Homogeneous mixture. Solvent. Solute. Types of solutions. Concentration of solution. Ideal Solution. Non-ideal Solutions. Ionic Solutions. The attractions of water dipoles for ions pulls the ions out of the crystalline lattice and into aqueous solution.

1. Solutions and Enthalpy
Dr. K. Shahzad Baig
Memorial University of Newfoundland
(MUN)
Canada
Petrucci, et al. 2011. General Chemistry: Principles and Modern Applications. Pearson Canada Inc., Toronto, Ontario.
Tro, N.J. 2010. Principles of Chemistry. : a molecular approach. Pearson Education, Inc.

2. Some Terminology
Homogeneous mixture
Solvent
Solutes

3. Types of Solutions

4. Solution Concentration
Mass percent. (m/m)
Volume percent. (v/v)
Mass/volume percent (m/v)
Isotonic saline is prepared by dissolving 0.9 g
of NaCl in 100 mL of water and is said to be:
0.9% NaCl (mass/volume)

5. • Molarity: moles of solute/liter of solution
• Percent by mass: grams of solute/grams of solution (then multiplied by 100%)
• Percent by volume: milliliters of solute/milliliters of solution (then multiplied by 100%)
• Mass/volume percent: grams of solute/milliliters of solution (then multiplied by 100%)
𝑴𝒐𝒔𝒕 𝒄𝒐𝒏𝒄𝒆𝒏𝒕𝒓𝒂𝒕𝒊𝒐𝒏 𝒖𝒏𝒊𝒕𝒔 𝒂𝒓𝒆 𝒆𝒙𝒑𝒓𝒆𝒔𝒔𝒆𝒅 𝒂𝒔
=
𝐴𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
𝐴𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 𝑜𝑟 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
=

6. Very low solute concentrations
• Parts per million (ppm): grams of solute/grams of solution
(then multiplied by 106 or 1 million)
(µg/g, mg/L)
• Parts per billion (ppb): grams of solute/grams of solution
(then multiplied by 109 or 1 billion)
(ng/g, mg/L)
• Parts per trillion (ppt): grams of solute/grams of solution
(then multiplied by 1012 or 1 trillion)
(pg/g, ng/L)
ppm, ppb, ppt ordinarily are used when expressing extremely low concentrations
1.0 L @ 1.0 g/mL = 1000 g

7. Mole Fraction and Mole Percent

8. Molarity and Molality
Molarity varies with temperature (expansion or contraction of solution).
Molality is based on mass of solvent (not solution!) and is independent of temperature

9. Enthalpy of Solution
Solution formation can be considered to take place in three steps:
1. Move the molecules of solvent apart to make room for the solute
molecules. DH1 > 0 (endothermic)
2. Separate the molecules of solute to the distances found between them in
the solution. DH2 > 0 (endothermic)
3. Allow the separated solute and solvent molecules to mix randomly. DH3 <
0 (exothermic)
DHsoln = DH1 + DH2 + DH3

10. For dissolving
to occur,
the magnitudes
of ΔH1 + ΔH2
and of ΔH3
must be
roughly
comparable.

11. Intermolecular Forces in Ideal Solution
• An ideal solution exists when all intermolecular forces are of comparable strength,
DHsoln = 0.
• When solute–solvent intermolecular forces are somewhat stronger
than other intermolecular forces, DHsoln < 0.
• When solute–solvent intermolecular forces are somewhat weaker
than other intermolecular forces, DHsoln > 0.
• When solute–solvent intermolecular forces are much weaker than
other intermolecular forces, the solute does not dissolve in the
solvent.
Energy released by solute–solvent interactions is insufficient to separate solute particles
or solvent particles.

12. Non-ideal Solutions
In this solution, forces between ethanol and water are
stronger than other intermolecular forces
Adhesive forces greater than cohesive forces.

13. At the limit these solutions are heterogeneous
Adhesive forces are less stronger than cohesive forces

14. Ionic Solutions
The forces causing an ionic solid to dissolve in water are ion–dipole forces, the attraction
of water dipoles for cations and anions.
The extent to which an ionic solid dissolves in
water is determined largely by the competition
between:
– interionic attractions that hold ions in a
crystal, and
– ion–dipole attractions that pull ions into
solution.
The attractions of water dipoles for ions pulls the ions out of the crystalline lattice and
into aqueous solution.
Positive ends of dipoles attracted to anions.

15. Energy of Hydration