Lecture 1: States of Matter - Intermolecular Forces (PDF)

Lecture 1 States of Matter – Intermolecular Forces Let’s start studying of chemistry by understanding matter and its physical states at the macroscopic and molecular level through studying the intermolecular forces that affect the physical states of gases, liqui...

Lecture 1 States of Matter – Intermolecular Forces Let’s start studying of chemistry by understanding matter and its physical states at the macroscopic and molecular level through studying the intermolecular forces that affect the physical states of gases, liquids and solutions. Classification of Matter: A sample of matter can be a gas, a liquid, a solid or a plasma. These four forms, called the states of matter, differ in some of their observable properties. A gas (also known as vapor) has no fixed volume or shape; rather, it uniformly fills its container. A gas can be compressed to occupy a smaller volume, or it can expand to occupy a larger one. A liquid has a distinct volume independent of its container, assumes the shape of the portion of the container it occupies, and is not compressible to any appreciable extent. A solid has both a definite shape and a definite volume and is not compressible to any appreciable extent. A plasma forms at high temperature, nuclear reactors or inside stars gases are ionized to form a mixture of roughly equal number of positively and negatively charged particles. The negative charge is usually carried by electrons, each of which has one unit of negative charge. The positive charge is typically carried by atoms or molecules that are missing those same electrons. In some rare but interesting cases, electrons missing from one type of atom or molecule become attached to another component, resulting in a plasma containing both positive and negative ions. Characteristic Properties of the States of Matter: The following figure compares the common three states of matter. The state of a substance depends largely on the balance between the kinetic energies of the particles (atoms, molecules, or ions) and the interparticle energies of attraction. The kinetic energies, which depend on temperature, tend to keep the particles apart and moving. The interparticle attractions tend to draw the particles together. Substances that are gases at room temperature have much weaker interparticle attractions than those that are liquids; substances that are liquids have weaker interparticle attractions than those that are solids. The different states of matter adopted by the halogens at room temperature—iodine is a solid, bromine is a liquid, and chlorine is a gas—are a direct consequence of a decrease in the strength of the intermolecular forces as we move from I2 to Br2 to Cl2. Figure. Gases, liquids, and solids. Chlorine, bromine, and iodine are all diatomic molecules as a result of covalent bonding. However, due to differences in the strength of the intermolecular forces, they exist in three different states at room temperature and standard pressure: Cl2 is a gas, Br2 is a liquid, and I2 is a solid. Chlorine, Cl2 Bromine, Br2 Iodine, I2 Particles are far apart; Particles are closely packed Particles are closely packed possess complete freedom but randomly oriented; inan ordered array; of motion; kinetic energies retain freedom of motion; positionsare essentially of particles >> energies of kinetic energies of particles fixed; energies of particle– particle–particle attraction similar to energies of particle attraction >> particle–particle attraction kinetic energies of particles. Intermolecular interactions: Intermolecular forces (Also known as van der Waals forces, after description by Johannes D. van der Waals (1837–1923)) refer to the forces between individual particles (atoms, molecules, ions) of a substance. These forces are quite weak relative to intramolecular forces, that is, covalent and ionic bonds within compounds. For example, 927 kJ of energy is required to decompose one mole of water vapor into H and O atoms. This reflects the strength of intramolecular forces (covalent bonds). Only 40.7 kJ is required to convert one mole of liquid water into steam at 100°C. This reflects the lower strength of the intermolecular forces of attraction between the water molecules, compared to the covalent bonds within the water molecules. The attractive forces between water molecules are mainly due to hydrogen bonding. There are four types of intermolecular forces (vdW forces): Permanent Dipole-Permanent Dipole (Kessom Forces) Permanent Dipole-Induced Dipole (Debye Forces) Induced Dipole – Induced Dipole (London or Dispersion Forces) Hydrogen Bonding Permanent Dipole-Permanent Dipole: Due to the difference in electronegativity between two atoms bonded by covalent bond, a dipole moment originates due to presence of δ+ on the atom with the lowest electronegativity and δ- on the atom with the highest electronegativity. Permanent dipole–dipole interactions occur between polar covalent