Solutions Ask a chemist, they always have

Solutes and Solvents  Solution: a homogenous mixture  Solute: thing that dissolves  Solvent: thing that does the dissolving (found in the largest amounts) If the solvent is water, then it is called an aqueous solution If the solvent is water, then it is called an aqueous solution

Solubility: Ionic Compounds  Ions form, separate (dissociate) and move throughout the solution The forces that hold the ions together are overcome by the ions’ attractions to polar water. The forces that hold the ions together are overcome by the ions’ attractions to polar water. Ion- dipole interactionIon- dipole interaction  Because ions are present, ionic solutions can conduct a current Current is just movement of electrons Current is just movement of electrons

Figure 15.1: Dissolving of solid sodium chloride. Solvation animation Animation with Audio

Figure 15.2: Polar water molecules interacting with positive and negative ions of a salt.

 “Like dissolves like”  Typically, hydrogen bonding occurs between the substance being dissolved and the polar water molecules  Example: Sugar in water Sugar in water Ethanol in water Ethanol in water Solubility: Polar Compounds

Figure 15.3: The polar water molecule interacts strongly with the polar O—H bond in ethanol.

Figure 15.4: Structure of common table sugar. Get interactions between water molecules the polar regions on the sugar (the Os), and some hydrogen bonding at the -OH groups

“They go together like oil and water.” (things that don’t dissolve or mix)  Anything nonpolar will not mix well with anything polar Examples: Oil spill Salad dressing  Can mix when shaken (LDF) and then may separate out (other forces)

How much is too much?  There is a limit to the amount of a substance dissolved Saturated : the solution holds as much solute as possible at that temperature. Saturated : the solution holds as much solute as possible at that temperature. Unsaturated : solution has not reached the limit Unsaturated : solution has not reached the limit

 Can you have too much? YES! Supersaturated: have as much solute dissolved as possible, then cooled and all the solute stays dissolved. Supersaturated: have as much solute dissolved as possible, then cooled and all the solute stays dissolved. In other words…the solution contains more dissolved solid than a saturated solution created at the same temperature. In other words…the solution contains more dissolved solid than a saturated solution created at the same temperature. These can begin crystallization of the solute at the slightest change These can begin crystallization of the solute at the slightest change

Energy changes and Solvation  Any chemical change (including solvation) requires a change in energy Energy removed from or added to the reactants from the surroundings Energy removed from or added to the reactants from the surroundings NaOH (s)  Na + (aq) + OH - (aq) ΔH= kJ/mol NaOH (s)  Na + (aq) + OH - (aq) ΔH= kJ/mol (that’s 44.5kJ released, so exothermic, per mole of NaOH) Because you are breaking the ionic bond, energy must be either released when breaking the bond, or consumed when making the new ionsBecause you are breaking the ionic bond, energy must be either released when breaking the bond, or consumed when making the new ions ALL changes in formula indicate a change in energy. ALL changes in formula indicate a change in energy. However, sometimes the energy change is so small, you can’t tell that a change has occurred However, sometimes the energy change is so small, you can’t tell that a change has occurred

Molarity ( M )  Most common way to express concentration  Molarity is the number of moles of solute dissolved in each liter of solution  Formula M = moles of solute M = moles of solute liters of solution liters of solution  Dependent on temperature  The higher the molarity the stronger the concentration

Molality ( m )  Another way to calculate concentration  Formula m = moles solute. m = moles solute. kilograms of solvent  Not dependent on temperature  The higher the molality the stronger the concentration

Normality ( N )  3rd way to calculate concentration  Typically used with acids and bases Indicates amount of H + and OH - available for acid/base reactions Indicates amount of H + and OH - available for acid/base reactions  Formula N = equivalents. N = equivalents. 1 liter of solution  Equivalent weight: the mass in grams of acid/base that gives 1 mole of H + /OH -  The higher the normality the stronger the concentration

Colligative properties  In the winter, why do we throw salt when it snows?  Why does Emeril add salt to boiling water when cooking pasta?  It has to do with Intermolecular Forces (IMFs)

Intramolecular forces (bonds)  Hold atoms together in molecules  Have high energy associated with them it’s difficult to break molecules into their individual atoms it’s difficult to break molecules into their individual atoms  Different types based upon what is going on with the electrons (electron clouds) Ionic: electrons are transferred between atoms Ionic: electrons are transferred between atoms Covalent Covalent electrons are shared between atomselectrons are shared between atoms Can be polar or nonpolarCan be polar or nonpolar

Intermolecular forces (IMFs)  Hold molecules together  Much weaker than intramolecular forces Intramolecular bonds are usually 100x or even 1000x stronger Intramolecular bonds are usually 100x or even 1000x stronger *(kJ are units of energy like Calories; 1Cal= 4.184kJ) 1000cal= 1Cal 1000cal= 1Cal 1cal =4.184J 1cal =4.184J

Figure 14.2: Intermolecular forces exist between molecules. Bonds exist within molecules.

