Science Applied Department (ASD) Centre for Foundation Studies and Extension Education (FOSEE)
PCE 0015 Chemistry for Engineers Foundation in Engineering ONLINE NOTES
Chapter 4
CHEMICAL BONDING
FOSEE , MULTIMEDIA UNIVERSITY (436821-T) MELAKA CAMPUS, JALAN AYER KEROH LAMA, 75450 MELAKA, MALAYSIA. Tel [phone] Fax [phone] URL: http://fosee.mmu.edu.my/~asd/
PCE0015
Chemistry For Engineers
Chapter 4
4.1 4.2 4.3 4.4 4.5 4.6
Lewis Dot Symbols The Ionic Bond The Covalent Bond Writing Lewis Structures Exception to the Octet Rule The hydrogen Bond
Upon completion of this chapter, you should be able to : 1. Apply Lewis dot symbols. A Lewis dot symbol shows the number of valence electrons possesses by an atom of a given element. They are useful mainly for representative elements. 2. Explain the ionic bond. An ionic bond is the product of the electrostatic forces of attraction between positive and negative ions. An ionic compound consists of a large network of ions in which positive and negative charges are balanced. 3. Explain the covalent bond. In a covalent bond, two electrons (one pair) are shared by two atoms. In multiple covalent bond, two or three pairs of electrons are shared by two atoms. 4. Write Lewis structure. The arrangement of bonding electrons and lone pairs around atoms in a molecule is represented by a Lewis structure. 5. Understand and apply the octet rule. The octet rule predicts that atoms form enough covalent bonds to surround themselves with eight electrons each. There are exceptions to the octet rule. 6. Explain the hydrogen bond. Hydrogen bond is a special type of dipole-dipole interaction (attraction between polar molecule) between the hydrogen atom in a polar bond, such as N H, O H, or F H, and an electronegative O, N, or F atom.
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PCE0015
Chemistry For Engineers
Chapter 4
4.1
LEWIS DOT SYMBOLS
Lewis dot symbol consists of the symbol of an element and one dot for each valence electron in an atom of the element. Figure below shows the Lewis dot symbols for the representative elements and the noble gases. Note that except for helium, the number of valence electrons each has is the same as the group number of the element. Elements in the same group have similar outer electron configurations and hence similar Lewis dot symbols.
4.2
THE IONIC BOND
An ionic bond is the electrostatic force that holds ions together in an ionic compound. For example, the reaction between lithium and fluorine to form lithium fluoride. The electron configuration of lithium is 1s22s1, and that of fluorine is 1s22s22p5. When lithium and fluorine atoms come in contact with each other, the outer 2s1 valence electron of lithium is transferred to the fluorine atom. Using Lewis dot symbols, the reaction represented as: Li + F 1s22s1 1s22s22p5 Li+ F (or LiF)
1s2 1s22s22p6
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Chemistry For Engineers
Chapter 4
The ionic bond in LiF is the electrostatic attraction between the positively charged lithium ion and the negatively charge fluoride ion. The compound itself is electrically neutral.
4.3
THE COVALENT BOND
Covalent bond is a bond in which two electrons are shared by two atoms . Covalent compounds are compounds that contain only covalent bonds. For the sake of simplicity, the shared pair of electrons is often represented by a single line. Thus, the covalent bond in the hydrogen molecule can be written as H H. Covalent bonding between many-electron atoms involves only the valence electrons. Example, for fluorine molecule, F2, the electron configuration for F is 1s12s2s2p5. The 1s electrons are low in energy and stay near the nucleus most of the time and they do not participate in bond formation. Thus each F atom has seven valence electrons, the formation of F2 molecule can be represented as follows: F + F FF or F F
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PCE0015
Chemistry For Engineers
Chapter 4
Note that only two valence electrons participate in the formation of F2. The other, nonbonding electrons are called lone pair pairs of valence electrons that are not involved in covalent bond formation. Thus each F in F2 has three lone pairs of electron. The structure used to represent covalent compound such as F2 is called Lewis structure. A Lewis structure is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms. Only valence electrons are shown in a Lewis structure. In the F2 molecules, the F atom achieves the stable noble gas configuration by sharing electrons : FF 8e 8e The formation of these molecules illustrates the octet rule : an atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons . By sharing electrons in a covalent bond, the individual atoms can complete their octets. Atoms can have different types of covalent bonds. In a single bond, two atoms are held together by one electron pair. If two atoms share two pairs of electrons, the covalent bond is called a double bond. Double bonds are found in molecules of carbon dioxide (CO2) and ethylene (C2H4). A triple bond arises when two atoms share three pairs of electrons, as in the nitrogen molecule (N2). Comparison of the Properties of Covalent and Ionic Compounds. Ionic and covalent compounds differ markedly in their general properties, physical for example appearance, melting point, boiling point, density, etc because of differences in the nature of their bonds. The intermolecular force that hold molecules of a covalent compound are not held together tightly. Where as the electrostatic forces holding ions together in an ionic compound are usually very strong, so ionic compound are solids at room temperature and have high melting point. Many ionic compound are soluble in water, and the aqueous solutions conduct electricity, because the compounds are strong electrolytes. Most covalent compounds are insoluble in water, or if they do dissolve, their aqueous solutions generally do not conduct electricity. Molten ionic compounds conduct electricity because they contain mobile cations and anions; liquid or molten covalent compounds do not conduct electricity because no ions are present.
