1. Solutions and Intermolecular Forces
Honors Chemistry
Mr. Kinton
Enloe High School

3. 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

4. Mixtures Homogeneous Heterogeneous Appear uniform throughout
Also known as solutions
Examples: air, brass, rubbing alcohol
Appear visibly different
The composition and properties are different
Examples: sand, rocks, wood

5. Mixtures
Classify the following as either an element, compound, heterogeneous mixture or solution:
Air
Iron
Carbon dioxide
Soil

6. 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

7. 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

8. 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

9. 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

10. Types of Intermolecular Forces

11. 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

12. 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

13. Types of Intermolecular Forces

14. 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

16. 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

18. 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

19. 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

20. 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

21. 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

22. 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

23. 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

25. 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

27. Energy Changes and Solutions
To determine the overall energy change in the solution formation process use the relationship
ΔHsoln = ΔHsolute + ΔHsolvent + ΔHmix
ΔHsolute: Corresponds to the separation of the solute
ΔHsolvent : Corresponds to the separation of the solvent
ΔHmix : Corresponds to the energy released by mixing

28. 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!

29. 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

30. 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

31. 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

32. 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

33. 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

34. 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

35. 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

37. 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?

38. 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

39. 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)

40. 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)

41. 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

42. Electrolytes
Identify the following as strong, weak, or non electrolytes:
H2SO4
CH3OH N2
LiOH
FeCl3

43. 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

44. Concentration
Calculate the molarity of a solution made by dissolving g of sodium sulfate in enough water to form a 125 mL solution.
What is the molar concentration of each ion present in a M aqueous solution of calcium nitrate?

45. 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)

46. Concentration
How many grams of Na2SO4 are required to make L of M Na2SO4?

47. 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

48. 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

49. 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?