Chapter 9 Covalent Bonding

I. The Covalent Bond A. Why do atoms bond? When two atoms need to gain electrons, they can share electrons to acquire a noble- gas configuration

B. What is a covalent bond? In a covalent bond, shared electrons are considered to be part of the complete outer energy level of both atoms involved Generally occurs when elements are relatively close to each other on the periodic table A majority form between nonmetallic elements

A molecule is formed when two or more atoms bond covalently Examples of covalent bond – carbohydrates, sugars, proteins, fats, and DNA

1. Formation of a covalent bond There are seven elements that naturally occur as diatomic molecules (H 2, N 2, O 2, Fl 2, Cl 2, Br 2, and I 2 )

Opposite charges of the protons of one atom and electrons of another atom are attracted to each other Like charges of the protons of each atom repel each other As a result, the most stable arrangement of atoms exists at the point of maximum attraction

C. Single Covalent Bonds When a single pair of electrons is shared, a single covalent bond forms Group 7A elements have seven valence electrons and can form one covalent bond with another element Group 6A element shave six valence electrons and can form two covalent bond and so on.

1. The sigma bond Single covalent bonds are also called sigma bonds, σ Occurs when the electron pair is shared in an area centered between the two atoms The valence atomic orbital of one atom overlaps or merges with the valence atomic orbital of another atom

Sigma bonds can form from the overlap between s orbitals, between an s and p orbital, or between 2 p orbitals

D. Multiple Covalent Bond Some atoms attain a noble-gas configuration by sharing more than one pair of electrons between two atoms Carbon, nitrogen, oxygen, and sulfur most often form multiple bonds

A double covalent bond occurs when two pairs of electrons are shared, O 2 A triple covalent bond occurs when three pairs of electrons are shared, N 2

1. The pi bond (π) Formed when parallel orbitals overlap to share electrons Electron pairs occupy the space above and below the line that represents where the two atoms are joined together Always accompanies a sigma bond when forming double and triple bonds

One pi bond in a double bond Two pi bonds in a triple bond

E. Strength of Covalent Bonds Depends on how much distance separates bonded nuclei Bond length is the distance between the two nuclei at the position of maximum attraction - Determined by atomic radius and how many electron pairs are shared

As the number of shared electron pairs increases, bond length decreases Single bonds are weaker than double bonds which are weaker than triple bonds

Energy is released when a bond forms (exothermic) Energy is absorbed when a bond breaks (endothermic)

The amount of energy required to break a specific covalent bond is tis bond dissociation energy Bond energy is directly related to bond length

II. Naming Molecules A. Naming Binary Molecular Compounds Many binary molecular compounds were discovered and given common names long before the modern naming system was developed Binary molecular compounds are composed of two different nonmetals and do not contain metals or ions

1. The first element in the formula is always named first, using the entire element name 2. The second element in the formula is named using the root of the element and adding the suffix –ide 3. Prefixes are used to indicate the number of atoms of each type that are present in the compound - The first element never uses the prefix mono-

B. Naming Acids 1. Naming binary acids Contains hydrogen and one other element Use the prefix hydro- to name the hydrogen part of the compound The rest of the name consists of a form of the root of the second element plus the suffix –ic, followed by the word acid

2. Naming oxyacids Any acid that contains hydrogen and an oxyanion If the oxyanion suffix is –ate, it is replaced with the suffix –ic If the oxyanion suffix is –ite, it is replaced with –ous They hydrogen in an oxyacid is not part of the name

III. Molecular Structures A. Structural Formulas Uses letter symbols and bonds to show relative positions of atoms 1. Predict the location of atoms - Hydrogen is always a terminal atom because it can share only one pair of electrons

2. Find the total number of electrons available for bonding 3. Determine the number of bonding pairs by dividing the number of electrons available for bonding by two 4. Place one bonding pair between the central atom and each of the terminal atoms 5. Place lone pairs around each terminal atom bonded to the central atom to satisfy the octet rule

6. If the central atom is not surrounded by four electron pairs, it does not have an octet - You must convert one or two of the lone pairs on the terminal atoms

B. Resonance Structures It is possible to have more than one correct Lewis structure when a molecule or polyatomic ion has both a double bond and a single bond Resonance is a condition that occurs when more than one valid Lewis structure can be written

Resonance structures differ only in the position of the electron pairs Molecules or ions behaves as if it has only one structure - Bond lengths are shorter than single bonds but longer than double bonds

C. Exceptions to the Octet Rule 1. Some molecules have an odd number of valence electrons and cannot form an octet around each atom 2. Some molecules have less than eight valence electrons and form a coordinate covalent bond with an atom with a lone pair of electrons 3. Some molecules have a center atom with more than eight valence electrons and forms an expanded octet

IV. Molecular Shape A. VSEPR Model The shape of the molecule determines whether or not molecules can get close enough to reach Based on an arrangement that minimizes the repulsion of shared and unshared pairs

Repulsions among electron pairs result in atoms exist at fixed angles, or bond angles - Supported by experimental evidence Lone pairs occupy a slightly larger orbital than shared electrons Shared bonding orbitals are pushed together slightly by lone pairs

1. Linear 180° 2 shared pairs 2. Trigonal planar 120° 3 shared pairs 3. Tetrahedral 109.5° 4 shared pairs 4. Trigonal pyramidal ° 3 shared pairs, 1 lone pair 5. Bent 104.5° 2 shared pairs, 2 lone pairs 6. Trigonal bipyramidal 90/120° 5 shared pairs 7. Octahedral 90° 6 shared pairs

C. Hybridization A hybrid results from combining two of the same type of object, and it has characteristics of both Atomic orbitals undergo hybridization during bonding to form new hybrid orbitals

V. Electronegativity and Polarity A. Electronegativity Difference and Bond Character Electronegativity indicates the relative ability of an atom to attract electrons in a chemical bond Chemical bonds can be predicted using the electronegativity difference of the elements in the bond

For identical atoms, with an electronegativity difference of zero, the electrons in the bond are equally shared, or covalent Chemical bonds between different atoms are never completely ionic or covalent - Depends on how strongly atoms are attracted to electrons

Unequal sharing of electrons between different elements result in a polar covalent bond Large differences indicate an electron was transferred from one atom to another, resulting in ionic bonds An electronegativity difference of 1.70 is considered 50% covalent and 50% ionic Ionic bonds generally have electronegativities above 1.70

B. Polar Covalent Bonds In a polar covalent bond, electrons spend more time around one atom than another Partial changes occur at the ends of the bond, δ + and δ - Referred to as a dipole (two poles)

1. Molecular polarity Nonpolar molecules are not attracted by an electric field Polar molecules tend to align with an electric field because of their dipole

2. Polar molecule or not? Symmetric molecules are usually nonpolar Asymmetrical molecules are usually polar

3. Solubility of polar molecules The ability of a substance to dissolve in another substance Polar molecules and ionic compounds are usually soluble in polar substances Nonpolar molecules dissolve only in nonpolar substances

C. Properties of Covalent Compounds Differences in properties are a result of differences in attractive forces Weak forces of attractions are known as intermolecular forces, or van der Waal forces In nonpolar substances, weak attractions are called dispersion forces

The force in polar molecules is stronger and is called a dipole-dipole force A hydrogen bond is a strong intermolecular force that is formed between a hydrogen and a dipole end of another atom Physical properties such as melting point, boiling point, hardness, etc. are determined by these intermolecular forces