Covalent Bonding Honors Chemistry PDF

Here is the transcription of the provided text, formatted into markdown. # Covalent Bonding Honors Chemistry Mrs. Calder ## Review of Matter * Some elements exist as atoms (monoatomic) such as Noble Gases * Metals & non-metals combine by transferring electrons to form salts IONIC COMPOUNDS *...

Here is the transcription of the provided text, formatted into markdown. # Covalent Bonding Honors Chemistry Mrs. Calder ## Review of Matter * Some elements exist as atoms (monoatomic) such as Noble Gases * Metals & non-metals combine by transferring electrons to form salts IONIC COMPOUNDS * Some elements do not give up electrons but have a tug-of-war over them and the sharing of electrons holds them together. These are covalent bonds. ## About Covalent Compounds A. The smallest form of them is called a molecule. (NOT FORMULA UNIT) B. Sometimes called "Molecular" compounds. The image displays atoms either sharing electrons or transferring electrons. *Sharing of Electrons* results in a **covalent bond** and formation of a **molecule.** *Transfer of Electrons* results in creating **positive and negative ions** and **ionic bond.** * Diatomic molecules consist of just 2 atoms of a single element (HOFBrINCI) * Molecular Compounds consist of atoms of different elements combined in a specific ratio ($H_2O$, $NH_3$) * Remember Ionic compounds were lattices * They have low melting points and boiling points compared to ionic compounds * They consist of 2 or more NON-METALS combined * Can be solids, liquids, or gases * Remember all ionic compounds were:??? * A molecular formula shows how many atoms of each element a molecule contains * UNLIKE IONIC COMPOUNDS, the molecular formula is not the lowest whole number ratio, but tells exactly how many atoms there are of each thing. Ex. Ethane $C_2H_6$ * Structural formulas show the arrangement of the atoms. The image contains three different ways to represent ammonia: 1. The molecular formula showing how many atoms of each element are present 2. The structural formula 3. The ball and stick model which shows the best 3-dimensional arrangement ## Covalent Bonds * They fulfill the octet rule by sharing electrons. See Chlorine example (next slide). * When only one pair of electrons is shared, it is called a SINGLE COVALENT BOND. * When drawing the structural formula, a pair of shared electrons is represented as a dash (-) A. Unshared electrons remain dots: $:\text{CI} \cdot \cdot \text{CI}: \rightarrow :\text{CI} : \text{CI}: \rightarrow : \text{CI} - \text{CI}:$ Another Example: Carbon joined with 4 hydrogens forming methane $CH_4$. ## Octet Rule * **Octet Rule**- a compound tends to form so that each atom has an octet (8) of electrons in its highest energy level by gaining, losing or sharing electrons * **Duet Rule**- applies to H and He ## Drawing Lewis Structures 1. find the number of valence electrons in each atom and add them up 2. draw the atoms next to each other in the way they will bond (if there is just one atom of an element make it the center) 3. add one bonding pair between each connected atoms (a dash = a pair) 4. add the rest of the electrons as dots until all have 8 *** ## Example 1 * $CH_3CI$ * methyl chloride * C: 4 x 1 = 4 * H: 1 x 3 = 3 * CI: 7 x 1 = 7 * total = 14 electrons * carbon is central ## Example 2 * $NH_3$ * ammonia * N: 5 x 1 = 5 * H: 1 x 3 = 3 * total = 8 * N is central $H - \underset{\cdot \cdot}{N} -H$ $|$ $H$ ## More Steps: * If you have no more electrons to add, and some are short (they don't all have octets), take any electrons not involved in a bond and create double bonds or triple bonds * (ONLY CARBON, OXYGEN, NITROGEN and a few others can do this) *** ## Example 3 * $N_2$ * nitrogen gas * N: 5 x 2 = 10 * 10 electrons * $\cdot \cdot N \equiv N \cdot \cdot$ ## Example 4 * $CH_2O$ * formaldehyde * C: 4 x 1 = 4 * H: 1 x 2 = 2 * O: 1 x 6 = 6 * total = 12 * C is central $H - \underset{||}C-H$ $\underset{\cdot \cdot}{O}$ ## Exceptions to Octet Rule * **Electron Deficient**: less than 8 * Boron: 3 in outer energy level * Beryllium: 2 in outer energy level * **Exceed Octet**: more than 8 * anything in $3^{rd}$ period or heavier * because may use the empty d orbital * add extras to middle atom * ex: S, P, I * Try $SF_6$, $SF_4$, $CIF_3$, $PCI_5$ H H : B : H **Note:** the image displays the Lewis structure for $SF_6$, $SF_4$, $CIF_3$, $PCI_5$. For polyatomic ions, put the charge outside brackets. * Example: Ammonium Ion $NH_4^+$ A) Boron Tetrafluoride Ion ($BF_4^-$) B) Sulfate ($SO_4^-$) C) Carbonate ($CO_3^{2-}$) **Note:** the image displays the Lewis structure for $NH_4^+$. ## BASIC RULES TO REMEMBER * C, N, O, F ALWAYS obey octet rule * B and Be tend to be electron deficient * $2^{nd}$ row elements can NOT exceed an octet * $3^{rd}$ row & up can exceed octet using d orbital but try to satisfy octet first. **NOT ON TEST** ## A Coordinate Covalent Bond... * When one atom donates both electrons in a covalent bond. * Carbon monoxide (CO) is a good example: * Both the carbon and oxygen give another single electron to share. **The images further detail coordinate covalent bonds**: The coordinate covalent bond is shown with an arrow and is commonly shown in carbon monoxide $:\text{C}::\bar{}\bar{}\bar{}::O:$ ## Resonance Structures * They are drawn when there are 2 or more valid Lewis structures * It allows us to envision the bonding. * In reality, no back and forth charges occur. The real bond is a hybrid of the 2 structures * Ex. $NO_2$ * There are resonance structures for the Ozone ion ($O_3$) and Acetate ion ($C_2H_3O_2^{-1}$) **Note:** *neither* structure is correct, it is actually a hybrid of the two. ## Resonance in Ozone Polyatomic ions - note the different positions of the double bond. Resonance in a carbonate ion ($CO_3^{2-}$): displayed with Lewis structures Resonance in an acetate ion ($C_2H3O_2^{-1}$): displayed with Lewis structures ## Bond Dissociation Energy * This is the energy to break a covalent bond. * The stronger the covalent bond, the higher the energy * It can explain why some covalent compounds are so stable (because energy is too high to break) ## VSEPR Theory (Valence shell electron pair repulsion) A. States that repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible. B. In the case below, $CO_2$ and $H_2O$ both have only 3 atoms, but the shape of water is bent. C. For VSEPR we are looking at the central atom only and determining the 3D shape based on the # of bonds and lone pairs (unshared e- pairs) around that atom. ## Linear * 2 Bonds, 0 unshared electrons (around central atom). The bond angle between the atoms attached to center is 180°. This way the bonds are as far away from each other as possible. * 2 pair electrons – 0 unshared (both in bonds) * Example: $BeH_2$ **Note:** Beryllium Hydride $H-Be-H$ and depiction of linear molecular geometry @180$^\circ$. ## Another example of linear * NOTE: Even though these are double bonds, we ignore the $2^{nd}$ or $3^{rd}$ in a double/triple. **Note:** Carbon Dioxide. * Look at center element and count the number of elements to which it is bonded. IGNORE the fact that it is a double bond for this! * Depiction of linear molecular geometry @180$^\circ$. * Ex: $H_2O$. 2 bonds, 2 lone pairs of electrons around central atom. **Bent** * The 2 lone electron pairs exert a little extra repulsion on the two bonding hydrogen atoms to create as slight compression to a ~104$\degree$ bond angle. * 4 pairs of total electrons – 2 bonds, 2 lone pairs. * ALSO **3 pairs with 1 unshared is called bent, but larger angle** **Note:** Depiction of Water molecule. * Example: $BH_3$. **Trigonal Planar (120$\degree$)** * Now there are 3 bonds around central atom. No lone pairs. * The hydrogen atoms are as far apart as possible at 120°. * The molecule all in a plane and is two dimensional. Flat. * 3 