Unit 4 Molecular Compounds PDF

**Unit 4 Molecular Compounds** **Key Terms** covalent bond molecule atomic element molecular element diatomic element ionic compound molecular compound Lewis structure bonding pair lone pair single, double, triple bond nonpolar covalent bond polar covalent bond electronegativity elec...

**Unit 4 Molecular Compounds** **Key Terms** covalent bond molecule atomic element molecular element diatomic element ionic compound molecular compound Lewis structure bonding pair lone pair single, double, triple bond nonpolar covalent bond polar covalent bond electronegativity electron hog bond dipole VSEPR theory electron-group geometry molecular shape linear trigonal planar tetrahedral pyramidal bent polar molecule nonpolar molecule ionic bonding dipole--dipole attraction hydrogen bond dispersion forces Problem Checklist 1. Describe the formation of a covalent bond and compare to an ionic bond. 2. Distinguish between atomic element, molecular element, ionic compound, and molecular compound. 3. Name and write formulas for molecular elements and molecular compounds: know the 7 diatomic elements and their names, and use Greek prefixes to name molecular compounds. 4. Given a chemical formula, distinguish as ionic or covalent and write the name. Given a chemical name, distinguish as ionic or covalent and write the formula. 5. Draw the Lewis structure for a molecule given the formula and count lone pairs and bonding pairs. 6. Describe the difference between a nonpolar covalent bond and polar covalent bond. 7. Draw bond dipoles for polar covalent bonds. 8. Use electronegativity values to predict bond type (nonpolar covalent, polar covalent, ionic). 9. Predict the shape of a molecule given the chemical formula. 10. Predict the polarity of a molecule given the chemical formula. 11. Predict the attractive forces in a substance given the chemical formula. Textbook Problems +-----------------------------------+-----------------------------------+ | Textbook Section | Page \# Questions | +===================================+===================================+ | 6.5 Molecular Compounds: Sharing | p.193 6.47, 6.49, 6.51, 6.53, | | Electrons | 6.55 | | | | | Understanding the Concepts | p.215 6.101 | | | | | Additional Practice Problems | p.216 6.121, 6.123, 6.125, 6.127, | | | 6.155 | +-----------------------------------+-----------------------------------+ | 6.6 Lewis Structures for | p.198 6.57(not d), 6.59 (not d), | | Molecules | 6.61 | | | | | Understanding the Concepts | p.215 6.105 | | | | | Additional Practice Problems | p.216 6.129, 6.131(not d), 6.157 | +-----------------------------------+-----------------------------------+ | 6.7 Electronegativity and Bond | p.202 6.65, 6.67, 6.69, 6.71, | | Polarity | 6.73 | | | | | Understanding the Concepts | p.215 6.109 | | | | | Additional Practice Problems | p.216 6.133, 6.135, 6.137, 6.139 | +-----------------------------------+-----------------------------------+ | 6.8 Shapes and Polarity of | p.207 6.75, 6.79, 6.81, 6.85 | | Molecules | | | | p.216 6.141, 6.143 (bc only), | | Additional Practice Problems | 6.145, 6.147 | +-----------------------------------+-----------------------------------+ | 6.9 Intermolecular Forces in | p.209 6.87, 6.89(not b), 6.91, | | Compounds | 6.93, 6.95 | | | | | Understanding the Concepts | p.215 6.111(not b) | | | | | Additional Practice Problems | p.217 6.149 | | | | | Challenge PRoblems | p.217 6.159 (not b) | +-----------------------------------+-----------------------------------+ | Chapter 6 Answers | p.218 | +-----------------------------------+-----------------------------------+ Contents {#contents.TOCHeading} ======== [6.5 Molecular Compounds: Sharing Electrons 2](#_Toc147480200) [Covalent Bonds 3](#covalent-bonds) [Names and Formulas of Molecular Compounds 3](#names-and-formulas-of-molecular-compounds) [Ionic or Covalent? 