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Home , Electron pair, Octet rule, Unpaired electron

CHAPTER NOTES – CHAPTER 16

Covalent Bonding

Goals : To gain an understanding of :

1. and electron dot notation. 2. Stable electron configurations. 3. Covalent bonding. 4. Polarity of bonds and

NOTES:

Valence are the electrons in the highest energy level of an . For example, in the calcium atom ( 1s22s22p63s23p64s2) the 4s2 electrons are the valence electrons. In the titanium atom (electron configuration 1s22s22p63s23p64s23d2) The 4s2 electrons are still the valence electrons -they are in the highest energy level. In the phosphorus atom (electron configuration 1s22s22p63s23p3) the 3s23p3 are the valence electrons.

The numbers of the representative elements are:

• Group 1 - 1 valence electron • Group 2 - 2 valence electrons • Group 13 - 3 valence electrons • Group 14 - 4 valence electrons • Group 15 - 5 valence electrons • Group 16 - 6 valence electrons • Group 17 - 7 valence electrons • Group 18 - 8 valence electrons

The Noble gases (Group 0) have a stable electron configuration (s2p6) with 8 electrons filling the outer s and p orbitals. This stability comes from the low energy state of this configuration and also accounts for the low reactivity of these elements (most elements react with other elements to get to a lower, more stable energy state). For example the (Group 7A) have 7 valence electrons (s2p5) and want to gain one electron to get the low energy, stable electron configuration of the noble gases. The elements in group 6 (s2p4) want to gain 2 electrons to get the low energy, stable electron configuration of the noble gases. The Group 1A elements (s1) want to lose their outer electron to empty their outer shell and get a stable electron configuration. For example if (1s22s22p63s1) loses its 3s1 electron it will have filled s and p orbitals in its outer energy level.

Gilbert Lewis, in 1916, proposed the : react by changing their number of electrons so as to acquire the stable electron configuration of a (s2p6).

An exception to the octet rule is the electron configuration of . Helium(1s2) is a noble gas, only it has only one orbital, the s orbital. It is filled and therefore stable and elements close to it (, and sometimes ) try to acquire its electron configuration by losing or gaining electrons).

Covalent bonding is a type of bonding that occurs between . Nonmetals have many valence electrons and wish to gain electrons in order to fill their outer energy level to become stable (fulfill the octet rule - get the outer electron 2 6 configuration of a noble gas, s p ). For example fluorine (F2) is a diatomic consisting of two fluorine atoms covalently bonded together. Each atom is one electron shy of the noble gas electron configuration (valence electrons are s2p5) and so they each share one of each other's electrons to fill up their outer energy level. In a Bohr model it would look like this:

The electron dot structure would look like:

We can translate the above dot structure into a structural formula - a formula which represents the bonding pair of electrons as a line. All other valence electrons are ignored.

In a quantum mechanical representation it would look like this :

The outer electrons occupy the same orbital, outlined in red, giving each atom the s2p6 noble gas electron configuration. Note the electrons in the outer orbital (outlined in red) have opposite spins according to the Pauli exclusion principle.

2 2 Here is another example using dioxide (CO2). Carbon has four valence electrons (s p ) and needs four more (to get to s2p6) and each of the atoms has 6 valence electrons (s2p4) and needs to share two more. I will skip the Bohr model as it is not the most accurate representation of modern atomic theory.

Electron dot structure:

Structural formula:

Quantum mechanical model:

Note each model represents four bonding pairs of electrons and each model shows the outer energy level filled - having the s2p6 electron configuration with the shared electrons.

Another observation is that carbon is no longer in its ground state. Note in the above diagram of the quantum mechanical model that carbon has only one electron in its 2s orbital. In order to get four unpaired electrons (to bond with the two oxygen atom's four unpaired electrons) from carbon one of the 2s electrons moves into the 2p orbitals to form what are called sp3 orbitals - one s orbital combined with 3 p orbitals). This is called hybridization and is shown below.

• Single bonds are formed when one pair of electrons forms the bond (e.g fluorine - see above). • Double bonds are formed when two pairs of electrons form bonds between atoms (e.g. CO2 - see above). • Triple bonds are formed when three pairs of electrons form bonds between atoms. An example is (N2). A nitrogen atoms has the outer electron configuration of s2p3 and therefore needs three electrons and forms a triple bond when bonding with another nitrogen atom as shown below.

An exception to the octet rule is when atoms bond to attain the electron configuration of the noble gas helium (1s2). This is also a very stable, low energy electron configuration. An example of this is hydrogen (H2).

Polyatomic are groups of atoms that are covalently bonded together that have, as a group, gained or lost electrons to form an . The covalent bonds holding them together are strong so they usually act as a single unit and do not break apart 2- or change during chemical reactions. An example is the ion (SO4 ). It consists of one sulfur atom and 4 oxygen atoms that have gained two electrons and therefore have a 2- charge. The dot structure is shown below. Sulfur and oxygen are group 16 elements and have six valence electrons (s2p4). Sulfur's electrons are in black, oxygen's are in blue and the electrons gained from another source (probably from a metallic element) are in red. The whole polyatomic ion in shown in brackets with the charge indicated.

