Atomic Structure and Properties Study Guide Part 2

Molecular and Ionic Compound Structure and Properties Around 7‒9% of the questions on your AP exam will cover Molecular and Ionic Compound Structure and Properties. Types of Chemical Bonds Chemical bonds are the attractive forces between atoms that hold them together, including sharing electrons and other electrostatic forces. These forces can be intramolecular, as in they occur within a molecule, or intermolecular, as in occurring between molecules. Intramolecular forces are stronger than intermolecular forces. The main types of intramolecular binding forces are summarized in the following table. Bond Description Covalent The sharing of electrons between a pair of atoms, also called a molecular bond. Covalent bonds can result in atoms having complete valence shells, stabilizing the atoms. These are usually gasses, liquids, or soft solids and include most biological molecules. The first covalent bond is called a sigma bond (  ) and additional bonds are called pi bonds (). The sigma bond is the most stable. Ionic bonds describe the electrostatic attraction between cations and anions or with charged polyatomic ions. Ionic bonds involve the transfer of electrons from a metal to a non-metal, resulting in full valence shells for both atoms. Ionic bonds tend to be easily separated by external forces, like the addition of polar solvents, causing substances to break apart into anions and cations. Solid, ionically bonded structures form ordered and strong lattices that are nonconductive and have high melting and boiling points. This group includes salts. Metallic bonds are the sharing of free electrons in a metallic structure, composed of charged cations. In metallic bonds, valence electrons are only loosely bound to the nucleus and form a sort of cloud around metallic cations, allowing them to be shared. This results in a highly stable configuration of positively charged cations in a sea of negatively charged electrons. Metallic bonds account for certain features of metals like malleability, conductivity, and strength. Ionic Metallic 12 Example CO2 NaCl Copper wire, aluminum foil One way to predict whether a bond is ionic or covalent is to look at the electronegativity of the atoms involved. Electronegativity describes how much an atom attracts bonding electrons. The scale runs from 4 (fluorine) at the high end to 0.7 (cesium and francium) at the low end. On the periodic table, electronegativity increases from left to right and bottom to top (excluding noble gasses). Another way to think of electronegativity is as a quantitative way to consider how factors like the atomic structure, valence shells, and Coulomb’s law affect attraction between ions. To determine bond type, remember that when the difference between electronegativity is high (greater than 1.7), ionic bonds form. When the difference is low (1.7 and lower), covalent bonds form. Covalent bonds formed between atoms with similar electronegativities (<0.5) are nonpolar covalent bonds. This means that electrons are shared equally across the atoms. A polar covalent bond forms when there is a larger difference in electronegativity across atoms. The atom with the greater charge pulls electrons closer to it, forming a dipole. In a dipole, electrons spend more time at one end of the molecule, giving that side a slightly more negative charge ( ) and the other end a slightly positive charge ( +). This can result in molecules that are charged, or polar, and can interact with other molecules through these charges. The size of a dipole is measured by the dipole moment, ý. When the difference between the dipole is very high, the molecule at the + end of the bond essentially transfers its electrons to the more electronegative atom, forming an ionic bond. The direction of a dipole is ind icated in a bond as an arrow from the less electronegative atom to the more electronegative atom. Intramolecular Force and Potential Energy During covalent bond formation, unstable atoms with incomplete valence shells share electrons to increase their stability. Attractive forces between the nuclei and electron clouds of unstable atoms attract each other to form a bond. This is counteracted by repulsive forces between each atom’s nuclei and between each atom’s electron clouds. This balance between repulsive and attractive intramolecular forces can be shown in a potential energy diagram. 13 In this diagram, when the two molecules are far apart there is no potential energy, or stored energy. The closer the molecules get to each other, the lower the potential energy due to attractive forces between the molecules (remember Coulomb’s law). The distance where this potential energy is lowest is the average length of the covalent bond that forms between them. In other words, covalent bond lengths are optimal when the potential energy of the bond is lowest. You’ll notice in the diagram that when the molecules get even closer together, the potential energy rapidly shoots back up. This is due to repulsive intramolecular forces between the atoms. In the case of double and triple bonds, the bond length becomes shorter and the strength of the bond becomes higher; thus, you can expect to see shorter lengths between nuclei and higher bond energies. Structure of Ionic Solids Ionic solids form in a manner that reduces the repulsive forces between similarly charged ions and increases the attractive force between oppositely charged ions. This forms a chemical lattice with an orderly arrangement. The lattice energy (E) is the energy needed to separate 1 mol of an ionic solid into its gas phase. This is proportional to the product of the electrical charges of the ions (Q) divided by the sum of the radii of the atoms (r): ā∝ Ā1 Ā2 ÿ 14 Note that when the ions are oppositely charged, the lattice energy is negative. This indicates energy will be released when the lattice forms due to attraction between the molecules. The more negative the lattice energy is, the stronger the ionic bonds will be. These strong bonds holding ionic solids together explains some of the features of ionic solids such as crystal structures and brittle properties. Structure of Metals and Alloys Metallic solids are formed by an array of metallic cations surrounded by a sea of freely moving shared electrons. The loose shared electron sea helps to understand some of the features of pure metals, such as conductance, malleability, and