1. Reactions in Aqueous Solutions
General properties of aqueous solutions
Precipitation Reactions
Acid-Base Reactions
Oxidation-Reduction Reactions
Concentration of Solutions
Gravimetric Analysis
Acid-Base Titrations
Redox Titrations

2. General Properties of Aqueous Solutions: definitions
Solution: a homogeneous mixture of two or more substances
Solute: the substance present in a smaller amount
Solvent: the substance present in a larger amount
Aqueous Solution: the solute is initially a solid or a liquid and the solvent is water

3. Electrolytic Properties
Electrolyte: a substance whose dissolution in water results in a solution that conducts electricity
Non-electrolyte: a substance whose dissolution in water results in a solution that does not conduct electricity
Dissociation: when the compound is broken down into cations and anions
Hydration: when an ion is surrounded by water in a particular way
Ionization: when acids and bases separate into ions
Reversible Reaction: when a reaction can occur in both directions
Chemical equilibrium: when no net change in a reaction can be observed
Examples of Electrolytes and Non-electrolytes
Strong electrolytes: HCl, NaOH, and ionic compounds
Weak electrolytes: water (very weak), HF
Non-electrolytes: methanol, glucose and urea

4. Precipitation Reactions
A precipitate: the insoluble solid that forms from a reaction
Precipitate reactions form precipitates
Solubility: the maximum amount of solute that can dissolve in a given amount of solvent and a specific temperature
Insoluble compounds: carbonates, phosphates, chromates, sulfides (not when with compounds with alkali metal ions and ammonium ion)
Soluble compounds: compounds with alkali metal ions,ammonium ion, nitrates, bicarbonates, chlorates, halides (not with silver, mercury or lead) and sulfates (not with silver, calcium, strontonium, barium, mercury or lead)
Other insoluble compounds: hydroxides, except when with compounds with alkali metal ions and barium ion

5. Molecular, Ionic and Net Ionic Equations
Molecular equations: the formulas are written as if all species existed as molecules or whole units
Example: NaCl (aq) + AgNO (aq)  AgCl (s) + NaNO (aq)
Ionic equations: all the dissolved species are expressed as free ions
Example: Na (aq) + Cl (aq) + Ag (aq) + NO (aq) AgCl (s) + Na (aq) + NO (aq)
Net ionic equations: only the undissolved species are expressed on both sides of the reaction
Example: Ag (aq) + Cl (aq)  AgCl (s)

6. Acids and Bases Properties of Acids Sour taste
Color change of litmus paper blue to red
Reacts with some metals to produce hydrogen gas and with carbonates and bicarbonates to produce carbon dioxide
Aqueous acidic solutions conduct electricity
Bronsted Acid: proton donor
Properties of Bases
Bitter taste
Slippery feel
Color change of litmus paper red to blue
Aqueous basic solutions conduct electricity
Bronsted Base: proton acceptor

7. Acid-Base Neutralization
A reaction between a base and an acid is called a neutralization reaction.
The most common occurrence of this is when and acid and base yield water and a salt.
In an example we can use the most widely laboratory used acid and base: hydrochloric acid and sodium hydroxide. This reaction yields water and table salt (sodium chloride).
HCl +NaOH  NaCl + water

8. Oxidation-Reduction Reactions
Electron-transfer reactions are known as oxidation-reduction reactions or redox reactions
Each step is called a half-reaction where the loss/gain of electrons is clearly shown
An oxidizing agent is one that accepts electrons and the reducing agent is one that donates electrons

9. Oxidation numbers (or states)
Oxidation numbers or oxidation states are the number of charges the atom would have in a molecule, or an ionic compound, if electrons were transferred completely
Rules of oxidation numbers:
free elements, or uncombined elements, always have an oxidation state of zero
-monoatomic ions have an oxidation state equal to that of their charge
-for the most part oxygen has an oxidation state of -2, except for in hydrogen peroxide and peroxide where the oxidation state is -1
-usually the oxidation number of hydrogen is +1, except for when with binary compounds where it is -1
-the oxidation state of fluorine is always -1, for chlorine bromine and iodine the state is negative when as halide ions but positive when with oxygen
-polyatomic ions must have charges that add up to the net charge
-lastly, oxidation numbers do not have to be integers

10. Combination reactions are generally represented by A + B  C, in this type of reaction, two or more substances combine to form one product. Decomposition reactions are generally represented by C  A+B, here one substance decomposes into at least two simpler substances.

11. Displacement reactions
There are three different types of displacement reactions: hydrogen displacement, metal displacement and halogen displacement
In a hydrogen displacement reaction hydrogen is yielded as one of the products
metal is yielded in a metal displacement reaction
a halogen, such as fluorine, chlorine, bromine or iodine, is produced in a halogen displacement reaction

12. Disportionation reactions
Disportionation reactions have a unique quality, which is that an element is both oxidized and reduced.
In order for this to happen the element must have at least three oxidation states, one such element is oxygen.
When hydrogen peroxide is decomposed into water and oxygen three different oxidation numbers exist for oxygen in the reaction. When oxygen is part of hydrogen peroxide, its oxidation number is -1, in water its oxidation number is -2 and by itself as a free element its oxidation state is 0.

13. Concentration of a solution
The concentration of a solution depends on how much solute is present in a specific amount of solvent or solution.
For this section we will focus on one expression of concentration which is molarity (M), or molar concentration.
This is determined by how many moles of solute are in how many liters of solution M = mol / L

14. Obtaining desired molarity
We would multiply the desired molarity by the desired final volume to get the necessary number of moles.
Then we would have to convert moles to grams.
Using the final number of grams necessary we would place it all in a flask large enough to hold and measure the desired final volume
by adding water until we reach that final volume we would attain our desired molarity.

15. Dilution
dilution of a solution is making a concentrated solution less concentrated.
In order to do this we must realize that there is 1 mole of solute in 1 liter of solution.
We take a portion of the concentrated solution and dilute it enough to make the desired molarity.
Initial molarity multiplied by initial volume is equal to final molarity multiplied by final volume

16. Gravimetric Analysis
Gravimetric Analysis is a very accurate technique of measuring the amount of substance in a sample by measuring the mass of the formed precipitate, as long as the entire reaction is totally completed.
The precipitate also needs to be completely insoluble, otherwise part of the precipitate would not be solid, but aqueous and would have filtered through.

17. Acid-Base Titrations
During a titration a standard solution (a solution of accurately known concentration) is added gradually to another solution of unknown concentration, until the chemical reaction between the two solutions is complete.
Through titration we can find the exact concentration of the original solution of unknown concentration by calculations as long as we know the original volumes.
In acid-base titrations we must use indicators to let us know when the equivalence point has been reached, or in other words the titration is complete and the acid has been thus completely neutralized by the base.

18. Redox Titrations
Oxidation-Reduction titrations are very similar to acid-base titrations, only an equivalence point is reached when the reducing agent is completely oxidized.
Again a good indicator is necessary, and sometimes the oxidizing agent can act as an indicator because it changes color drastically when it is oxidized versus reduced.