<style> /* Width of text in slides */ .slides { width: 1200px !important; } /* line space bullet list */ .reveal .slides section, .reveal .slides section > section { line-height: 1.5; } /* headers */ .reveal h1 { margin-bottom: 1em; font-weight: bold; font-size: 32pt } .reveal h2 { margin-bottom: 1em; font-weight: bold; font-size: 30pt } .reveal h3 { margin-bottom: 1em; font-weight: bold; font-size: 28pt } .reveal h4 { margin-bottom: 1em; font-weight: bold; font-size: 26pt } .reveal h5 { margin-bottom: 1em; font-weight: bold; font-size: 24pt } .reveal h6 { margin-bottom: 1em; font-weight: normal; font-size: 24pt } /* text */ .reveal { font-size: 24pt; font-weight: normal; } /* logo as background */ .container { background: no-repeat url(https://www.mcc.edu/MottStrong/images/mott-strong-3c-logo.png) top 20px left 10px; background-size: 5% !important; } code {color: #000000} .reveal table th { font-size: 1.20em} .reveal table td { font-size: 1.00em} /* Remove Border arround images */ .reveal section img { background:none; border:none; box-shadow:none; object-fit: contain; margin: 0; padding: 0; max-width: 100%; } .reveal figure img { box-shadow:none; background:none; border:none; max-width: 100%; object-fit: contain; } </style> $\require{mhchem}$ $\require{cancel}$ OpenStax Chemistry 2e Chapter 7 Section 2 === Michael Stogsdill Mott Community College --- ## Learning Objectives - Describe the formation of covalent bonds - Define electronegativity and assess the polarity of covalent bonds --- ## Covalent Compounds - Actually exist as separate molecules - These electrically neutral molecules have weaker interactions than ionic compounds - They have low melting points and boiling points - Exist as solids, liquids, and gases at STP - Most covalent compounds are insoluble in water. - They are poor conductors of electricity and heat. --- <img src="https://doc.chemnotes.org/uploads/4038c9e6-4124-4a26-9ecd-176edf005cd8.png" alt="Volumetric flask filled with cartoon diatomic molecules in the gas phase." style="zoom:100%;" /> --- ## Covalent Bonding - **Covalent bonding** is the result mutual attraction of atoms for a *shared* pair of electrons. - *Covalent bonding* occurs between elements with fairly close *ionization energies* and *electron affinities*. - Bond Length is determined by the energetics of overlapping orbitals. --- ## Covalent Bond Length - When *covalent bonds* form, the outer orbitals of the atoms overlap *releasing* energy. - The attraction of a shared electrons is balanced with the repulsion of the two atoms’ nuclei. - The length of a covalent bond is the distance between nuclei that results in the *lowest potential energy*. --- <img src="https://openstax.org/apps/archive/20250226.165223/resources/685c9dc768cbf53712bef9f2da980a2ca09825ef" alt="A graph is shown with the x-axis labeled, “Internuclear distance ( p m )” while the y-axis is labeled, “Energy ( J ).” One value, “0,” is labeled midway up the y-axis and two values: “0” at the far left and “0.74” to the left, are labeled on the x-axis. The point “74” is labeled, “H bond H distance.” A line is graphed that begins near the top of the y-axis and to the far left on the x-axis and drops steeply to a point labeled, “negative 7.24 times 10 superscript negative 19 J” on the y-axis and 74 on the x-axis. This low point on the graph corresponds to a drawing of two spheres that overlap considerably. The line then rises to zero on the y-axis and levels out. The point where it almost reaches zero corresponds to two spheres that overlap slightly. The line at zero on the y-axis corresponds to two spheres that are far from one another." style="zoom:50%;" /> Note: **Figure 7.4** The potential energy of two separate hydrogen atoms (right) decreases as they approach each other, and the single electrons on each atom are shared to form a covalent bond. The bond length is the internuclear distance at which the lowest potential energy is achieved. --- ## Exothermic vs. Endothermic - We've seen that *exothermic reactions* release