ductility. When metals are mixed with other metals or non-metals, alloys are formed. When two metals of similar size are mixed, a substitutional alloy is formed. The percentages of metals in a substitutional alloy can be varied to alter the properties of the alloy, such as increasing strength and conductivity while decreasing corrosion. When a pure metal is mixed with a much smaller metal or non -metal, an interstitial alloy is formed. Interstitial alloys are stronger, harder, and less ductile than their pure host metals. Lewis Diagrams Lewis dot diagrams are a quick way to represent the valence electrons of an atom, where each dot around a molecule represents a valence electron. A covalent bond is represented by dots between atoms, or by dashes corresponding to the number of covalent bonds. According to the octet rule, most atoms are considered stable when this valence shell is occupied by eight atoms; thus, the optimal configuration is to have 2 electrons on each of the 4 sides or 4 total bonds. Hydrogen atoms only need two electrons in their valence shells. 15 When drawing a dot diagram for a compound, first arrange the ions so they form a symmetrical structure. Remember that hydrogen is always on the outside of the structure, never central. Then calculate the total number of valence electrons the molecule has vs. the number of molecules needed (generally 8 for every molecule except hydrogen). The difference between the number of electrons the molecule has and the number needed tells you how many electrons are shared. Then divide shared electrons by two to find out how many are between elements in this molecule. You then draw the bonds and complete the structure by adding the remaining electrons so that each atom has the correct number of valence electrons. Lewis dot diagrams do not handle unpaired electrons well or show the difference between the strength of bonds; these cases are better described by models that take into account 3-D structure. Resonance and Formal Charge Resonance structures are not adequately represented by dot diagrams. These are instance where two forms of a molecule are possible with the same connectivity but with differences in the distribution of electrons. For instance, ozone can exist as O=O −O or O−O=O. In reality, these would exist as partial bonds, but for the purpose of demonstration would be written as follows: O=O−O  O−O=O. You may come across instances where there are multiple possible Lewis structures. To determine which is the most appropriate, you can calculate the formal charge of the atoms in the chemical compound: Formal charge = (# valence electrons) 2 (# nonbinding electrons + bonding electrons ) 2 The most favorable Lewis structures will have small formal charges on each element: more electronegative elements will have lower formal charges, like charges will be far apart, and unlike charges will be near. The total formal charge for the molecule is the sum of the formal charges of the atoms. Note that you may be asked on the exam to compare the relative contribution of different resonance structures possible for a molecule. In that case, the resonant structures that are more common are those that satisfy the octet (or duet) rule for all atoms that have a formal charge on each atom closest to zero, and where more electronegative elements have lower formal charges. 16 VSEPR and Bond Hybridization Valence shell repulsion theory, or VSEPR, asserts that electron pairs repel each other and unbonded electrons are even more repulsive to bonded pairs. This causes molecules to take on distinct shapes depending on the number of bonds and hybridized orbitals. For example, a molecule with two sigma bonds will take on a linear shape so that both bonds will be as far apart as possible. A molecule with four bonds will take on a tetrahedral shape to optimize distance. You will need to estimate molecular shapefor the exam using VSEPR. To determine the shape of a chemical compound using VSEPR: 1. Draw out the Lewis dot diagram. 2. Count the number of electron-dense regions. These include each element in the compound as well as each pair of electrons. 3. Mentally arrange these electron dense regions to allow equal space between them; this gives the overall shape of the structure 4. Take into account electronic geometry, due to the number of lone electron pairs, as well as molecular geometry, due to the number of bonds. 17 Another way to represent molecular geometry is through Valence bond theory. In this theory, covalent bonds form through the mixing of orbitals to form hybrid orbitals. The number of hybrid orbitals formed is equal to the number of contributing orbitals. The types of orbitals that contribute determine the hybrid orbital type (e.g., s and p orbitals hybridizing means that you will have two sp orbitals). 18 Sample Molecular and Ionic Compound Structure and Properties Questions When two chlorine atoms are brought together, they form a Cl –Cl bond and energy is released. The potential energy for this system as a function of Cl –Cl distance is shown above. According to this graph, what amount of energy is required to break the Cl –Cl bond? A. 17 eV B. 8 eV C. 0 eV D. 28 eV Explanation: The correct answer is B. The right side of the graph asymptotically approaches 8 eV, which is the energy required to separate the two Cl atoms infinitely far apart. 20 Which of the following hydrocarbons is expected to have single bonds between all carbon atoms? A. propane (C3H8) B. butene (C4H8) C. hexene (C6H10) D. octyne (C8H14) Explanation: The correct answer is A. All of the bonds between the carbon atoms are the same (single bonds). The ending –ene signifies there are double bonds between carbon atoms along with single bonds and the ending –yne signifies there are triple bonds, which can be seen by drawing the Lewis diagram. Which of the following statements best explains why LiBr is expected to boil at a higher temperature than SF4? A. LiBr contains ionic bonds and SF4 has dipole-dipole forces B. The density of SF4 is greater than that of LiBr C. SF4 has a higher molar mass than LiBr D. LiBr is linear and SF4 is not Explanation: The correct answer is A. LiBr forms an ionic bond, because Li is a very active metal and Br is a very electronegative nonmetal. SF4 is a polar molecule with dipole-dipole forces between the molecules. Generally, species with ionic bonds have higher boiling points than species with dipole-dipoleforces. 21