Reading Notes for Chapter 9


These are Dr. Bodwin's reading notes for Chapter 9 of "Introduction to Chemistry". I am using a local .pdf copy that was downloaded in August 2020.

Chapter Summary:

It's important to understand the behaviour of pure substances, but most of the things we encounter and use every day are mixtures. In some cases (Ideal gases), mixtures behave pretty much like pure substances, but in most cases, mixtures have vastly different properties.

Solution Definitions:

Solution = a homogeneous mixture of 2 or more substances.
Solute(s) = the minor component(s) of a solution
Solvent(s) = the major component(s) of a solution
If you visualize a "solution", you probably think about dissolving a solid solute in a liquid solvent, like adding table salt (NaCl) to water. Notice that there is nothing in the definition of solution that specifies states of matter.

Solubility:

Substances with similar intermolecular forces (IMFs) tend to interact with each other.
"Soluble" vs. "miscible" - these terms kind of mean the same thing, but they are usually applied to different situations. If we're adding a small amount of solute to a large amount of solvent and it dissolves, we typically call that "soluble". If I have 2 different samples (usually liquids or solutions) that combine when mixed, then we typically call that "miscible".
Review solubility rules and trends from when we looked at precipitation reactions.
Solubility is not a "yes" or "no". Sodium chloride is soluble in water. Can we dissolve a bucket of salt in a cup of water?
Saturated solutions are solutions that are at their solubility limit.

Concentration:

Concentration is a way to describe the relative amounts of substances in a mixture. There are a lot of different units that can be used to express concentrations, and they are all used in different settings and different applications, but remember that all concentration units are just expressing the relative amounts of different components of a mixture.

Fraction-based Concentration Units:
This is the general term for a concentration unit that relates the amount of a component to the total amount of solution.
Mass-Mass fraction = mass of solute divided by total mass of solution. The masses must be in the same unit.
Mass-Volume fraction = mass of solute divided by total volume of solution. These are usually expressed as "grams per milliliter", but always read carefully!
Volume-Volume fraction = volume of solute divided by total volume of solution. The volumes must be in the same unit.
Mass-Mass and Volume-Volume have the advantage that it doesn't matter what the mass or volume unit is, the answer will be the same as long as the same unit is used for each.
If you multiply a fractional concentration by 100, you get a percent. ("per cent" is from "parts per hundred"...)
For very small concentrations, you can multiply the fractional concentration by 1,000,000 to get parts per million, ppm. (ppm video)
For very, very small concentrations, you can multiply the fractional concentration by 1,000,000,000 to get parts per billion, ppb.
Fractional and percent concentration unitss are very common in healthcare settings.

Molarity:
In many cases, we need to know the number of moles of a substance in a solution. We might be able to get that from the mass in a fraction-based concentration unit, but because it's so common, we have a concentration unit that directly gets to moles.
Molarity, M = moles of solute divided by liters of solution.
(moles from concentration video)

Dilution:

Dilutions happen all the time, don't make them harder than they are.

For just about any dilution, we can use "C1V1=C2V2" where C is concentration and V is volume.
(dilution video)

Colligative Properties:

Colligative properties rely upon the number of solute particles in solution. Counting solution particles requires an understanding of how different things dissolve.
  1. Molecular solutes - Molecular solutes remain in 1 piece when they dissolve. Sugar, glucose, ethylene glycol...
  2. Ionic solutes - break into their ions. The number of ions is the number of particles. DO NOT break apart polyatomic ions. Sodium chloride (NaCl) dissolves to form 1 sodium ion and 1 chloride ion (2 particles per formula unit). Sodium nitrate (NaNO3) dissolves to form 1 sodium ion and 1 nitrate ion (2 particles per formula unit). Potassium sulfate (K2SO4) dissolves to form 2 potassium ions and 1 sulfate ion (3 particles per formula unit).
Vapor pressure decreases when a solute is added.
Freezing point decreases when a solute is added.
Boiling point increases when a solute is added.

Osmosis is one of the most important forces in biological systems. Solvent passes through a semi-permeable membrane from the side of (relatively) low concentration to the side of (relatively) high concentration. Osmosis occurs in an attempt to equalize concentrations across a semi-permeable membrane.
Osmotic pressure is the (theoretical) pressure required to stop osmosis.
(https://courses.lumenlearning.com/introchem/chapter/osmotic-pressure/)
The "i" term your textbook uses to describe osmolarity is just the number of solute particles per solute formula unit, just like what's described above for ionic solutes. In different settings, this is called the "van't Hoff factor".

If a solution is forced through a semipermeable membrane, this process is called reverse osmosis.

Colligative properties can be calculated... for this course we will just talk about direction and extent of change. For example: Which solution would you expect to have the lower freezing point, 1.28M NaCl(aq) or 1.73M NaCl(aq)? {It's 1.73M NaCl(aq) because there are more solute particles per unit volume.}



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