molecules because of the attraction of the δ+ atoms of one molecule to the δ- atoms of another molecule. Fig. Dipole–dipole interactions among polar molecules. Each polar molecule is shaded with regions of highest negative charge (δ-) in red and regions of highest positive charge (δ+) colored blue. Attractive forces are shown as blue arrows, and repulsive forces are shown as red arrows. Stronger attractions and repulsions are indicated by thicker arrows. Molecules tend to arrange themselves to maximize attractions by bringing regions of opposite charge together while minimizing repulsions by separating regions of like charge. These forces are weaker than ion–ion forces. Average dipole–dipole interaction energies are approximately 4 kJ per mole of bonds. They are much weaker than ionic and covalent bonds, which have typical energies of about 400 kJ per mole of bonds. Substances in which permanent dipole–dipole interactions affect physical properties include bromine fluoride, BrF, and sulfur dioxide, SO2 Permanent Dipole-Induced Dipole: A polar molecule can induce a dipole moment in a nonpolar molecule or an atom. When the two are brought together, the arrangement of electrons in the nonpolar molecule will be disrupted. As a result, it will acquire a partially positive charge and a partially negative charge. The two molecules will attract one another through dipole-induced dipole forces. It is a weaker interaction than the regular dipole-dipole interaction. Induced Dipole – Induced Dipole: Dispersion forces are weak attractive forces that are important only over extremely short distances. These forces are present between all types of molecules in condensed phases but are weak for small molecules. Dispersion forces are the only kind of intermolecular forces present among symmetrical nonpolar substances such as CO2, O2, N2, Br2, H2, and monatomic species such as the noble gases. Without dispersion forces, such substances could not condense to form liquids or solidify to form solids. Condensation of some substances occurs only at very low temperatures and/or high pressures.  Dispersion forces result from the attraction of the positively charged nucleus of one atom for the electron cloud of an atom in nearby molecules. This induces temporary dipoles in neighboring atoms or molecules. As electron clouds become larger and more diffuse, they are attracted less strongly by their own (positively charged) nuclei. Thus, they are more easily distorted, or polarized,by adjacent atoms or molecules. Polarizability increases with increasing numbers of electrons and therefore with increasing sizes of molecules. Therefore, dispersion forces are generally stronger for molecules that are larger or that have more electrons. Fig. Dispersion forces. “Snapshots” of the charge distribution for a pair of helium atoms at three instants. Learning Check: Why is the boiling point of the halogen in each period greater than the noble gas? Molecular shape affects intermolecular attraction: Hydrogen Bonding: Hydrogen bonds are a special case of strong dipole–dipole interaction. They are not really chemical bonds in the formal sense. Strong hydrogen bonding occurs among polar covalent molecules containing H bonded to one of the three small, highly electronegative elements—F, O, or N. Hydrogen bonds result from the electrostatic attractions between δ+ atoms of one molecule, in this case H atoms, and the δ - atoms of another molecule. The small sizes of the F, O, and N atoms, combined with their high electronegativities, concentrate the electrons of these molecules around these δ- atoms. This causes an H atom bonded to one of these highly electronegative elements to become quite positive. The δ + H atom is attracted to a lone pair of electrons on an F, O, or N atom other than the atom to which it is covalently bonded. The molecule that contains the hydrogen-bonding δ+ H atom is often referred to as the hydrogen bond donor; the δ- atom to which it is attracted is called the hydrogen-bond acceptor.  Typical hydrogen-bond energies are in the range 15 to 20 kJ/mol, which is four to five times greater than the energies of other dipole–dipole interactions. As a result, hydrogen bonds exert a considerable influence on the properties of substances. Hydrogen bonding is responsible for the unusually high melting and boiling points of compounds such as water, methyl alcohol, and ammonia compared with other compounds of similar molecular weight and molecular geometry. Hydrogen bonding between amino acid subunits, for example, is very important in establishing the three-dimensional structures of proteins. Table. Approximate Contributions to the Total Energy of Interaction Between Molecules, in kJ/mol Q: Identify the types of intermolecular forces that are present in a condensed phase (liquid or solid) sample of each of the following: (a) Water, H2O, (b) Iodine, I2 and (c) Nitrogen dioxide, NO2 The Dissolution Process: A solution is formed when one substance disperses