Why do we care?  The strength of the IMFs determine the state of matter Solid, liquid, or gas* Solid, liquid, or gas* *Not plasma- intramolecular bonds are broken to get plasmas *Not plasma- intramolecular bonds are broken to get plasmas

shapevolumedensityenergy* motion with Energy* level of organization* strength of IMFs* Gas indefinite variable with P and T variable with volume change high high; molecules freely moving with great distance compared to molecular size between them very lowlow Liquid indefinite constant ** moderate high; molecules freely moving past each other but in close proximity to each other low moderate Solid definite constant ** low low; vibration only as molecules are basically fixed in place high * all at room temperature, ~25C **small variations occur due to temperature changes, very little variable with pressure changes Solids, Liquids, and Gases

 Things with strong IMFs tend to be solids at room temperature  Things with weak IMFs tend to be gases at room temperature  Medium IMFs tend to be in between- liquids, yes, but with varying characteristics liquids, yes, but with varying characteristics  Amorphous solids: long transition between solid and liquid states- gets soft, then melts (like wax)  Crystalline solids: definite, clear melting point (no soft transition- ie: ice)

Energy Changes Accompanying Changes of State   Think back: Each change of state is accompanied by a change in the energy of the system Whenever the change involves the disruption of intermolecular forces, energy must be supplied

Energy Changes Accompanying Changes of State   The disruption of intermolecular forces accompanies the state going towards a less ordered state (higher entropy) As the strengths of the intermolecular forces increase, greater amounts of energy are required to overcome them during a change in state Takes more energy to go from a liquid to a gas than from a solid to a liquid   Removing energy allows the molecules to “self- organize”, and results in an more ordered state Lower entropy

Heat of Fusion   The melting process for a solid is also referred to as fusion The enthalpy change associated with melting a solid is often called the heat of fusion (Δ H fus ) Ice ΔH fus = 6.01 kJ/mol Δ H is a change (Δ) in enthalpy (H), a measure of energy that is much like heat, but takes into account a few other factors

Heat of Vaporization   The heat needed for the vaporization of a liquid is called the heat of vaporization (Δ H vap ) Water Δ H vap = kJ/mol   Vaporization requires the input of heat energy

  Less energy is needed to allow molecules to move past each other than to separate them totally, so ΔH fus < Δ H vap

The heating/cooling curve for water heated or cooled at a constant rate.

 Think of IMFs like magnets: stronger magnets hold things more firmly together The more firm the connections, the less molecular motion can occur with the same amount of Energy added The more firm the connections, the less molecular motion can occur with the same amount of Energy added Adding (or removing) energy from the system can overcome (or increase) the IMFs, and cause a change in state Adding (or removing) energy from the system can overcome (or increase) the IMFs, and cause a change in state Add Energy, move from S -> L -> GAdd Energy, move from S -> L -> G Remove Energy, move from G -> L -> SRemove Energy, move from G -> L -> S

Vaporization and Vapor Pressure  The molecules in a sample of a liquid move at various speeds (average speed is the temperature; some have more energy, some have less, but the overall KE is temperature) (average speed is the temperature; some have more energy, some have less, but the overall KE is temperature)  Sometimes molecules at the surface have sufficient speed to overcome the attractive forces and leave the liquid surface (thus vaporizing)

Figure 14.9: Microscopic view of a liquid near its surface.

Dynamic equilibrium  Dynamic equilibrium is the state where there is simultaneous and equal vaporization and condensation of the substance  In a closed container, at some pressure, the amount that vaporizes will equal the amount condensing on the surface of the liquid This is the equilibrium vapor pressure This is the equilibrium vapor pressure

VP and IMFs  Stronger IMFs equal lower vapor pressures Less likely to evaporate Less likely to evaporate  Weaker IMFs equal higher vapor pressures Substance with very low IMFs and therefore high vapor pressures evaporate very quickly and easily Substance with very low IMFs and therefore high vapor pressures evaporate very quickly and easily Called volatile substanceCalled volatile substance  Mass and shape important, just like with boiling point Heavier = lower VP Heavier = lower VP ex: oilex: oil Lighter= higher VP Lighter= higher VP ex: alcoholex: alcohol More volatileMore volatile  Think propane (C 3 H 8 ) v. gasoline (C 8 H 18 )