4.4 WRITING LEWIS STRUCTURES
The basic steps are as follows: 1. Write the skeletal structure of the compound, using chemical symbols and placing bonded atoms next to another. In general, the least electronegative atom occupies
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PCE0015
Chemistry For Engineers
Chapter 4
the central position. Hydrogen and fluorine usually occupy the end positions in the Lewis structure. 2. Count the total number of valence electrons present. 3. Draw a single covalent bond between the central atom and each of the surrounding atoms. Complete the octets of the atoms bonded to the central atom. 4. After completing step 1 3, if the central atom has fewer than eight electrons, try adding double or triple bonds between the surrounding atoms and the central atom, using the lone pairs from the surrounding atoms to complete the octet of the central atom.
4.5 EXCEPTION TO THE OCTET RULE
The octet rule applies mainly to the second-period elements. Exceptions to the octet rule fall into three categories characterized by an incomplete octet, an odd number of electrons, or more than eight valence electrons around the central atom. The Incomplete Octet. In some compounds the number of electrons surrounding the central atom in a stable molecule is fewer than eight. For example beryllium, the electron
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PCE0015
Chemistry For Engineers
Chapter 4
configuration of it is 1s22s2; it has two valence electrons in the 2s orbital. In the gas phase, beryllium hydride (BeH2) exists as discrete molecules. The Lewis structure of BeH2 is: H Be H Only four electrons surround the Be atom, and there is no way to satisfy the octet rule for beryllium in this molecule. Odd-Electron Molecules. Some molecules contain an odd number of electrons. Eg. Nitric oxide (NO). NO As we need an even number of electrons for complete pairing (to reach eight), the octet rule clearly cannot be satisfied for all the atoms in any of these molecules. Odd-electron molecules are sometimes called radicals. Many radicals are highly reactive. The reason is that there is a tendency for the unpaired electron to form a covalent bond with an unpaired electron on another molecule. The Expanded Octet. Atoms of elements in and beyond the third period of the periodic table form some compounds in which more than eight electrons surround the central atom. In addition to the 3s and 3p orbitals, elements in the third period also have 3d orbitals that can be used in bonding. These orbitals allow an atom to form an expanded octet. For example : sulfur hexafluoride, a very stable compound. The electron configuration of sulfur is [Ne]3s23p4. In SF6, each of sulfurs six valence electrons forms a covalent bond with a fluorine atom, so there are twelve electrons around the central sulfur atom.
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PCE0015
Chemistry For Engineers
Chapter 4
4.6
THE HYDROGEN BOND
Normally, the boiling point of a series of similar compound containing elements in the same periodic group increase with increasing molar mass. This increase in boiling point is due to the increase in dispersion forces for molecules with more electrons. Hydrogen compounds of group 4A follow this trend (refer to figure below).
The lightest compound, CH4 has the lowest boiling point, and the heaviest compound, SnH4, has the highest boiling point. However, hydrogen compounds of the elements in Groups 5A, 6A, and 7A do not follow this trend. In each of these series, the lightest compound (NH3, H2O, and HF) has the highest boiling point, contrary to the expectation based on molar mass. This is because, there are stronger intermolecular attraction in NH3, H2O, and HF, called the hydrogen bond, which is a special type of dipole-dipole interaction (attraction between polar molecule) between the hydrogen atom in a polar bond, such as N H, O H, or F H, and an electronegative O, N, or F atom. The interaction is written : A H B or A H A
A and B represent O, N, or F; A H is one molecule or part of a molecule and B is a part of another molecule; and the dotted line represents the hydrogen bond. Note that the O, N, and F atoms all possess at least one lone pair that can interact with the hydrogen atom in hydrogen bonding. The average energy of a hydrogen bond is quite large for a dipole-dipole interaction (up to 40 kJ/mol). Thus, hydrogen bonds have a powerful effect on the structures and properties of many compounds.
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