Solutions and Intermolecular Forces Honors Chemistry Mr. Kinton Enloe High School Classification of Matter Revisited Mixture: A combination of 2 or more substances and each substance retains its own chemical identity
Unlike pure substances, mixtures can be separated using physical means (filtration, distillation, chromatography) 2 different types of mixtures: homogeneous and heterogeneous Mixtures Homogeneous
Heterogeneous Appear uniform Appear visibly different The composition and throughout
Also known as solutions Examples: air, brass, rubbing alcohol properties are different Examples: sand, rocks, wood Mixtures
Classify the following as either an element, compound, heterogeneous mixture or solution: Air Iron Carbon dioxide Soil Matter Vocabulary Physical Property: measurable attribute of matter that will not change the identity or composition of matter
Chemical Property: used to describe how matter may change or react to form other substances Intensive Property: attribute that does not depend on the amount of substance present. Useful in identifying a substance Matter Vocabulary Extensive Property: attribute that depends on the
amount of matter present. Physical Change: a change that impacts the physical appearance but not the composition of the substance Chemical Change: a change that causes the substance to be transformed into a chemically different substance Intermolecular Forces
In chemical bonding we examined intramolecular forces, which focused on the forces between atoms or ions within a chemical compound (ionic, covalent, metallic) Intermolecular forces examines how the attractions between molecules influences their physical properties Intermolecular forces are weaker than intramolecular forces
Types of Intermolecular Forces London Dispersion Forces: caused by the movement of electrons As the electrons move the repel each other creating a momentary charge difference Weakest IMF and present in all molecules Polarizability: the ease with which the charge distribution can be distorted by an external electric field
Types of Intermolecular Forces Types of Intermolecular Forces The strength of London Dispersion forces tends to increase with increasing molecular size Dispersion forces also tend to increase for molecules that have a greater mass
Types of Intermolecular Forces Dipole-dipole: dipoles interact by the positive end of one polar molecule being attracted the negative end of another polar molecule This is similar to the concept of ionic bonding but much weaker The molecules MUST be polar to have dipole-dipole forces
Stronger than London Dispersion forces Types of Intermolecular Forces Types of Intermolecular Forces Dipole-dipole interactions increase as the molecules that are interacting with one another increase in polarity. The greater the dipole moment the stronger the dipole-dipole interactions will be
Types of Intermolecular Forces Hydrogen Bonding: special type of intermolecular attraction between the H atom in a polar bond and an unshared electron pair on a nearby F, O, or N atom Strongest IMF and will have a profound impact on the physical properties of your compound Intermolecular Forces and Physical
Properties Melting point and boiling point are 2 physical properties that are most influenced by intermolecular forces As the strength of the intermolecular force present increases so does the melting and boiling point of that substance Also influences that density of the substance: This helps explain why water in the liquid phase is more dense than
water in the solid phase Ion-Dipole Forces Interaction that exists between an ion and the partial charge on the end of a polar molecule Must contain an ion and polar molecule to have this intermolecular interaction Ion-dipole forces increase in strength as the charge
of the ion increases and as the dipole moment of the polar molecule increases Solutions, Suspensions, Colloids Solution: mixture of substances that has a uniform composition Suspension: mixture where the particles are large enough to be seen and given time will settle out in the mixture
Colloids: mixture containing particles that a larger than normal solutes but small enough to remain suspended in the mixture Solutions When a solution forms it is a physical change, the chemical identity of the substances involved is not being changed Solutions can form in the solid, liquid, or gas
phase Most solutions in chemistry occur in water and are in the aqueous phase Solutions Vocabulary Solute: substance dissolved in a solvent to form a solution. Normally present in a smaller amount Solvent: the dissolving medium of a solution. Normally present in a greater amount
Solvation: surrounding of the solute by the solvent Hydration: surrounding of the solute by water Solution Process A solution is formed when one substance disperses uniformly throughout another When a solution forms in the solid or liquid state we must account for the intermolecular attractive forces between the
molecules Intermolecular forces also act on solute particles and the solvent In order for a solution to form there must be attractive forces between your solute and solvent in the solution Energy and the Solution Process When a solution forms, 3 energy changes take place: Energy is required to separate the solute particles Energy is required to separate the solvent particles
Energy is released as solvation takes place Separation of the solute and solvent are both endothermic When solvation takes place, that process is exothermic Energy Changes and Solutions To determine the overall energy change in the solution formation process use the relationship HHsoln = HHsolute + HHsolvent + HHmix HHsolute: Corresponds to the separation of the solute
HHsolvent : Corresponds to the separation of the solvent HHmix : Corresponds to the energy released by mixing Factors the Influence Solution Formation The speed of solution formation can be impacted by 3 factors: Surface Area of the solute Stirring the mixture Heating the solvent
Do not confuse rate of solution formation with solubility! Solubility The amount of substance that can be dissolved in a solvent at a given temperature to form a saturated solution Saturated solution: dissolved and undissolved solute are in equilibrium
Unsaturated solution: solutions containing less solute than a saturated ones Supersaturated solution: solution containing more solute than a saturated one Solubility Crystallization: process in which a dissolved solute comes out of solution and forms a crystalline solid
Dynamic Equilibrium: state of where the forward and reverse processes are occurring at the same rate In a saturated solution the rate of dissolving is equal to the rate of crystallization Factors that Affect Solubility Pressure will only affect the solubility of a gas in a solvent. As pressure increases so does the solubility of the gas
Example: carbonated beverages Solids and liquids do not experience a change due to pressure due to how the molecules are arranged Factors that Affect Solubility Solute-Solvent interactions that are more similar in nature will lead to greater solubility like dissolves like
Miscible: liquids that mix in all proportions (acetone in water) Immiscible: liquids that do not dissolve in one another (gasoline in water Factors that Affect Solubility Temperature or the measure of the average kinetic energy of the molecules in a sample Solubility of solutes will increase as the temperature
of the solution increases This is because there is more energy available to break down the intermolecular forces within the molecule Solubility Curves A graph that will give the solubility of various substances in some amount of water at a given temperature
Any amount of solute below the line indicates an unsaturated solution Any amount of solute above the line in which all of the solute is dissolved is supersaturated Solubility Curves If the amount of solute is above the line but not all is dissolved, the solution is saturated
Any point that falls on a line represents where that solute is saturated at a given temperature If the amount of water changes it will influence how much solute can be dissolved Solubility Curves of Pure Substances g r a m s s o lu te p e r 1 0 0 g r a m s H 2 O
150 140 KI 130 120
110 NaNO3 100 90 KNO3
80 70 NH4Cl NH3 60
50 KCl 40 NaCl 30 20
KClO3 10 Ce2(SO4)3 0 0 10
20 30 40 50 Te m per atu re/Celsuis
60 70 80 90 100 Solubility Curve
At 10oC how many grams of KClO3 will dissolve? At 10oC, how many grams of KClO3 will dissolve in 50 mL of water? At 30oC, how many grams of NaCl will dissolve in 200 mL of water? Electrolytes
An aqueous solution that contains ions Electrolytes are able to conduct electricity to some extent Strong electrolytes completely separate into their individual ions in solution Weak electrolytes will only partially separate into individual ions Ionic compounds will dissociate and molecular compounds will ionize
Strong Electrolytes Exist almost completely as ions in solution All soluble ionic compounds are strong electrolytes HCl, HBr, HI, HClO3, HClO4, HNO3 and H2SO4 Group 1A metal hydroxides and Ca(OH)2, Sr(OH)2, and Ba(OH)2 Strong electrolytes will not revert back into their ionic or molecular state and will separate as follows: HCl(aq) H+(aq) + Cl-(aq)
Weak Electrolytes Solutions where a small portion of the solute exists in ion form Weak acids and weak bases are considered weak electrolytes When they form solution it occurs as follows: HC2H3O2(aq)
H+(aq) + C2H3O2-(aq) Non Electrolyte Substance that does not form ions in solution Any soluble molecular compound will be a non electrolyte Will not conduct electricity in water
Electrolytes Identify the following as strong, weak, or non electrolytes: H2SO4 CH3OH N2 LiOH FeCl3
Concentration The amount of solute dissolved in a given quantity of solution or solvent Molarity (M): unit of concentration expressed by the moles of solute in a liter of solution When examining the concentration of ions, it will depend on the number of ions in the chemical formula
Concentration Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate in enough water to form a 125 mL solution. What is the molar concentration of each ion present in a 0.025 M aqueous solution of calcium nitrate?
Concentration and Molarity Molarity serves as a conversion factor between volume of solution and the moles of solute We can use dimensional analysis to help us convert between any of these 3 quantities and even use that to do solution stoichiometry (next unit) Concentration
How many grams of Na2SO4 are required to make 0.350 L of 0.500 M Na2SO4? Dilution Process used to lower the concentration of a solution by adding water Stock solutions are what labs purchase and then chemists dilute the stock solution to a more manageable concentration
Dilution When doing a dilution with concentrated acids or bases always add the acid/base to water Adding water to concentrated acids/bases causes spattering due to the heat generated MconcVconc =MdilVdil
Dilution How many milliliters of 3.0 M H2SO4 are needed to make 450 mL of 0.10 M H2SO4? If 10.0 mL of a 10.0 M stock solution of NaOH is diluted to 250 mL, what is the concentration of the new solution?
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