energy, while *endothermic reactions* absorb energy. - We can now tie the energy absorbed and released to the making (releases energy) and breaking (requires energy) of *chemical bonds*. --- ## Electronegativity - **Electronegativity** is a measure of the tendency of an atom to attract electrons (or electron density) towards itself. - It determines how the shared electrons are distributed between the two atoms in a bond. - The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity. ---  Note: **Figure 7.6** The electronegativity values derived by Pauling follow predictable periodic trends, with the higher electronegativities toward the upper right of the periodic table. --- ## Electronegativity vs. Electron Affinity - The **electron affinity** of an element is a measurable physical quantity. - The energy released or absorbed when an isolated gas-phase atom acquires an electron, measured in kJ/mol. - Electronegativity describes how tightly an atom attracts electrons in a bond. - It is a dimensionless quantity that is calculated, not measured. - The relative scale of electronegativity from 0 to 4 is arbitrary. --- ## Linus Pauling - Pauling derived the first electronegativity values by comparing the amounts of energy required to break different types of bonds. - He chose an arbitrary relative scale ranging from 0 to 4.  Note: **Figure 7.7** Linus Pauling (1901–1994) made many important contributions to the field of chemistry. He was also a prominent activist, publicizing issues related to health and nuclear weapons. --- ## Question 15 <p style="text-align:left">From its position in the periodic table, determine which atom in each pair is more electronegative:</p> <p style="text-align:left">a. Br or <span><!-- .element: class="fragment highlight-green" -->Cl</span></p> <p style="text-align:left">b. N or <span><!-- .element: class="fragment highlight-green" -->O</span></p> <p style="text-align:left">c. S or <span><!-- .element: class="fragment highlight-green" -->O</span></p> <p style="text-align:left">d. P or <span><!-- .element: class="fragment highlight-green" -->S</span></p> <p style="text-align:left">e. Si or <span><!-- .element: class="fragment highlight-green" -->N</span></p> <p style="text-align:left">f. Ba or <span><!-- .element: class="fragment highlight-green" -->P</span></p> <p style="text-align:left">g. N or <span><!-- .element: class="fragment highlight-green" -->K</span></p> Note: a. $\ce{Cl}$ b. $\ce{O}$ c. $\ce{O}$ d. $\ce{S}$ e. $\ce{N}$ f. $\ce{P}$ g. $\ce{K}$ --- ## Question 17 <p style="text-align:left">From their positions in the periodic table, arrange the atoms in each of the following series in order of increasing electronegativity:</p> <p style="text-align:left">a. C, F, H, N, O &emsp; &emsp; <span style="color:green"><!-- .element: class="fragment" data-fragment-index="1" -->H &lt; C &lt; N &lt; O &lt; F</span></p> <p style="text-align:left">b. Br, Cl, F, H, I &emsp; &emsp; <span style="color:green"><!-- .element: class="fragment" data-fragment-index="2" -->H &lt; I &lt; Br &lt; Cl &lt; F</span></p> <p style="text-align:left">c. F, H, O, P, S &emsp; &emsp; <span style="color:green"><!-- .element: class="fragment" data-fragment-index="3" -->H &lt; P &lt; S &lt; O &lt; F</span></p> <p style="text-align:left">d. Al, H, Na, O, P &emsp; &emsp; <span style="color:green"><!-- .element: class="fragment" data-fragment-index="4" -->Na &lt; Al &lt; H &lt; P &lt; O</span></p> <p style="text-align:left">e. Ba, H, N, O, As &emsp; &emsp; <span style="color:green"><!