrid orbitals in BF3.

Figure 11.4The sp3 hybrid orbitals in CH4.

Figure 11.5The sp3 hybrid orbitals in NH3.

Figure 11.5 continuedThe sp3 hybrid orbitals in H2O.

Figure 11.6The sp3d hybrid orbitals in PCl5.

Figure 11.7The sp3d2 hybrid orbitals in SF6.

Valence Bond : Hybrid Orbitals

EXAMPLEDetermine the geometry (electronic and molecular) and the hybrid orbital used by the central atom)of the following molecules or ions:CO2 6. SOCl2 11. SO32- H2O7. SOF412. PCl5NH3 8. NH4+13. ClF5HF9. NO2-114. CO32-PCl3 10. SO3 15. C2H4

*Polar Molecules: The Influence of Molecular GeometryMolecular geometry affects molecular polarity.Due to the effect of the bond dipoles and how they either cancel or reinforce each other.Polar Molecules must meet two requirements:One polar bond or one lone pair of electrons on central atom.Neither bonds nor lone pairs can be symmetrically arranged that their polarities cancel.

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EXAMPLEDetermine the following molecules or ions are polar or non-polar:CO2 6. SOCl2 11. SO32- H2O7. SOF412. PCl5NH3 8. NH4+13. ClF5HF9. NO2-114. CO32-PCl3 10. SO3 15. C2H4

Homework No. 2

MFLewis StructureElectronic GeometryMolecular GeometryHybrid OrbitalCN-1NO2SeCl4PO43-XeF4SF4H3O+BF3O3O2

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