Chapter 8: Covalent Bonding

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1 Chapter 8: Covalent Bonding Section 8.1 Section 8.2 Section 8.3 Section 8.4 Section 8.5 The Covalent Bond Naming Molecules Molecular Structures Molecular Shapes Electronegativity and Polarity Review Vocabulary chemical bond: the force that holds two atoms together oxyanion: a polyatomic ion in which an element (usually a nonmetal) is bonded to one or more oxygen atoms ionic bond: the electrostatic force that holds oppositely charged particles together in an ionic compound atomic orbital: the region around an atom s nucleus that defines an electron s probable location electronegativity: the relative ability of an atom to attract electrons in a chemical bond New Vocabulary covalent bond molecule Lewis structure sigma bond pi bond endothermic reaction exothermic reaction oxyacid structural formula resonance coordinate covalent bond polar covalent bond 8.1 The Covalent Bond Why do atoms bond? The stability of an atom, ion or compound is related to its energy:. Metals and nonmetals gain stability by electrons (gaining or losing) to form ions that have stable noble-gas electron configurations. Another way atoms can gain stability is by valence electrons with other atoms, which also results in noble-gas electron configurations. Atoms in non-ionic compounds share. The chemical bond that results from sharing electrons is a. A molecule is formed when two or more atoms bond covalently. The majority of covalent bonds form between atoms of nonmetallic elements Diatomic molecules (H 2, N 2, F 2, O 2, I 2, Cl 2, Br 2 ) exist because the two-atom molecules are more stable than the individual atoms. The most stable arrangement of atoms exists at the point of maximum net attraction, where the atoms bond covalently and form a molecule.

2 Single Covalent Bonds When only one pair of electrons is shared, the result is a. The figure shows two hydrogen atoms forming a hydrogen molecule with a single covalent bond, resulting in an electron configuration like helium In a dots or a line are used to symbolize a single covalent bond. The halogens the group 17 elements have 7 valence electrons and form single covalent bonds with atoms of other non-metals. Atoms in group 16 can share two electrons and form two covalent bonds. Water is formed from one oxygen with two hydrogen atoms covalently bonded to it. Atoms in group 15 form single covalent bonds, such as in ammonia. Atoms of group 14 elements form single covalent bonds, such as in methane. Sigma bonds are single covalent bonds. Sigma bonds occur when the pair of shared electrons is in an area centered between the two atoms. Multiple Covalent Bonds Double bonds form when two pairs of electrons are shared between two atoms.

3 Triple bonds form when three pairs of electrons are shared between two atoms A multiple covalent bond consists of one and at least one The pi bond is formed when parallel orbitals overlap and share electrons. The pi bond occupies the space above and below the line that represents where the two atoms are joined together. The Strength of Covalent Bonds The strength depends on the distance between the two nuclei, or bond length. As length increases, strength decreases The amount of energy required to break a bond is called the bond. The the bond length, the the energy required to break it. An is one where a greater amount of energy is required to break a bond in reactants than is released when the new bonds form in the products. An is one where more energy is released than is required to break the bonds in the initial reactants. Section 8.2 Naming Molecules Naming Binary Molecular Compounds Ex. N2O The first element is always named first using the entire element name, N is the symbol for nitrogen.

4 The second element is named using its root and adding the suffix -ide, O is the symbol for oxygen so the second word is oxide. Prefixes are used to indicate the number of atoms of each element that are present in the compound, There are two atoms of nitrogen and one atom of oxygen so the first word is dinitrogen and the second word is monoixide. Prefixes are used to indicate the number of atoms of each element in a compound. Many compounds were discovered and given common names long before the present naming system was developed (water, ammonia, hydrazine, nitric oxide). Naming Acids Binary Acids (An acid that contains hydrogen and one other element) Ex. HCl The first word has the prefix hydro- to name the hydrogen part of the compound. The rest of the word consists of a form of the root of the second element plus the suffix ic, HCl (hydrogen and chlorine) becomes hydrochloric. The second word is always acid, Thus, HCl in a water solution is called hydrochloric acid. An oxyacid is an acid that contains both a hydrogen atom and an oxyanion. Ex. HNO3 Identify the oxyanion present. The first word of an oxyacid s name consists of the root of the oxyanion and the prefix per- or hypo- if it is part of the name and a suffix. If the oxyanion s name ends with the suffix ate, replace it with the suffix ic. If the name of the oxyanion ends with suffix ite, replace it with suffix ous, NO3 the nitrate ion, becomes nitric. The second word of the name is always acid, HNO3 (hydrogen and nitrogen ion) becomes nitric acid. An acid, whether a binary acid or an oxyacid, can have a common name in addition to its compound name.