5](#ionic-or-covalent) [6.6 Lewis Structures for Molecules 6](#lewis-structures-for-molecules) [Lewis Structures, Molecules 6](#lewis-structures-molecules) [Diatomic Elements 7](#diatomic-elements) [Sharing Electrons Between Atoms of Different Elements 8](#sharing-electrons-between-atoms-of-different-elements) [Drawing Lewis Structures 8](#drawing-lewis-structures) [Double and Triple Bonds 9](#double-and-triple-bonds) [6.7 Electronegativity and Bond Polarity 10](#electronegativity-and-bond-polarity) [Electronegativity and Bond Types 10](#electronegativity-and-bond-types) [Dipoles and Bond Polarity 11](#dipoles-and-bond-polarity) [Summary: Variations in Bonding 13](#summary-variations-in-bonding) [Predicting Bond Type Using Electronegativity 13](#predicting-bond-type-using-electronegativity) [6.8 Shapes and Polarity of Molecules 14](#shapes-and-polarity-of-molecules) [Central Atoms with Two Electron Groups 15](#central-atoms-with-two-electron-groups) [Central Atoms with Three Electron Groups 15](#central-atoms-with-three-electron-groups) [Central Atoms with Four Electron Groups 16](#central-atoms-with-four-electron-groups) [Polarity of Molecules 19](#polarity-of-molecules) [Nonpolar molecules 19](#nonpolar-molecules) [Polar molecules 19](#polar-molecules) [6.9 Intermolecular Forces in Compounds 20](#intermolecular-forces-in-compounds) [Dipole--Dipole and Hydrogen Bonds 20](#dipoledipole-and-hydrogen-bonds) [Dispersion Forces 21](#dispersion-forces) [Intermolecular Forces and Melting Points 22](#intermolecular-forces-and-melting-points) []{#_Toc147480200.anchor} 6.5 Molecular Compounds: Sharing Electrons ========================================== Unlike ionic, molecular compounds need prefixes to indicate chemical formula. Covalent Bonds -------------- - **Molecular compounds** form when - atoms of two or more nonmetals **share electrons** and form a **covalent bond**. - valence electrons are **shared by nonmetal** atoms to achieve stability. - A molecule is a discrete group of atoms in a definite proportion. - Ex. Water H~2~O, Carbon dioxide CO~2~, Glucose C~6~H~12~O~6~ Names and Formulas of Molecular Compounds -------------------------------------------------------------- - first nonmetal given its element name - second nonmetal uses the first syllable of the name followed by *ide*. - When a subscript indicates two or more atoms of an element, a prefix is shown in front of its name. - Several compounds may be formed from the same two nonmetals: - The number of oxygen atoms present is indicated by the prefix. - When the vowels *o* and *o* or *a* and *o* appear together, the first vowel is omitted. **Learning Check** Name the molecular compound B~2~O~3~. Write the molecular formula for diphosphorus pentoxide. Select the correct name for each compound. 1\. SiCl~4~ A. silicon chloride B. tetrasilicon chloride C. silicon tetrachloride 2\. P~2~O~5~ A. phosphorus oxide B. phosphorus pentoxide C. diphosphorus pentoxide 3\. Cl~2~O~7~ A dichlorine heptoxide B. dichlorine oxide C. chlorine heptoxide Ionic or Covalent? ------------------ A compound is usually - **K**~2~O: **K** is a **metal**; the compound is ionic. potassium oxide - **N**~2~O: **N** is a **nonmetal**; the compound is covalent. dinitrogen oxide **Create a flow chart for naming ion vs covalent compounds** 6.6 Lewis Structures for Molecules ======================================================== - A molecule is represented by a Lewis structure: - valence electrons are arranged to give octets. **The Simplest Molecule, H~2~** - A hydrogen molecule, H~2~, - the positive charge of the nucleus attracts the electron of the other atom. - has a shared pair of electrons, *a covalent bond*, to give a noble gas arrangement of He to each H atom. - the molecule formed is more stable than the two individual H atoms. Lewis Structures, Molecules --------------------------- - A molecule is represented by a Lewis structure in which the valence electrons of all the atoms are arranged to give octets. - The shared electrons, or *bonding pairs*, are shown as two dots or a single line between atoms. - The nonbonding pairs, or *lone pairs*, are placed on the outside of the atoms.  - start with the electron-dot symbols for each fluorine atom. - each fluorine atom shares one electron to form a covalent bond, giving each fluorine an octet. Diatomic Elements -----------------  Sharing Electrons Between Atoms of Different Elements ----------------------------------------------------- Drawing Lewis Structures ------------------------ - To draw the electron-dot formula for methane, CH~4~, - draw the Lewis symbols for carbon and hydrogen. - carbon needs four hydrogen atoms to have an octet. - two bonding electrons between carbon and hydrogen can be drawn as a line. - the central atom is the atom in the formula with the least number of atoms.. Double and Triple Bonds ----------------------- - atoms share two pairs of electrons. - forms when there are not enough electrons to complete octets. - occurs when atoms share three pairs of electrons. - forms when there are not enough electrons to complete octets.  **Exceptions to the Octet Rule** - While the octet rule is useful, there are exceptions. - Hydrogen requires just two electrons to form a noble gas arrangement. - Nonmetals P, S, Cl, Br, and I can form compounds with 10 or 12 valence electrons. - The S atom has an octet in many compounds, but in SF~6~, there are 12 valence electrons or 6 bonds to the sulfur atom. 6.7 Electronegativity and Bond Polarity ======================================= Electronegativity is the pull an atom has on electrons.  - In H~2~ molecule, electrons are shared equally, and the bond is nonpolar. - In HCl molecule, electrons are shared unequally, and the bond is polar. Electronegativity and Bond Types -------------------------------- H-Si 2.1 − 1.8 = **0.3** Nonpolar covalent **polar covalent** bond occurs between nonmetal atoms. - unequal sharing of electrons. - moderate electronegativity difference (from 0.5 to 1.8). - becomes more polar as the difference in electronegativity increases Examples Electronegativity Atoms Difference Type of Bond O-Cl 3.5 − 3.0 = 0.5 Polar covalent Cl-C 3.0 − 2.5 = 0.5 Polar covalent O-S 3.5 − 2.5 = 1.0 Polar covalent Dipoles and Bond Polarity ------------------------- - The separation of charges in a polar bond is called a **dipole**. - The positive and negative ends of the dipole are located by using - the lowercase Greek letter delta with a positive or a negative charge. - an arrow that points from the positive to the negative end of the dipole.  - occurs between metal and nonmetal ions. - is a result of electron transfer. - has a large electronegativity difference (greater than 1.8). Summary: Variations in Bonding ------------------------------ - If the electronegativity difference is - between 0 and 0.4, the bond is nonpolar covalent. - between 0.5 and 1.8, the bond is polar covalent. - greater than 1.8, the bond is ionic, and the electrons are considered transferred. - between 1.8 and 2.0, can be either polar covalent or ionic. It depends on the elements involved in the bond: - Metal & non-metal = ionic - 2 non-metals = covalent Predicting Bond Type Using Electronegativity -------------------------------------------- **Learning Check** - Use the electronegativity difference to identify the type of bond (nonpolar covalent \[NP\], polar covalent \[P\], or ionic \[I\]) between the following: - A. K---N - B. N---O - C. Cl---Cl - D. H---Cl 6.8 Shapes and Polarity of Molecules ==================================== **Valence Shell Electron-Pair Repulsion Theory** (VSEPR) - describes the orientation of electron groups around the central atom. - states that electron groups are arranged as far apart as possible around the central atom. - states that the specific shape of a molecule is determined by the number of atoms attached to the central atom. **Example: Central atom bonded to Four Atoms** In a molecule of methane, CH4, the central C atom is bonded to four H atoms. The best geometry for minimal repulsion is tetrahedral, with bond angles of 109°. **\ ** Central Atoms with Two Electron Groups -------------------------------------- - two bonding electron groups are placed around the central atom, carbon. - the repulsion is minimized by placing the two groups on opposite sides of the carbon atom, giving this a **linear** arrangement with bond angles of 180°. - the shape with two electron groups around the central atom\ is **linear**.  Central Atoms with Three Electron Groups ---------------------------------------- - three electron groups are placed around the central atom, carbon. - the repulsion is minimized by placing the three groups as far apart as possible at bond angles of 120°. - the shape with three electron groups around the central atom is **trigonal planar**. - 1\. a single-bonded O atom - 2\. a double-bonded O atom - 3\. a lone pair of electrons - The repulsion is minimized by placing the three electron groups as far apart as possible. - The shape with two bonds and a lone pair on the central atom is **bent** with a bond angle of 120°.  