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CHAPTER NOTES – CHAPTER 16 Covalent Bonding Goals : to Gain CHAPTER NOTES – CHAPTER 16 Covalent Bonding Goals : To gain an understanding of : 1. Valence electron and electron dot notation. 2. Stable electron configurations. 3. Covalent bonding. 4. Polarity of bonds and molecules NOTES: Valence electrons are the electrons in the highest energy level of an atom. For example, in the calcium atom (electron configuration 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 valence electron 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 halogens (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 sodium (1s22s22p63s1) loses its 3s1 electron it will have filled s and p orbitals in its outer energy level. Gilbert Lewis, in 1916, proposed the octet rule : Atoms react by changing their number of electrons so as to acquire the stable electron configuration of a noble gas (s2p6). An exception to the octet rule is the electron configuration of helium. 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 (lithium, beryllium and sometimes hydrogen) try to acquire its electron configuration by losing or gaining electrons). Covalent bonding is a type of bonding that occurs between nonmetals. 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 molecule 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 carbon dioxide (CO2). Carbon has four valence electrons (s p ) and needs four more (to get to s2p6) and each of the oxygen 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 nitrogen (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. 2 An exception to the octet rule is when atoms bond to attain the electron configuration of the noble gas helium (1s ). This is also a very stable, low energy electron configuration. An example of this is hydrogen (H2). Polyatomic ions are groups of atoms that are covalently bonded together that have, as a group, gained or lost electrons to form an ion. 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 sulfate 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 resonance forms with the remaining valence electrons we find that there is always an unpaired electron: The results of this are that an orbital will have only one electron with a certain spin. This spinning electron creates a magnetic field 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 Electron Pair 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 covalent bond the electrons are evenly distributed between the two atoms sharing the electrons. This is because the electronegativities 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 electronegativity 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.
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