VP and Boiling  Vaporization occurs at any temperature, but occurs more rapidly as temperature increases Molecules at the surface would have to have more speed to overcome the IMFs Molecules at the surface would have to have more speed to overcome the IMFs Boiling is the point at which the vapor pressure equals the external pressure on the surface of the liquid Boiling is the point at which the vapor pressure equals the external pressure on the surface of the liquid Molecules are able to “escape” liquid phase b/c they have enough Energy to break the IMFsMolecules are able to “escape” liquid phase b/c they have enough Energy to break the IMFs Convert PE of IMFs to KE of motion in a gas Convert PE of IMFs to KE of motion in a gas

Boiling and VP, con’t  Liquids have some air dissolved in them in tiny invisible bubbles  As water vaporizes in the liquid, it is added to the bubbles  Also, the gas bubbles are expanding because they are being heated; this causes an increase in volume, but not mass At this point, 2 things are going on: At this point, 2 things are going on: This decreases density, causing the bubbles to float to the surfaceThis decreases density, causing the bubbles to float to the surface Also, as gas expands, the pressure increasesAlso, as gas expands, the pressure increases When the pressure of the bubble increases to greater than the vapor pressure at the surface, the liquid is boiling  All molecules must be vaporized before a further increase in temperature can occur Need to break all IMFs (convert all PE of IMFs before increasing KE of molecules) Need to break all IMFs (convert all PE of IMFs before increasing KE of molecules)

Figure 15.10: Pure water.

Boiling Point and Elevation  As elevation on the Earth’s surface increases, the atmospheric pressure decreases (smaller column of air pushing down on the area; therefore less pressure) (smaller column of air pushing down on the area; therefore less pressure)  Boiling point changes as the atmospheric pressure changes  If you could decrease the pressure without changing temperature, the substance would boil at a lower temperature A decrease in pressure results in a decrease in BP A decrease in pressure results in a decrease in BP

Figure 14.14: The formation of the bubble is opposed by atmospheric pressure.

Boiling point elevation By adding salt (or other compounds) to water, the temperature of boiling goes up  it boils at a higher temperature By adding salt (or other compounds) to water, the temperature of boiling goes up  it boils at a higher temperature Interrupts H bondingInterrupts H bonding Need more vapor molecules and greater pressure to get bubbles to formNeed more vapor molecules and greater pressure to get bubbles to form Takes more time to get vapors to add to bubblesTakes more time to get vapors to add to bubbles The molecules that do get into the bubbles need more energyThe molecules that do get into the bubbles need more energy Dependent on how much solute is added Dependent on how much solute is added

Figure 15.9: A bubble in the interior of liquid water surrounded by solute particles and water molecules.

Figure 15.10: Solution (contains solute).

Vapor Pressure Reduction  Vapor pressure changes as IMFs change  For the same reasons boiling point is disturbed  What would evaporate faster: Salt water Salt water Distilled water Distilled water WHY? WHY?

Boiling Point Elevation Calculations  Water with salt added boils at a higher temperature than pure water. By how much will the boiling point change if 100.g of salt is added to 500. g of water? K b water = 0.52 o C/ m  Formula: T b = K b m i T b = K b m i  K b : Molal Boiling Point elevation constant ( o C/ m )  i= = Pieces that the material dissociates into (for ionic compounds only)  (Keep I at 1 (one) for covalent compounds)

Freezing Point Depression and Boiling Point Elevation SolventFormula Melting Point (°C) Boiling Point (°C) K f (°C/m ) ( K b (°C/ m) WaterH2OH2O Acetic acidHC 2 H 3 O BenzeneC6H6C6H CamphorC 10 H 16 O Carbon disulfideCS CyclohexaneC 6 H EthanolC 2 H 5 OH

Freezing point depression By adding salt (or other solutes) to water, the temperature of freezing drops  it freezes at a lower temperature By adding salt (or other solutes) to water, the temperature of freezing drops  it freezes at a lower temperature Because H bonding is disturbedBecause H bonding is disturbed Dependent on how much solute is addedDependent on how much solute is added

Freezing Point Depression Calcs  Antifreeze protects cars from freezing and overheating. Calculate the freezing point depression of a solution of 100. g of ethylene glycol (C 2 H 6 O 2 ) antifreeze in kg of water. K f water = 1.86 o C/ m  Formula: T f = K f m i T f = K f m i  K f : Molal Freezing Point depression constant ( o C/ m )  i= Pieces that the material dissociates into (for ionic compounds only)  (Keep I at 1 (one) for covalent compounds)

Freezing Point Depression and Boiling Point Elevation SolventFormula Melting Point (°C) Boiling Point (°C) K f (°C/m) K b (°C/m) WaterH2OH2O Acetic acidHC 2 H 3 O BenzeneC6H6C6H CamphorC 10 H 16 O Carbon disulfideCS CyclohexaneC 6 H EthanolC 2 H 5 OH