bonds no lone pairs **Note:** Boron Hydride H-B-H H Depiction of Trigonal Planar Molecular Geometry @120$^\circ$. * Example: $NH_3$. **Trigonal Pyramidal** * There are 3 bonds & 1 lone pair around central atom. It pushes down MORE so the other 3 are closer to each other. Creates a slight compression to a ~107$^\circ$ bond angle. * 4 pairs total, 3 bonds, 1 lone pair **Note:** Ammonia - $NH_3$ * Depiction of trigonal pyramidal molecular geometry - tetrahedral electron pair Geometry @107.3$^\circ$. * Example: $CH_4$. **Tetrahedral (109.5$\degree$)** Methane H H-C-H H * We have 4 bonds. No unshared pair. The hydrogen atoms are as far apart as possible at ~109° degree bond angle. * 4 bonds, no lone pairs **Note:** Depiction of Tetrahedral Molecular Geometry. * NOT ON TEST **Trigonal Bipyramidal** **Note:** Example $PCl_5$ Model of a trigonal bipyramidal. * A trigonal bipyramid * 5 bonds, no unshared **Note:** Model of a trigonal planar **T-Shaped NOT ON TEST** * T- Shaped molecular geometry has 3 bonds, and 2 are lone pairs * 5 pairs, total. 3 bonds; 2 lone pairs * Ex: $BrF_3$ **Other Shapes:** **Note:** See Saw * Square Pyrimidal **Polarity:** * It is an uneven sharing of electrons in a covalent compound BASED ON ELECTRONEGATIVITY DIFFERENCE. * It results in partial charges indicated by the symbol. * Polarity can determine the shape of the molecule as well as how it behaves * The table below is a guide to determine if bonds will be covalent, polar covalent, or ionic: | Electronegativity Difference | Most likely type of bond | Example | | :---------------------------: | :-----------------------: | :------: | | 0-0.4 | Non Polar Covalent | $H_2$ | | 0.5-1.9 | POLAR | $HCI/H_2O$ | | 2.0 + | Ionic | $NaCl$ | Technically, 1.7-2.0 is tricky.. If a metal is involved, its ionic. If 2 nonmetals, polar covalent **Bond Polarity** * Written as: $\delta^+$ H–Cl $\delta^-$ * the positive and minus signs (with the lower case delta) denote partial charges. * Can also be shown: H–Cl * the arrow points to the more electronegative atom. ## Table of Electronegativities 1 | | | | | | | | | | | | | | | | | | | | :---: | :---: | :-------------------------: | :---------: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | ## Bond Polarities | Bond | Diff in EN | Negative atom | Type of Bond | | :----: | :--------: | :-----------: | :------------------: | | H - H | 0.0 | N/A | pure covalent | | C - H | 0.4 | C | (weakly) polar covalent | | O - H | 1.4 | O | polar covalent | | H - F | 1.9 | F | polar covalent | | S - O | 1.0 | O | polar covalent | | C - O | 1.0 | O | polar covalent | | Al - C | 1. 0 | C | polar covalent | | Na - Cl | 2.1 | Cl | ionic | | Mg - O | 2.3 | O | ionic | | Mg - C | 1.3 | C | polar covalent | ## Polar bond vs. Polar Molecule: * Just because a BOND is polar, doesn't mean the whole molecule is polar. * Ex. CO2 is linear and even though C-O is a polar bond, they pull equally in opposite directions so the MOLECULE is NOT polar. * But H2O is bent with lone pairs and that structure causes it to be a polar MOLECULE. * You will learn more in AP Chem on this | Naming/formulas for Covalent | | | | | | :-------------------------: | :--------: | :------: | :--------: | :-----: | | Compounds | | | | | | Mono | 1 | Неха | 6 | | | Di | 2 | Hepta | 7 | | | Tri | 3 | Octa | 8 | | | Tetra | 4 | Nona | 9 | | | Penta | 5 | Deca | 10 | | ## Examples * CO = * Carbon Monoxide * CO2 = * Carbon Dioxide * $P_2O_3$ = * Diphosphorus Trioxide * $BF_3$ * Boron Trifluoride * $PF_5$ = * Phosphorus Pentafluoride ## Network Solids: * Special covalent compounds where all atoms are covalently bonded to each other so that breaking the solid requires breaking covalent bonds, not just intermolecular forces. * Examples = Diamonds, Silicon carbide **Note:** The different configurations of molecules in diamond and graphite showing the tetrahedral arrangement. Very weak attractions between layers of network bonds.