03-StudyGuide-Concentration,Solubility,Gases,&ColligativeProperties PDF

Helpful Hints Solutions In general a solution consists of two major components: the solute and the solvent. The solute is the component in the lesser quantity while the solvent is the component in the greater quantity. Solutions can be composed of mixtures of various phase combinations. So...

Helpful Hints Solutions In general a solution consists of two major components: the solute and the solvent. The solute is the component in the lesser quantity while the solvent is the component in the greater quantity. Solutions can be composed of mixtures of various phase combinations. Solute liquid solid gas liquid solid gas liquid solid Phases Solvent gas gas liquid liquid liquid solid solid solid aerosol, solid foam, sol, solid solid Name emulsion solid foam fog aerosol froth suspension emulsion suspension fire insulating Example hair spray dust milk wet cement toothpaste plastics foam foam Liquid-Liquid Mixtures When dealing with mixtures of liquids the terms miscible and immiscible are used to reference substances that mix creating a homogeneous mixture and things that do not mix to form a homogeneous mixture, respectively. A major influence of miscibility is based on the intermolecular forces involved in the interaction of the solute with the solvent, as well as the solute and solvent with themselves; this is often expressed by the statement "like dissolves like" indicating that polar liquids are miscible with other polar liquids while nonpolar liquids are miscible with other nonpolar liquids. In general no outside factors such as temperature or pressure affect the miscibility of liquids. Gas-Liquid Solutions The solubility of a gas increases with increasing pressure as shown by Henry's Law where P is the pressure of the gas, kH is the Henry's Law constant of a given gas, and S is the molar solubility of the gas: ∙ In addition to pressure changes, temperature changes will also affect the solubility of a gas. Increasing temperature will generally cause a decrease in gas solubility because higher temperature means more kinetic energy for each molecule allowing them to escape into the vapor phase more readily. Solid-Liquid Mixtures When referring to a mixture of a solid and liquid it is typical to use the term soluble for a solid that dissolves and insoluble for a solid that does not dissolve. The underlying cause of whether something is soluble relates to whether the solvent and solute interactions are greater than those of the solvent and solute with themselves. Another factor that affects solubility which will be discussed in further detail later is the tendency of mixing based on the lower energy involved in a more disordered state, this tendency to mix is called entropy. Solutions of dissolved solids can be classified in three ways with respect to the amount of dissolved solute. A solution with less than the maximum amount of solute dissolved is referred to as an unsaturated solution, while a solution with the maximum amount of solute dissolved is referred to as a saturated solution. The final classification is less obvious which relates to having more than the maximum amount of solute dissolved and is referred to as a supersaturated solution. Supersaturation is an unstable state that can only be reached through a careful process that generally involves heating the solvent, dissolving a large amount of solute, and then slowly cooling the solution while ensuring that it is not disturbed to prevent precipitation. The following table is useful when determining whether an ionic compound is soluble in water: *It is important to be familiar with the general solubility rules, such as all nitrates are soluble; attempting to memorize all rules would be overly difficult, and unnecessary. Solubility is a measure of the maximum possible solute concentration and can be expressed in terms of various units of concentration, for example in terms of molarity (M). The table below gives several additional units of concentration that may be used. Dilution It is often necessary to prepare a needed solution from an initial highly concentrated solution, the stock solution; this process is called dilution. Dilution involves transferring a sample of the stock solution to another container and adding additional solvent. It is important to realize that the number of moles of solute in the transferred sample does not change, only the volume changes. For the case of molarity the following equation can be used for dilution calculations: C1V1 = C2V2 because n1 = n 2 The number of moles in the transferred sample does not change, only the volume changes. Pressure A key feature of gas molecules is they are in constant motion exerting a pressure as they interact with their containment boundaries. Units of pressure include, but are not limited to bar, atmospheres (atm), Torricellis (Torr), Pascals (Pa), pounds per square inch (psi), and millimeters of mercury (mmHg). Some common pressure relationships are given below: Pascal 101,325 Pa = 1 atm Torricellis 760 Torr = 1 atm Millimeters of mercury 760 mmHg = 1 atm Pounds per square inch 14.7 psi = 1 atm The Ideal Gas Law is a reasonable approximation based on a combination of individual gas laws such as Charles's, Boyle's, and Avogadro's Laws: where R = 0.08206 (L·atm)/(mol·K) The ideal gas law is versatile due to the fact that the ideal gas constant, R, is constant no matter how any of the other variables are adjusted. This means that the initial and final states of any system can be related as follows: Gas Mixtures The ideal gas law can be applied to a single gas or mixture of gases. For a mixture of gases the total pressure is a result of the partial pressure of each component. The components in a mixture of ideal gases do not interact; therefore each of the individual partial pressures can simply be summed to attain the total pressure of the system, as shown by Dalton's Law of Partial Pressures: ⋯ Similar to relating initial and final states of a gas, one component can be compared to a total mixture of gases: Concentration, Density, and Molar Mass of Gases With manipulation, the ideal gas law can be used to solve for variables that don’t necessarily appear in the general gas law. The main theme to get down is what each variable represents, and how the equation can be rearranged in order to match it to what is desired. Concentration (C): The units of concentration, more specifically, molar concentration ( ); can easily be determined by simple rearrangement of the ideal in order to solve for. Density (d) and Molar Mass (M): Density, , can be determined by rearrangement of the ideal gas law utilizing the molar mass, , in order to get units of mass (grams) into the equation, or by simple unit analysis density can be determined by using the molar mass and concentration.  The Kinetic Molecular Theory is a simple model for the behavior of gases. There are three major assumptions made in this theory: 1. Particle size is negligible 2. Average particle kinetic energy is proportional to temperature 3. Particle collisions are completely elastic Some important relationships related to this theory include: kinetic energy is proportional to T in Kelvin the root mean square velocity of a particle is related to the temperature in K and molar mass of the particle Graham’s law of effusion the rate of gas movement is inversely proportional to the square root of its molar mass The effects of particle volume and particle intermolecular forces are accounted for in the following equation, known as the van der Waals equation. This equation can be used to calculate the properties of a gas under nonideal conditions. In this equation a and b are constants that are dependent on the gas which correct for the intermolecular forces and the volume of the particles, respectively. Colligative properties are properties that depend on the number of particles dissolved in a solution, not on the type of particle. Some common examples are: - vapor pressure lowering - freezing point depression - boiling point elevation - osmotic pressure Since colligative properties depend on the number of dissolved particles, the effect of electrolytes is different than the effect of nonelectrolytes. Electrolytes result in a larger number of particles when dissolved in solution; therefore have a greater effect on colligative properties. Each of the above properties can be estimated based on the amount of dissolved solute using the following equations: vapor pressure with nonvolatile nonelectrolyte solute ° Psolution = vapor pressure of solution Xsolvent = mole fraction of solvent Posolvent = vapor pressure of pure solvent vapor pressure with volatile nonelectrolyte solute ° ° change in freezing point with nonelectrolyte solute ∆ ΔTf = change in freezing point Kf = freezing point depression constant change in boiling point with nonelectrolyte solute ∆ ΔTb = change in boiling point Kb = boiling point elevation constant osmotic pressure with nonelectrolyte solute ∙ 0.08206 ∙ T For electrolyte solutions the van’t Hoff factor (i) must be considered: The modified equations for solutions with electrolyte solutes are as follows: ° ∙ ∆ ∙ ∙ ∆ ∙ ∙ ∙ ∙