uniformly throughout another. A solution is defined as a homogeneous mixture, at the molecular level, of two or more substances in which phase separation does not occur. A solution consists of a solvent and one or more solutes. The solvent is the medium in which the solutes are dissolved. Solutes usually dissolve to give ions or molecules in solution. Solutions include different combinations in which a solid, liquid, or gas acts as either solvent or solute. Usually the solvent is a liquid. For instance, seawater is an aqueous solution of many salts and some gases such as carbon dioxide and oxygen. Many naturally occurring fluids contain particulate matter suspended in a solution. For example, blood contains a solution (plasma) with suspended blood cells. Solutions in which the solvent is not a liquid are also common. Air is a solution of gases with variable composition. Alloys are solid solutions of solids dissolved in a metal. The solvent is usually the most abundant species present. If we mixed 10 grams of alcohol with 90 grams of water, we would call alcohol the solute. If we mixed 10 grams of water with 90 grams of alcohol, water would be considered the solute. The dissolution process can occur with or without a chemical reaction: 1. With a chemical reaction: If the resulting solution is evaporated to dryness, solid sodium hydroxide, NaOH, is obtained rather than metallic sodium.  2. Without a chemical reaction: Evaporation of the water from the sodium chloride solution yields the original NaCl. The ease of dissolution of a solute depends on two factors that accompany the process: (1) The change in energy. (2) The change in disorder (called entropy change). The process is favored by (1) a decrease in the energy of the system, which corresponds to an exothermic process, and (2) an increase in the disorder, or randomness, of the system (The entropy of a system increases if its degree of disorder increases, or if its energy becomes dispersed over a greater number of particles). The energy change that accompanies a dissolution process is called the heat of solution, Δ Hsolution. This change depends mainly on how strongly the solute and solvent particles interact. The main interactions that affect the dissolution of a solute in a solvent follow. a. Weak solute–solute attractions favor solubility. b. Weak solvent–solvent attractions favor solubility. c. Strong solvent–solute attractions favor solubility. \ We can imagine the solution process as having three components, each with an associated enthalpy change: Fig. Enthalpy changes accompanying the solution process. Step (a): The intermolecular or interionic attractions among solute particles in the pure solute must be overcome to dissolve the solute. This part of the process requires an absorption of energy (endothermic). Step (b): Separating the solvent molecules from one another to “make room” for the solute particles also requires the absorption of energy (endothermic). Step (c): As the solute particles and solvent molecules interact in the solution, energy is released (exothermic). The overall dissolution process is exothermic (and favored) if the amount of heat absorbed in hypothetical Steps a and b is less than the amount of energy released in Step c (ΔHsolution0). However, many solids do dissolve in liquids by endothermic processes. The reason such processes can occur is that the endothermicity can be compensated by a large increase in disorder of the solute (remember: the second factor in the dissolution process) during the dissolution process. The solute particles are highly ordered in a solid crystal, but are free to move about randomly in liquid solutions. Likewise, the degree of disorder in the solvent increases as the solution is formed, because solvent molecules are then in a more random environment. They are surrounded by a mixture of solvent and solute particles. Most dissolving processes are accompanied by an overall increase in disorder. Dissolution of Solids in liquids: We have to ask ourselves some questions to understand the dissolution process deeply: Why does NaCl dissolve in water spontaneously? By contrast naphthalene (C10H8) doesn’t dissolve in water and dissolve easily in benzene (C6H6). Why does glucose (C6H12O6) dissolve in water spontaneously although it is not an ionic compound?  Fig. Dissolution of the ionic solid NaCl in water, the disorder factor is usually favorable to solubility What makes some ionic compounds dissolve in water spontaneously and release heat, while other compounds absorb heat and also dissolve in water? Why do some ionic compounds dissolve sponateously in water? To answer all these questions we have to learn about crystal lattice, solvation and hydration energies. The crystal lattice energy: is defined as the energy change accompanying the formation of one mole of formula units in the crystalline state from constituent particles in the gaseous state. This process is always exothermic; that is, crystal lattice energies are always negative.  