This module expands upon the last by continuing our discussion of chemical equations. This material is covered in Chapter 4 of Brown and introduces the basic chemistry of aqueous chemical solutions and then focuses on three aqueous reactions: precipitation, neutralization and reduction-oxidation. Please bring your lab manual to class as I'll be referring to the periodic table frequently.
Topics for Lecture 4: [Links to slide sets]
1. Ionic compounds (except for precipitates) dissociate in water, forming ions.
2. Most aqueous solution chemistry occurs via exchange reactions.
a. Electrolytes are conductive solutions of ions.
b. A solution is a homogenous mixture of solute in solvent.
c. Molarity is a measurement of solution concentration (moles solvent/L solute).
d. Precipitates (solid & insoluble ionic compounds) form because the attractive forces of a few ionic
compounds are greater than the attractive force of water for the ions of those compounds.
e. Most neutralization reactions produce water and a salt, and have the net ionic equation: proton +
hydroxide → water. If the base has CO3 or S as an anion, the water product of neutralization is
replaced by a gas.
f. Complete and net ionic equations show the net effect of chemical reactions; show which atoms change
connectivity. Solids, water and gases do not dissociate.
g. Reduction and oxidation are linked and due to the transfer of electrons. Oxidation is loss and
reduction is gain.
h. Oxidation of metals by acids or salts occurs by displacement reaction.
i. Solutes dissolve or dissociate in solvents in a process called solvation.
ii. Salts and strong acids are strong electrolytes, while weak acids are weak electrolytes.
iii. Molecular compounds are diluted, rather than dissociated, and are not electrolytes.
iv. The ratio of ions produced by dissociation is dictated by molecular formula.
v. Conversion factors are used to convert mass to moles to molecules to atoms to volume and back:
(g/mol), (6.02 x 1023 molecules or atoms/mole), (moles/L) and formula subscripts to convert
molecules to atoms and vice versa.
vi. Dilution with solvent decreases the concentration of solute, but solute molecules do not dissociate.
vii. The solubility guidelines (precipitation chart) are used to determine which compounds precipitate
(are insoluble) and are organized by anion.
viii. Acids can dissolve precipitates.
ix. Acids donate protons and lower pH while bases accept protons and increase pH.
x. Acids can be mono-, di- or triprotic and bases can be mono-, di-, or tribasic.
xi. There are seven strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO3, HClO4.
xii. Conjugate acid-base pairs are formed when acids lose protons and/or when bases gain protons.
xiii. Titration is a technique used to determine the concentration of one reactant. It requires a balanced
chemical equation, one reactant of known concentration, and a visible or discernible endpoint.
xiv. Spontaneous redox reactions cause corrosion and occurs at electrodes.
xv. Oxidation numbers are assigned to each atom in a redox equation. If an atom’s oxidation number
increases from reactant to product that atom is oxidized. If an atom’s oxidation number decreases
from reactant to product, that atom is reduced.
xvi. The activity series shows the ease of oxidation of metals including hydrogen. In order to oxidize
elemental metals, ions must be located below the elemental metal in the activity series.
xvii. Precious metals are prized because they resist oxidation (corrosion) and endure through time.
Links & items of interest:
Resources for students: