Liquids and Solids
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Intermolecular Forces are very significant when dealing with liquids and solids.
Three Types of Chemical Bonds Ionic Bonds Between Metal and Nonmetal Nonmetal takes electron of metal and forms anion, leaving the metal as a cation. The two ions then attract each other
Covalent Bonds Between Nonmetal and Nonmetal The two molecules share an electron pair Metallic Bonds Between Metal and Metal Characteristics of Matter Solids Have definite shape incompressible Usually have higher density than liquids Nearly Are not fluid Diffuse only very slowly through solids Have an ordered arrangement of particles that are very close together; particles usually only have vibrational motion
Have strong IMFs, but have low kinetic energies Liquids Have no definite shape (assume shapes of containers) Have definite volume (Very slightly compressible)
Have high density Are fluid Diffuse through other liquids Consist of disordered clusters of particles that are quite close together; particles have random motion
Gases Have no definite shape (fill containers completely) Are compressible Have low density Are fluid Diffuse rapidly
Consist of extremely disordered particles with a lot of empty space between then; particles have rapid, random motion
Have high kinetic energies, but have weak IMFs Intermolecular Vs. Intramolecular Intramolecular forces (Chemical bonding) are the attractive forces between adjacent atoms within a molecule
Intermolecular forces (van der Waals forces) are the attractive forces between adjacent molecules in a condense sample
Intermolecular forces are much weaker than intramolecular bonding Types of IMF Ion-Ion Interactions Attractive forces between cations and anions in ionic compounds Strength measured by Coulomb’s Law The higher the melting point, the stronger the ionic bond Dipole-Dipole interactions These result from permanent dipoles of polar covalent molecules. Its strength depends on the polarity of the molecules.
IMFs among nonpolar molecules (Dispersion Forces) Dispersion forces result from temporary dipoles of molecules Also known as London Dispersion Forces The strength of dispersion forces depends on the size or polarizability of the molecule (The larger a molecule is, the higher the polarizability, and thus the greater dispersion force)
Dispersion forces are the only IMFs that exist among nonpolar molecules forces also exist among polar molecules, and usually Dispersion play the major role
Hydrogen Bonding Formed when an extremely electronegative element is bonded with hydrogen. The molecule formed will leave the hydrogen atom almost naked as the attractive forces of either Fluorine, Oxygen, and Nitrogen will strongly attract the electron. The hydrogen will then attract another lone pair, resulting in a hydrogen bond.
Hydrogen bonding is not chemical bonding, but is a special type of dipole-dipole IMF
Hydrogen bonding is stronger than normal van der Waals forces, but still much weaker than chemical bonds
To form a H-bonding, both H-bonding donor (X-H, X = F, O, N) and H-bonding accepter (-Y, Y = F, O, N) must be present in the same system.
Properties of Liquids Viscosity Viscosity is the ability of resistance to flow The viscosity of a liquid depends on its intermolecular forces, the greater the IMF, the greater the viscosity
Surface Tension Surface tension is a measure of the unequal attractions that occur at the surface of a liquid. The molecules at the surface are attracted unevenly
Capillary Action Forces – IMF among liquids Adhesive Forces – Forces between liquid molecules and the surface Cohesive of the container
When the diameter of the glass tube is very small, A.F.>>C.F., causing water to creep up inside the tube, forming a meniscus.
Evaporation Evaporation (Vaporization) is the process by which some molecules on the surface of a liquid overcome the attractive forces (IMF) from other molecules and go into the gas phase.
The rate of evaporation depends on temperature a closed container, at a given temperature, eventually the rate of In evaporation = the rate of condensation. This is called dynamic equilibrium
When in equilibrium, the total number of molecules in gas phase would remain the same, so the pressure of the vapor over the liquid would also keep constant
The pressure of the vapor over a liquid at equilibrium is called the Vapor Pressure of the liquid at a given temperature T.
Vapor Pressure The higher the temperature, the higher the VP The higher the IMF, the lower the VP
Boiling Point The BP is the temperature at which the liquid starts to boil since its vapor pressure at BP equals the external pressure
Normal boiling point is when Pexternal = 1 atm BP increases with increasing IMF BP increases with increasing external pressure Distillation is a method used to separate a mixture of two or more different liquids with different boiling points
Heat Transfer Involving Liquids Heat transfer is involved for a liquid when there is a temperature change or phase change
Specific heat is the amount of heat needed for 1 gram of liquid to raise its temperature by 1° C [J/g*°C]
Molar heat capacity is the amount of heat needed for 1 mole of liquid to raise its temperature by 1° [J/mol*°C]
Heat of vaporization is the amount of heat needed for 1 mole/gram of liquid to be converted to vapor at its boiling point (ΔHvap) [kJ/mol]
Heat of condensation is the amount of heat released by 1 mole/gram of vapor to be converted to liquid at its condensing point (ΔHcond)
[kJ/mol]
Volatile liquids have weak IMF, Nonvolatile liquids have strong IMF Properties of Solids Melting Point The MP is the temperature at which the solid melts (same as FP) normal MP is when Pexternal = 1 atm The External pressure only has small effects on MP/FP
Heat Transfer involving Solids The molar heat/enthalpy of fusion is the amount of heat required to melt one mole of a solid at its melting point
The heat/enthalpy of solidification of a liquid is equal in magnitude to the heat of fusion.
Sublimation/Vapor Pressure of Solids Sublimation is the process where a solid directly becomes a gas Deposition is the process where a gas turns into a solid Amorphous vs. Crystalline
Amorphous solids DO NOT have definite shape and definite melting points
Crystalline solids DO have definite shape and definite melting points In a crystalline solid, all particles are arranged in a well-defined order in 3-D space and inter-particle attractive forces are equal throughout the whole crystal
In an amorphous solid, the arrangement of the particles is irregular and inter-particle attractive forces are different in different positions
Structures of Crystals All particles (atoms, molecules, or ions) are arranged in the same order throughout in the whole crystal
A unit cell is the smallest unit of a crystal that can show all the characteristics of the arrangement of the particles. Huge number of unit cells are stacked in 3-D space to form a crystal.
Considering each particle as a geometric point in 3-D space, the resulting 3-D array of all such points is called the lattice of the crystal
Isomorphous – referring to different substances with the same lattice
Polymorphous – referring to same substances with different lattices Seven crystal systems
Cubic, tetragonal, orthorhombic, monoclinic, triclinic, hexagonal, rhombohedral
Metallic
Ionic
Metal ions in “electron cloud”
Metallic bonds attraction between cations and ’
e
s
Soft to very hard; good thermal and electrical
Anions, cations Electrostatic Hard; brittle; poor thermal and electrical conductors; high melting points (400 to 3000°C)
conductors; wide range of melting points (39 to 3400°C) Li, K, Ca, Cu, Cr, Ni ( metals)
NaCl, CaBr2, K2SO4 (typical salts)
Molecular
Covalent
Molecules (or atoms) Atoms Covalent bonds Dispersion, dipole–dipole, and/or hydrogen bonds Very hard; poor thermal and electrical conductors;* Soft; poor thermal and electrical conductors; low
high melting points (1200 to 4000°C)
melting points (272 to 400°C) C (diamond), SiO2 (quartz) CH4 (methane), P4, O2, Ar, CO2, H2O, S8
Solutions
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Solutions Solutions are homogeneous mixtures In a solution, one component (dissolving medium) is called a solvent and all others (dissolved species) are solutes
Spontaneity of the Dissolution Process Dissolution - A solute dissolves in a solvent without any chemical reaction
Spontaneity – How easy or difficult can a solute dissolve in a solvent? For a given solute, the spontaneity of its dissolution can be quantitatively measured with its solubility – How many grams of solute can at most dissolve in 100 grams of solvent?
If solubility is high, the solute is soluble in the solvent If solubility is low, the solute is insoluble in the solvent Two factors that affect dissolutions Energy change of dissolution – Heat of solution ΔH Solution
If ΔHSolution < 0, dissolution is favored If ΔHSolution > 0, dissolution is NOT favored
Change in disorder/randomness of the solution, ΔS If ΔS < 0, disorder , dissolution is favored If ΔS > 0, disorder , dissolution is NOT favored
Dissolution is always favored by this factor (ΔS>0)
Spontaneity of the dissolution of a solute in a solvent depends on the energy change, ΔHSolution
General Trends Weak solute–solute attractions favor solubility. solvent–solvent attractions favor solubility. Weak Strong solvent–solute attractions favor solubility.
Dissolution of Ionic Solids in Water Step 1 – Expansion of solute Step 2 – Expansion of solvent Step 3 – Solvation of solute ions Hydration energy is defined as the energy change involved in the (exothermic) hydration of one mole of gaseous ions.