-- .element: class="fragment" data-fragment-index="5" -->Ba &lt; H &lt; As &lt; N &lt; O</span></p> --- ## Pure Covalent vs. Polar Covalent Bonds  Note: **Figure 7.8** As the electronegativity difference increases between two atoms, the bond becomes more ionic. --- ## Pure Covalent - We have been aware of diatomic elemental species - $\ce{H2}$, $\ce{N2}$, $\ce{O2}$, $\ce{F2}$, $\ce{Cl2}$, $\ce{Br2}$ - The electronegativity difference between these atoms is zero. Electrons are shared evenly between the atoms. --- ## Polar Covalent - When two different types of atoms are bonded together, a difference in electronegativity exists. - When this difference is below the threshold to form an *ionic bond*, a **polar covalent bond** is formed. - The electrons are not shared evenly. *Electron density* is drawn *towards the more electronegative atom* resulting in a partial negative charge ($\delta^{-}$) on the atom with greater electronegativity and a partial positive charge ($\delta^{+}$) on the atom with lower electronegativity - This unequal charge distribution gives rise to a **dipole**. --- <img alt="Two diagrams are shown and labeled “a” and “b.” Diagram a shows a small sphere labeled, “H” and a larger sphere labeled, “C l” that overlap slightly. Both spheres have a small dot in the center. Diagram b shows an H bonded to a C l with a single bond. A dipole and a positive sign are written above the H and a dipole and negative sign are written above the C l. An arrow points toward the C l with a plus sign on the end furthest from the arrow’s head near the H." src="https://openstax.org/apps/archive/20250226.165223/resources/a645a9b227fe84dfa7439d3b7ab7397c86f6420b" style="zoom:150%" > Note: **Figure 7.5** (a) The distribution of electron density in the HCl molecule is uneven. The electron density is greater around the chlorine nucleus. The small, black dots indicate the location of the hydrogen and chlorine nuclei in the molecule. (b) Symbols δ+ and δ– indicate the polarity of the H–Cl bond. --- ## Question 12 What information can you use to predict whether a bond between two atoms is covalent or ionic? ---- <p style="text-align:left;">If the elements in a compound are the same or close to one another in the periodic table, the compound is likely to be covalent; if they are far apart, then the compound will likely be ionic. If electronegativity values are available, the difference in electronegativity can indicate whether a bond is likely to be ionic or covalent.</p> --- ## Question 14 Explain the difference between a nonpolar covalent bond, a polar covalent bond, and an ionic bond. ---- <p style="text-align:left;">In an ionic bond, an electron or electrons are transferred from one atom to another, making one negative and the other positive. The electrostatic attraction holds them together. In a covalent bond, electrons are shared between the two atoms, resulting in a bond. A polar covalent bond is one in which the electrons are shared unequally. In this case, one atom has a greater electronegativity and a greater attraction for the electrons in the bond. In a pure covalent bond, the electrons are shared equally. In this case, the two atoms have very similar, if not the same, electronegativities and attract the electrons in the bond equally.</p> --- ## Bond Polarity and Electronegativity Difference | Bond | ΔEN | Polarity | | ---- | :--: | :-----------: | | C–H | 0.4 | ^δ−^ C−H ^δ+^ | | S–H | 0.4 | ^δ−^ S−H ^δ+^ | | C–N | 0.5 | ^δ+^ C−N ^δ−^ | | N–H | 0.9 | ^δ−^ N−H ^δ+^ | | C–O | 1.0 | ^δ+^ C−O ^δ−^ | | O–H | 1.4 | ^δ−^ O−H ^δ+^ | **Table 7.1** --- ## Question 21 <p style="text-align:left">Identify the more polar bond in each of the following pairs of bonds:</p> <p style="text-align:left">a. <span><!-- .element: class="fragment highlight-blue" -->HF</span> or HCl</p> <p style="text-align:left">b. NO or <span><!-- .element: class="fragment highlight-blue" -->CO</span></p> <p style="text-align:left">c. SH or <span><!-- .element: class="fragment highlight-blue" -->OH</span></p> <p style="text-align:left">d. <span><!-- .element: class="fragment highlight-blue" -->PCl</span> or SCl</p> <p style="text-align:left">e. CH or <span><!-- .element: class="fragment highlight-blue" -->NH</span></p> <p style="text-align:left">f. SO or <span><!-- .element: class="fragment highlight-blue" -->PO</span></p> <p style="text-align:left">g. <span><!-- .element: class="fragment highlight-blue" -->CN</span> or NN</p> Note: a. $\ce{HF}$ b. $\ce{CO}$ c. $\ce{OH}$ d. $\ce{PCl}$ e. $\ce{NH}$ f. $\ce{PO}$ g. $\ce{CN}$