Molecular Compounds. Chapter 5. Covalent (Molecular) Compounds

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1 Molecular Compounds Chapter 5 Covalent (Molecular) Compounds Covalent Compound- a compound that contains atoms that are held together by covalent bonds Covalent Bond- the force of attraction between atoms that results from the sharing of one or more pairs of electrons to obtain an octet (or duet) Covalent compounds usually contain only nonmetals in their formulas 1

2 Formation of a water molecule Covalent Bond for H 2 Orbitals of valence electrons overlap Both electrons shared by each nuclei Each nuclei now was a stable duet 2

3 Repulsive and Attractive Forces Both nuclei repel each other Both electrons repel each other Each nuclei attracts both electrons (stronger) These forces are related to nuclei distance Bond length- optimum distance between nuclei The Seven Diatomic Elements 3

4 Rationale for Covalent Bonding Atoms of electron rich elements share electrons to produce: Electron configuration of the nearest noble gas Molecular compounds with strong bonds An electron rich element is one that has half or more of its outer valence shell filled All nonmetals Hydrogen Covalent Bonds and the Periodic Table Covalent bonds can form between unlike atoms as well as between like atoms, making possible a vast number of 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 4

5 Typical Number of Covalent Bonds Formed by Main Group Elements Exceptions (numbers in parentheses above) for 3rd period and greater due to use of d orbitals for bonding AND Boron 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 5

6 Coordinate Covalent Bonds Coordinate covalent bond- formed from the overlap of a filled orbital on one atom with a vacant orbital of the other atom Both electrons of the bond come from one atom Molecular Formulas, Structural Formulas and Lewis Structures Molecular formulas- the numbers and kinds of atoms in one molecule Structural formulas- a picture that use lines to show how atoms are connected by covalent bonds Lewis structures- a picture that indicates the ways in which the valance electrons are distributed in a molecule 6

7 Lewis Structures Lewis Structures- a picture indicating the manner in which the valence electrons are distributed 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 the dot structure of the free atom Common Bonding Patterns 7

8 Guidelines For Lewis Structures Determine the TOTAL number of ALL valence electrons by adding the valence electrons in ALL the atoms in the molecule; call this number the 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 this number the DEMAND Guidelines For Lewis Structures Determine the TOTAL number of bonds required for the molecule Total # of bonds = (DEMAND - SUPPLY)/ (2e - /bond) Arrange the atoms in the correct skeletal structure 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 8

9 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 for H) 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 decrease of bonds on another atom Examples- phosphate, sulfate, SO 2 9

10 Molecular Geometry The geometry (shape) of a molecule about a given atom depends on the number of groups surrounding 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 as possible Predict shapes using the valence-shell electronpair (VSEPR) model Valence-shell Electron-pair Repulsion (VSEPR) Model Negatively charged electron clouds in bonds and lone pair repel each other Bonded groups and lone pair electrons orient themselves so that they are as far away form one another as possible This leads to specific geometric shapes See Table 5.1 for relationship between number of bonds, number of lone pairs and molecular shapes 10

11 Molecular Geometry Relationships Examples of Molecular Shapes 11

12 Molecules Without Lone Pairs Molecules With Lone Pairs 12

13 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 bond toward 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 13

14 Electronegativity and the Periodic 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 an unequal 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 14

15 Sample Calculations for Electronegativity 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 it is necessarily polar overall A molecule is polar if the center of partial positive charges does not coincide with the center of partial negative 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 15

16 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 16

17 Properties of Molecular Compounds Most molecules do not break apart when molecular compounds melt, boil or dissolve in water Exception: acids ionize in water Molecular compounds may be solids, liquids or gases at room temperature Interactions (intermolecular forces) may be very weak or relatively strong (e.g., quartz = SiO 2 ) See Table 5.3 for comparison to ionic compounds 17

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