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Molecular Compounds

Chapter 5

Covalent (Molecular) Compounds

• Covalent Compound- a compound that containsatoms that are held together by covalent bonds

• Covalent Bond- the force of attraction betweenatoms that results from the sharing of one or morepairs of electrons to obtain an octet (or duet)

• Covalent compounds usually contain onlynonmetals in their formulas

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Formation of a water molecule

Covalent Bond for H2

• Orbitals of valence electrons overlap• Both electrons shared by each nuclei• Each nuclei now was a stable duet

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Repulsive and Attractive Forces

• Both nuclei repel eachother

• Both electrons repel eachother

• Each nuclei attracts bothelectrons (stronger)

• These forces are relatedto nuclei distance

• Bond length- optimumdistance between nuclei

The Seven Diatomic Elements

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Rationale for Covalent Bonding

• Atoms of electron rich elements shareelectrons to produce:– Electron configuration of the nearest noble gas– Molecular compounds with strong bonds

• An electron rich element is one that hashalf or more of its outer valence shell filled– All nonmetals– Hydrogen

Covalent Bonds and thePeriodic Table

• Covalent bonds can form between unlike atoms as wellas between like atoms, making possible a vast numberof molecular compounds.

• Examples– Water consists of 2 hydrogens covalently bound to an oxygen

atom– Ammonia consists of 3 hydrogens covalently bound to a

nitrogen atom– Methane consists of 4 hydrogens covalently bound to a

carbon atom

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Typical Number of Covalent BondsFormed by Main Group Elements

• Exceptions (numbers in parentheses above) for 3rd periodand greater due to use of d orbitals for bonding ANDBoron with only 6 instead of 8 electrons

Types of Covalent Bonds• Single Bonds

– A:B (A-B) one pair of electrons shared by 2 atoms• Double Bonds

– A::B (A=B) two pair of electrons shared by 2 atoms• Triple Bonds

– A:::B (A=B) three pair of electrons shared by 2 atoms• Bond Length- distance between centers of the bonded atoms

– Single > double > triple• Bond Strength

– Triple > double > single

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Coordinate Covalent Bonds• Coordinate covalent bond- formed from the

overlap of a filled orbital on one atom with avacant orbital of the other atom– Both electrons of the bond come from one atom

Molecular Formulas, StructuralFormulas and Lewis Structures

• Molecular formulas- the numbers and kindsof atoms in one molecule

• Structural formulas- a picture that use linesto show how atoms are connected bycovalent bonds

• Lewis structures- a picture that indicatesthe ways in which the valance electrons aredistributed in a molecule

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Lewis Structures• Lewis Structures- a picture indicating the

manner in which the valence electrons aredistributed in a molecule– Electrons in bonds are represented by lines or pairs of dots– Nonbonding electrons are represented by pairs of dots (lone

pair electrons)– Each atom is given a noble gas configuration, if possible

* 2 electrons around H* 8 electrons around all other atoms

– Each atom is given its preferred number of bonds, if possible* Preferred number of bonds = number of unpaired dots in thedot structure of the free atom

Common Bonding Patterns

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Guidelines For Lewis Structures

• Determine the TOTAL number of ALL valenceelectrons by adding the valence electrons inALL the atoms in the molecule; call this numberthe SUPPLY– For ions, be sure to add one e- for each negative

charge or subtract one e- for each positive charge• Determine the number of electrons NEEDED by

each atom to fulfill the octet/duet rule; call thisnumber the DEMAND

Guidelines For Lewis Structures• Determine the TOTAL number of bonds

required for the moleculeTotal # of bonds = (DEMAND - SUPPLY)/ (2e-/bond)

• Arrange the atoms in the correct skeletalstructure– The central atom is usually written first or used only

once in the formula– If in doubt try using the atom that has the greatest

number of unpaired electrons in its dot structure

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Guidelines For Lewis Structures• Place lines between atoms to represent bonds

– Each atom must be connected at least once– Form double or triple bonds if necessary to give the

central atom a noble gas configuration– The total number of bonds must equal the number

you calculated previously• Fill in the remaining valence electrons with dots

to achieve an octet around each atom (duet forH)

Guidelines For Lewis Structures

• Complicated molecules– Normal bonding number for certain atoms can be

exceeded (S, N, P, X) or reduced (O)– But the sum of bonds for the molecule will not

change– Therefore, the number of bonds on one atom may

increase but must be counterbalanced by a decreaseof bonds on another atom

– Examples- phosphate, sulfate, SO2

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Molecular Geometry• The geometry (shape) of a molecule about a

given atom depends on the number of groupssurrounding that atom its Lewis structure

• Group- an atom or a lone pair of electrons• Groups are arranged about a central atom in

such a way as to keep the groups as far apart aspossible

• Predict shapes using the valence-shell electron-pair (VSEPR) model

Valence-shell Electron-pairRepulsion (VSEPR) Model

• Negatively charged electron clouds in bondsand lone pair repel each other

• Bonded groups and lone pair electrons orientthemselves so that they are as far away formone another as possible

• This leads to specific geometric shapes• See Table 5.1 for relationship between number

of bonds, number of lone pairs and molecularshapes

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Molecular Geometry Relationships

Examples of Molecular Shapes

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Molecules Without Lone Pairs

Molecules With Lone Pairs

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Another Sample Problem• Describe the geometry about the indicated atom• Yellow = S, Black = C, Blue = N, White = H

Electronegativity• Electronegativity- A measure of the ability of an

atom to attract electrons in a covalent bondtoward itself– Values range form 0.7 to 4.0 (See Fig. 5.7)– Increase from left to right across a period– Decrease from top to bottom down a group– Typical Values: F (4.0), O (3.5), N (3.0),

C (2.5), S (2.5), P (2.1), H (2.1)– Differences in electronegativities lead to differences

in bond polarity

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Electronegativity and thePeriodic Table

Bond Polarity• Nonpolar Bond- a covalent bond in which the electrons

are shared equally between the bonded atoms– Bonded atoms have the same electronegativity

• Polar Bond- a covalent bond in which there is anunequal sharing of electrons between the bonded atoms– Bonded atoms have unequal electronegativities– The atom having the higher electronegativity has a slight

negative charge (δ-)– The atom having the lower electronegativity has a slight

positive charge (δ+)– The greater the difference in electronegativity, the more polar

the bond

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Sample Calculations forElectronegativity

• Differences up to 1.0 indicate more polar bonds• Differences over 2.0 indicate more ionic bonds

Molecular Polarity

• A molecule with polar covalent bonds does not mean itis necessarily polar overall

• A molecule is polar if the center of partial positivecharges does not coincide with the center of partialnegative charges within the molecule– If both atoms of a diatomic molecule are identical, the

molecule is nonpolar– If all attached groups around a central atom are identical, the

molecule is nonpolar– If a molecule is symmetrical, the molecule is nonpolar

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Examples of Polar and Nonpolar Molecules

Naming of Binary Molecular Compounds• Binary compounds contain two different elements• Give least electronegative element first• Use prefixes to indicate the number of atoms each

element has– Omit “mono” if only one atom of the first element is present– Omit the second “o” if the name of second element begins

with an “o”• Name the second element with its root name followed

by “ide”

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Properties of Molecular Compounds

• Most molecules do not break apart whenmolecular compounds melt, boil or dissolve inwater– Exception: acids ionize in water

• Molecular compounds may be solids, liquidsor gases at room temperature– Interactions (intermolecular forces) may be very

weak or relatively strong (e.g., quartz = SiO2)– See Table 5.3 for comparison to ionic compounds