5 The name of a molecular compound reveals its composition and is important in communicating the nature of the compound. 8.3 Molecular Structures A structural formula uses to show relative positions of atoms. Drawing Lewis Structures Predict the location of certain atoms, the atom that has the least attraction for shared electrons will be the central atom in the molecule (usually, the one closer to the left side of the periodic table). All other atoms become terminal atoms. Note: Hydrogen is always a terminal atom. Determine the number of electrons available for bonding, the number of valence electrons. Determine the number of bonding pairs, divide the number of electrons available for bonding by two. Place the bonding pairs, place a single bond between the central atoms and each of the terminal atoms.

6 Determine the number of bonding pairs remaining, Subtract the number of bonding pairs in step 4 from the number of bonding pairs in step 3. Place lone pairs around terminal atoms, except hydrogen, to satisfy the octet rule. Any remaining pairs will be assigned to the central atom. Determine whether the central atom satisfies the octet rule, If not, convert one or two of the lone pairs on the terminal atoms into a double bond or a triple bond between the terminal atom and the central atom. Remember: carbon, nitrogen, oxygen and sulfur often form double and triple bonds. Atoms within a polyatomic ion are covalently bonded. The procedure for drawing Lewis structures is similar to drawing them for covalent compounds. Difference is, you need to determine the number of electrons available for bonding, find the number of electrons available in the atoms present and then subtract the ion charge if the ion is positive or add the ion charge if the ion is negative Resonance Structures Resonance is a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion. This figure shows three correct ways to draw the structure for (NO3)-1. Two or more correct Lewis structures that represent a single ion or molecule are resonance structures. The molecule behaves as though it has only one structure. The bond lengths are identical to each other and intermediate between single and double covalent bonds. Exceptions to the Octet Rule Some molecules do not obey the octet rule. A small group of molecules might have an odd number of valence electrons. NO2 has five valence electrons from nitrogen and 12 from oxygen and cannot form an exact number of electron pairs. A few compounds form stable configurations with less than 8 electrons around the atom a suboctet. A coordinate covalent bond forms when one atom donates both of the electrons to be shared with an atom or ion that needs two electrons.

7 A third group of compounds has central atoms with more than eight valence electrons, called an expanded octet. Elements in period 3 or higher have a d-orbital and can form more than four covalent bonds. 8.4 Molecular Shapes The shape of a molecule determines many of its. Molecular geometry (shape) can be determined with the Valence Shell Electron Pair Repulsion model, or which minimizes the repulsion of shared and unshared atoms around the central atom. Electron pairs repel each other and cause molecules to be in fixed positions relative to each other. Unshared electron pairs also determine the shape of a molecule. Electron pairs are located in a molecule. Hybridization Hybridization is a process in which atomic orbitals, identical hybrid orbitals. Carbon often undergoes hybridization, which forms an sp3 orbital formed from one s orbital and three p orbitals. Lone pairs also occupy hybrid orbitals. Single, double, and triple bonds occupy only (CO2 with two double bonds forms an sp hybrid orbital).

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9 8.5 Electronegativity and Polarity Electron Affinity, Electronegativity, and Bond Character Electron affinity measures the tendency of an atom to. Noble gases are not listed because they generally do not form compounds. This table lists the character and type of chemical bond that forms with differences in electronegativity. Unequal sharing of electrons results in a polar covalent bond Bonding is often not clearly ionic or covalent This graph summarizes the range of chemical bonds between two atoms.

10 Polar Covalent Bonds Polar covalent bonds form when atoms pull on electrons in a molecule unequally. Electrons spend more time around one atom than another resulting in partial charges at the ends of the bond called a. Covalently bonded molecules are either. Non-polar molecules are not attracted by an electric field. Polar molecules align with an electric field. Compare water, H2O, and CCl4 Both bonds are polar. The molecular shapes, determined by VSEPR, is bent and tetrahedral, respectively. O H bonds are asymmetric in water, so has a definite end and definite end. Thus, polar. The C Cl bonds are symmetrical in CCl4. The electric charge measured at any distance from the center is identical on all sides and partial charges are balanced. Thus nonpolar. Note: If bonds are polar, asymmetrical molecules are polar and symmetrical molecules are nonpolar. is the property of a substance s ability to dissolve in another substance. Polar molecules and ionic substances are usually soluble in. Non-polar molecules dissolve only in.

11 Properties of Covalent Compounds Covalent bonds between atoms are, but attraction forces between molecules are. The weak attraction forces are known as. The forces vary in strength but are weaker than the bonds in a molecule or ions in an ionic compound. Non-polar molecules exhibit a weak dispersion force, or induced dipole. The force between two oppositely charged ends of two polar molecules is a. A is an especially strong dipole-dipole force between a hydrogen end of one dipole and a fluorine, oxygen, or nitrogen atom on another dipole. Many physical properties are due to intermolecular forces. Weak forces result in the relatively low melting and boiling points of molecular substances. Many covalent molecules are relatively solids. Molecules can align in a, similar to ionic solids but with less attraction between particles. Solids composed of only atoms interconnected by a network of covalent bonds are called. are two common examples of network solids.