ΔHSolution = (heat of solvation) - (crystal lattice energy) Solubilities of Ionic Solids in Water For ionic solids, when ionic charge increases, size decreases
Dissolution of Nonpolar Solids in Liquids Nonpolar solids are insoluble in polar solvents, but soluble in nonpolar solvents (like dissolves like)
Dissolution of Liquids in Liquid Solvent Spontaneity of dissolution of a liquid solute in a liquid solvent can be quantitatively measured by its miscibility
Polar is miscible in polar, nonpolar is miscible in nonpolar Polar is immiscible in nonpolar and vice versa Dissolution of Gases in Liquid Solvents Most gases are slightly soluble in liquids Some gases can dissolve easily in water because of the chemical reaction with water molecules
Rates of Dissolution and Saturation When the rate of dissolution = the rate of crystallization, a dynamic equilibrium is established at a given T the amounts of solute in solid and in liquid phase would keep constant A solution at equilibrium is said to be saturated phase Supersaturated solutions are metastable
Effect of Temperature on Solubility The effect of T on solubility depends on the sign (+/-) of the heat of solution, ΔHSolution
ΔHSolution > 0, endothermic, T, solubility ΔHSolution < 0, exothermic, T, solubility Dissolutions of moist ionic solids are endothermic, and so T , solubility Dissolution of gases in exothermic, and so T , solubility by temperature Dissolutions of most liquids is only slightly influenced because ΔHSolution ≈ 0
Effect of Pressure on Solubility The effect of pressure is negligible for dissolution of solid or liquid The solubility of gases in liquid solvents increase as the pressures of the gases increase
Henry’s Law – At a given temperature, the concentration of a gas in a liquid is directly proportional to the partial pressure of the gas over the solution
Colligative Properties of Solutions
Comparing to the pure solvent, physical properties of a solution would change more or less
Some property changes depend only on the number, but not the kind, of solute particles in a given amount of solvent – colligative properties
Vapor Pressure lowering Boiling point elevation Freezing point depression Osmotic pressure production
Molality Molality is a concentration unit based on the number of moles of solute per kilogram of solvent
Mole Fraction vs. Vapor Pressure Mole Fraction Mole fraction is the number of moles of one compound divided by the total moles of all the components of the solution
Vapor Pressure Lowering A nonvolatile solute is one that stay in the liquid phase, and will not become a gas (I.E. sugar)
Raoult’s Law – When a nonvolatile/non-ionizing solute dissolves in a solvent, the vapor pressure of the solvent in an ideal solution is directly proportional to the mole fraction of the solvent in the solution
If a solution consists of two volatile components, A and B, the relationship between partial vapor pressure and mole fraction for each component still obeys Raoult’s Law
Fractional Distillation distillation is a technique used to separate solutions that Fractional have two or more volatile components with different boiling points
Boiling/Freezing Point Elevation Boiling point elevation of a solution to its pure solvent is ΔT b = Kbm Colligative Properties of Electrolyte Solutions Colligative properties depend only on the numbers of solute particles in a given mass of solvent
In a solution, some cations and anions can undergo temporary association to form ion pairs
The formation of the ion pairs reduces the total number of the solute particles in the solution. Or, it reduces the effective molality of the solute
van’t Hoff Factor The ratio of the actual colligative property to the value that would be observed if no dissociation occurred
If the extent of dissociation increases, the extent of temporary association decreases, the total number of particles increases, the effective molality increases, and so the van’t Hoff factor, i, increases
Osmotic Pressure Osmotic pressure depends on the number, and not the kind, of solute particles in solution; it is therefore a colligative property. It is the pressure exerted by the solution due to osmosis.
Osmosis is the spontaneous process by which the solvent molecules pass through a semi- permeable membrane from a solution of lower concentration of solute into a solution of higher concentration of solute.
Osmotic pressure (π) is directly proportional to the molar concentration of the solution and its absolute temperature
Π = MRT where π = osmotic press ure, R = 0.0821 L*atm/mol*K, and M = n/v = molar concentration = moles/liter
Thermodynamics
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Introduction Thermodynamics is the science of energy Chemical Thermodynamics is the application of thermodynamic theory to the study chemical reactions
Heat Changes and Thermochemistry Energy is the capacity to do work or to transfer heat. Potential Energy – The energy that a system possesses by virtue of its position or composition
Kinetic Energy – The energy of motion Universe – System + Surroundings Two types of reactions Exothermic
Reactions that release energy in the form of heat are called exothermic reactions
Endothermic
A process that absorbs energy from its surroundings is called endothermic
First Law of Thermodynamics (Law of Conservation of Energy) Energy is neither created nor destroyed in chemical reactions and physical changes
Some important terms System The substances involved in the chemical and physical changes under investigation
Surroundings The environment around the system Thermodynamic state of a system Is defined by a set of conditions (T, P, composition, and physical state) that completely specifies all the properties of the system. Once the state has been specified, all other properties are fixed.
State functions The value of a state function depends only on the state of the system and not on the way in which the system came to be in that state.
A change in a state function describes a difference be- tween the two states. It is independent of the process or pathway by which the change occurs.
Internal Energy The sum of all energy contained within a system (KE + PE) Does not include KE or PE of the whole system Depends only on the state, internal energy is a state function For a chemical or physical process, ΔE = ΔEinitial - ΔEfinal = q + w q = heat, w = work Energy Exchanges Heat (q) – energy change due to different temperatures Work (w) – any energy exchanges other than q – P.V.E (Pressure Volume Energy, mechanical, light, electrical energies etc.)