The amount of energy involved in this process depends on the electrostatic attraction between ions in the solid. When these attractions are strong due to higher charges and/or small ion sizes, a large amount of energy is released as the solid forms, increasing its stability. The reverse of the crystal formation reaction is breaking the crystal into separated gas phase ions. So we can consider this step is the step (a) in the dissolution process which is endothermic and if the compound has a small value for crystal lattice energy, the dissolution is more favorable. If the solvent is water, the energy that must be supplied to expand the solvent (Step b) includes that required to break up some of the hydrogen bonding between water molecules. The third major factor contributing to the heat of solution is the extent to which solvent molecules interact with particles of the solid. The process in which solvent molecules surround and interact with solute ions or molecules is called solvation. When the solvent is water, the more specific term hydration is used. Hydration energy: (equal to the sum of Steps b and c) is defined as the energy change involved in the (exothermic) hydration of one mole of gaseous ions. Hydration is usually highly exothermic for ionic or polar covalent compounds because the polar water molecules interact very strongly with ions and polar molecules. In other words, the compounds that are able to form ion-dipole interaction or hydrogen bond with water. Heat of solvation is always exothermic and has a negative value. Crystal lattice energy is always negative, so the second term in the right side of equation is always positive because of the minus sign. Therefore, the dissolution process is the result of the net value of both terms. Nonpolar solids such as naphthalene, C10H8, do not dissolve appreciably in polar solvents such as water because the two substances do not attract each other significantly. Naphthalene dissolves readily in nonpolar solvents such as benzene because there are no strong attractive forces between solute molecules or between solvent molecules. In such cases, the increase in disorder controls the process. These facts help explain the observation that “like dissolves like.” Factors affect the crystal lattice energy: Charge of ions: increases the magnitude of crystal lattice energy. Size of ions : smaller sizes increase the magnitude of crystal lattice energy. For low-charge species, hydration energies and lattice energies are usually of about the same magnitude, so they often nearly cancel each other. As a result, the dissolution process is slightly endothermic for many ionic substances. Ammonium nitrate, NH4NO3, is an example of a salt that dissolves endothermically. Many ionic solids dissolve with the release of heat such as anhydrous sodium sulfate, Na2SO4. As the charge-to-size ratio increases for ions in ionic solids, the magnitude of the crystal lattice energy usually increases more than the hydration energy. This makes dissolution of solids that contain highly charged ions—such as aluminum fluoride, AlF3 too endothermic to be soluble in water. Dissolution of liquids in liquids: The attractive forces between solute-solute molecules is much weaker than that of the solid, so this factor is less important in mixing. The mixing process is always exothermic and is called in this case miscibility which equal to solubility. All intermolecular forces that control dissolution process are also considered in the liquid-liquid interactions which follows the same rule “like dissolves like”. Dissolution of Gases in liquids: We should expect that polar gases are most soluble in polar solvents and nonpolar gases are most soluble in nonpolar liquids. Although carbon dioxide and oxygen are nonpolar gases, they do dissolve slightly in water. CO2 is somewhat more soluble because it reacts with water to some extent to form carbonic acid, H2CO3. This in turn ionizes slightly in two steps to give hydrogen ions, bicarbonate ions, and carbonate ions. Oxygen, O2, is less soluble than CO2 in water, but it does dissolve to a noticeable extent due to Debye forces (Dipoles-induced dipole). The only gases that dissolve appreciably in water are those that are capable of hydrogen bonding (such as HF), those that ionize extensively (such as HCl, HBr, and HI), and those that react with water (such as CO 2 or SO3).  Exercises 1. What are the intermolecular forces in the following molecules? (a) Cl2, Cl2 (b) NH3, H2O (c) Cl2, H2O (d) HCHO, CH3OH (e) C6H12, CH3COCH3 (f) CO2, H2O 2. Explain how the hydrogen bond is formed in the following molecules. (a) HF, CH3OH (b) NH3, H2O (c) CH3OH, CH3COCH3 3. Explain why these compounds are soluble or not in following solvents. (a) HCl and H2O (soluble) (b) HF and H2O (Soluble) (c) Al2O3 and H2O (Insoluble) (d) SiO2 and H2O (slightly soluble) (e) Na2CO3 and C6H14 (Insoluble) 4. “Oil and water don’t mix” Explain the molecular basis of this quote.