Resonance is the phenomenon ion which two or more equally valid formulas (electron dot or structural) can be drawn for a molecule. This can be seen in the ozone (O3) molecule below (I have chosen to show it as a linear molecule. The actual shape is bent triatomic).

When the total number of valence electrons of a substance is odd it is impossible to arrange the electrons in such a way as to fulfill the octet rule. For example, nitrogen dioxide (NO2) has a total of 17 valence electrons (5 from the nitrogen and six from each oxygen). If we draw the structural formula of one of the forms with the remaining valence electrons we find that there is always an :

The results of this are that an orbital will have only one electron with a certain . This spinning electron creates a causing the substance to be attracted to an external magnetic field. These substances are called paramagnetic and can be detected experimentally by determining their mass in and out of a magnetic field. A magnetic field can be placed so that it will pull down on a paramagnetic substance giving it a greater apparent mass than it has outside of the magnetic field.

Substances which have an even number of electrons have orbitals with a pair of electrons with opposite spins. These opposite spins create magnetic fields which cancel each other out leaving no net magnetic effect on the substance. These substances are called diamagnetic.

The VSEPR theory stands for Valence Shell Repulsion theory. This theory explains the shape of molecules based on the idea that electron pairs repel each other so that the electrons involved in bonds try to get as far apart as possible. For example, CO2 has two oxygen atoms double bonded to the central carbon. The two pairs of electrons involved in the bond all have the same charge (negative) and so try to move as far apart as possible, making it a linear molecule. It would be wrong to draw the bonds closer together than they would have to be, as illustrated below.

While bonding pairs of electrons repel each other, unshaired pairs have a greater effect on the shape of a molecule. Their charge is not diluted between the attraction for opposite nuclei and so repel bonded pairs to a great extent. If we look again at the ozone molecule we can see why its actual shape is bent. The unshared pair of electrons at the top of the molecule repels the electrons in the bonds forcing them downward into a bent shape.

The molecular shape of a compound is important because the shape of a compound's molecule determines many of the chemical and physical properties of the compound.

Covalent bonds can be either polar or nonpolar. In a nonpolar the electrons are evenly distributed between the two atoms sharing the electrons. This is because the of the two atoms are very similar. Diatomic substances (made up of two of the same atoms, e.g. F2, O2, N2 etc.) are nonpolar because atoms of the same element will have identical electronegativities. In a polar covalent bond one of the atoms has a greater and therefore the shared electrons are pulled toward that atom creating a polar bond - a bond with a negative and positive side. Polar bonds can be illustrated by using the sigma (+ or -) symbol or by an arrow starting with a + sign as shown below. In the water molecule oxygen is more electronegative so it attracts the electrons to itself making the oxygen side of the bond negative and leaving the hydrogen side of the bond positive.

A nonpolar molecule has an even distribution of its electrons making it nonpolar - no positive or negative side or pole. A polar molecule has one or more positive or negative poles.

A dipole is a molecule that has two poles, a positive pole and a negative pole. Dipoles are formed between atoms with differing electronegativities.

Van der Waals forces are intermolecular forces. They are forces between molecules and not within molecules. The three van der Waals forces are :

London dispersion forces or dispersion forces - weakest - thought to be created by the motion of electrons. The more electrons a molecule has the greater its attraction for other molecules. Dipole-dipole interaction - this is an electrostatic force is created by the attraction of opposite charges. The positive side of one molecule will attract and form a weak bond with the negative side of another molecule.

Hydrogen bond - strongest of the van der Waals forces - formed when a hydrogen atom involved in a polar bond (making it positively charged) of one molecule attracts an unshared electron pair of another molecule. See below.

The halogens (Group 17) are all diatomic molecules. As you go down the group (F2, Cl2, Br2, I2) the number of electrons increases and so does the dispersion force. This means the intermolecular forces become greater which explains why fluorine and are gases (weaker intermolecular forces) bromine is a liquid (stronger intermolecular forces) and iodine is a solid (strongest intermolecular forces of the halogens listed above).

Covalent Compounds

Chemical bond - A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. Attractive forces between the nuclei and electrons must outweigh the repulsive forces between the nuclei of each element and the electrons of each element.

Introduction to Chemical Bonding

I. Types of Chemical Bonding A. 1. Chemical bonding that results from the electrical attraction between cations and anions 2. Electrons are transferred in pure ionic bonding B. Covalent Bonding 1. Results from the sharing of electron pairs between two or more atoms a. Nonpolar Covalent Bond - A covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of charge b. Polar Covalent Bond - A covalent bond in which the bonded atoms have an unequal attraction for the shared electrons and a resulting unbalanced distribution of charge

Covalent Bonding and Molecular Compounds

I. Important Definitions A. Molecule 1. A neutral group of atoms that are held together by covalent bonds B. Diatomic Molecule 1. A molecule containing only two atoms (H2 N2 O2 F2 Cl2 Br2 I2) memorize these 7 diatomic molecules! C. Molecular Compound 1. A whose simplest units are molecules D. Chemical Formula 1. Indicates the relative numbers of atoms of each kind of a chemical compound by using atomic symbols and numerical subscripts E. Molecular Formula 1. Shows the types and numbers of atoms combined in a single molecule of a molecular compound

II. Formation of a Covalent Bond A. Interatomic Forces 1. As atoms approach one another, there are a number of forces at work a. repulsion of one nucleus for another b. repulsion of electrons for other electrons c. attraction of electrons to nuclei B. Energy Considerations 1. Atoms approach each other a. Potential energy decreases as attractive forces dominate over repulsive forces 2. Atoms get too close a. Potential energy begins to increase as repulsive forces dominate over attractive forces

III. Characteristics of the Covalent Bond A. Bond Length 1. The distance between two bonded atoms at their minimum potential energy B. Bond Energy 1. The energy required to break of and form neutral isolated atoms H2 + 436 kJ ‰ 2H 2. Covalent Bonding in Hydrogen

IV. The Octet Rule A. The Octet Rule 1. Chemical compounds tend to form so that each atom, by gaining, losing or sharing electrons, has an octet of electrons in its highest occupied energy level B. Diatomic Fluorine

C. Exceptions to the Octet Rule 1. Hydrogen a. Two valence electrons (helium configuration) 2. Expanded Octet a. More than eight electrons b. d orbitals involved as well as s and p c. atoms in the 3rd row and beyond that have d orbitals available to hold extra electrons d. likes to be electron deficient (only 6 electrons)

V. Electron-Dot Notation (Lewis dot or Lewis structures) A. Electron-dot Notation 1. An electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element's symbol 2. Inner shell electrons are not shown 3. The valence electrons are the electrons involved in the formation of covalent bonds

VI. Lewis Structures A. Unshared Electron Pairs (Lone Pairs) 1. A pair of electrons that is not involved in bonding B. Lewis Structures 1. Formulas in which atomic symbols represent nuclei and inner-shell electrons, dot pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons (lone pairs)

C. Structural Formula 1. Formulas indicating the kind, number, arrangement, and bonds but not unshared pairs of the atoms in a molecule

D. Drawing Lewis Structures (trichloromethane, CHCl3 as an example) 1. Determine the type and number of atoms in the molecule 2. Determine the total number of valence electrons in the compound C 1 x 4e- = 4e- H 1 x 1e- = 1e- Cl 3 x 7e- = 21e- Total = 26e- valence electrons to be used 3. Arrange the atoms to form a skeleton structure for the molecule. If carbon is present, it is the central atom. Otherwise, the first element in the formula or the least electronegative element atom is central (except for hydrogen, which is never central). Then connect the atoms by electron-pair bonds.

4. Add unshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is surrounded by eight electrons. (Obey the octet and the duet rules)

5. Count the electrons in the structure to be sure that the number of valence electrons used equals the number available and be sure to check that you have obeyed the octet and duet rules

VII. Multiple Covalent Bonds A. Double Bonds 1. A covalent bond produced by the sharing of two pairs of electrons (4 electrons) between two atoms

2. Higher bond energy and shorter bond length than single bonds B. Triple Bonds 1. A covalent bond produced by the sharing of three pairs of electrons (6 electrons) between two atoms

VIII. Resonance Structures A. Resonance 1. Resonance refers to bonding in molecules or ions that cannot be correctly represented by a single

IX. A Comparison of Ionic and Molecular Compounds

A. Polyatomic Ions 1. A charged group of covalently bonded atoms

Molecular Geometry

I. VSEPR (Valence Shell Electron Pair Repulsion) Theory A. Molecular Polarity 1. The uneven distribution of molecular charge 2. Molecules with preferential orientation in an electric field

B. VSEPR Theory 1. Repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible C. VSEPR and Unshared Electron Pairs 1. Unshared pairs take up positions in the geometry of molecules just as atoms do 2. Unshared pairs have a relatively greater effect on geometry than do atoms 3. Lone (unshared) electron pairs require more room than bonding pairs (they have greater repulsive forces) and tend to compress the angles between bonding pairs.

II. Intermolecular Forces (IMF) A. Intermolecular Forces 1. Forces of attraction between molecules a. Generally weaker than bonds that join atoms in molecules b. Boiling point gives a rough estimate of intermolecular forces (1) high bp = large attractive forces (2) low bp = small attractive forces

B. Molecular Polarity and Dipole-Dipole Forces 1. Dipole: represented by an arrow with a head pointing toward the negative pole and a crossed tail at the positive pole

2. Dipole-Dipole forces a. The negative region of one molecule is attracted to the positive region of another molecule b. A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons C. Hydrogen Bonding (very strong IMF) 1. The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative (N O F) atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule. H2O is bonded to other H2O molecules by hydrogen bonding.

D. London Dispersion Forces (van de Waals Forces) 1. The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles a. Electrons moving in an atom may create temporary unbalanced distribution of charge b. All molecules experience London forces c. London forces are the only forces of attraction among noble-gas atoms and nonpolar molecules 2. London forces increase with the number of electrons in an atom or molecule a. Greater mass = Greater London forces