Central Atoms with Four Electron Groups --------------------------------------- - four electron groups are attached to H atoms around the central atom, carbon. - the repulsion is minimized by placing the four electron groups at corners of a tetrahedron with bond angles of 109°. - the shape with four bonds on the central atom is called **tetrahedral**. - four electron groups, three bonds to H atoms and one lone pair, are around the central atom, N. - the repulsion is minimized by placing the four electron groups at corners of a tetrahedron with bond angles of 109°. - the shape with three bonds and a lone pair on the central atom is called **trigonal pyramidal**.  - four electron groups, two bonds to H atoms and two lone pairs, are around the central atom, O. - the repulsion is minimized by placing the four electron groups at corners of a tetrahedron with bond angles of 109°. - the shape with two bonds and two lone pairs on the central atom is called **bent**.  **Study Check** Predict the shape of a molecule of H~2~S. State the number of electron groups and lone pairs and use VSEPR theory to determine the shape of the following molecules or ions as tetrahedral, trigonal pyramidal, or bent. PF~3~ H~2~O CCl~4~ Polarity of Molecules --------------------- 1. The difference in electronegativity between atoms must result in polar bonds 2. The shape of the atom must result in a polar molecules. ### Nonpolar molecules - H~2~, Cl~2~, and O~2~ are nonpolar because they contain nonpolar bonds. - Molecules with polar bonds can be nonpolar if the polar bonds (dipoles) cancel in a symmetrical arrangement, such as in CO~2~ and CCl~4~.  ### Polar molecules - the electrons are shared unequally in the **polar covalent bond**. - one end of the molecule is more negatively charged than the other. - the polar bonds in the molecule do not cancel each other.  **Learning Check** Identify each of the following molecules as polar or nonpolar. PBr~3~ HBr CH~4~ 6.9 Intermolecular Forces in Compounds ====================================== Intermolecular Forces are forces between molecules. - Strong forces will hold molecules together as a solid - Weaker forces will allow molecules to slide past each other in a liquid - The weakest forces will not hold molecules together, so they fly apart and form a gas - Temperature and pressure also affect whether a solid, liquid, or gas is formed. - Warmer temperatures favour the formation of a gas - Higher pressures favour the formation of a solid Dipole--Dipole and Hydrogen Bonds --------------------------------- **Dipole-Dipole attractions** - polar molecules exert intermolecular forces called dipole--dipole attractions. **Hydrogen Bonds** - Some polar molecules form strong dipole attractions called **hydrogen bonds** between **hydrogen** **atoms bonded to F, O, or N**, and a lone pair on F, O, or N. - Hydrogen bonds are the strongest force between molecules and play a major role in the shape of DNA and protein molecules.  **Hydrogen Bonds and Water** A hydrogen bond occurs between the partially positive hydrogen atom in one molecule of water and the partially negative oxygen atom in another water molecule. Dispersion Forces ----------------- - weak attractions between nonpolar molecules. - caused by *temporary dipoles* that develop when molecules bump into each other. - weak but make it possible for nonpolar molecules to form liquids and solids.  Intermolecular Forces and Melting Points ---------------------------------------- Melting points of compounds: - related to the strength of intermolecular forces between molecules or compounds. - Melting points are lower due to weak forces such as dispersion forces. - are higher due to stronger intermolecular forces such as hydrogen bonding. - are highest in ionic compounds due to the strong forces between positive and negative ions in the compound.  **\ ** **Learning Check** Identify the main type of intermolecular forces that are present in liquids of the following compounds: ionic bonds, dipole--dipole, hydrogen bonds, or dispersion forces.