UNIT IX: SOLUTION CHEMISTRY
UNIT 6: SOLUTION CHEMISTRY
– the study of chemical reactions that occur in solutions.
6.1 – The Nature of Solutions
Important Definitions:
Solution – homogeneous mixture
Solvent – component of solution in greater quantity
Solute – component of solution in lesser quantity
Soluble – solvent and solvent form a homogenous mixture
Insoluble – can’t dissolve
Saturated – a solution in which the solvent has dissolved as much solute as possible.
- in order for a solution to be saturated, some undissolved solid must be present
Unsaturated – solution in which more solute can still be dissolved
Solubility – the maximum amount of the solute which can dissolve in a given amount of solvent at a given temperature.
Solubility always requires 5 pieces of information
• Type of solute
• Amount of solute used
• Type of solvent
• Amount of solvent
• Temperature of the solution
Examples:
Page 194 #1-4
MOLECULAR POLARITY
Intermolecular forces – forces that exist between molecules, holding them together.
Dipole – a partial separation of charge which exists when one end of a molecule (or bond) is slightly positive while the other end is slightly negative.
Example:
Attraction of temporary dipoles between neighbouring molecules
← London Forces
Attraction of permanent dipoles between neighbouring molecules
← Dipole – dipole forces
Molecules with permanent dipoles ( POLAR
- Unequal sharing of electrons
- Asymmetrical
-
- __ melting point
- __ boiling point
Molecules without permanent dipoles( NON-POLAR
- Equal sharing of electrons
- Symmetrical
-
- __ melting point
- __ boiling point
**London forces all ALWAYS present – they just aren’t noticeable in ionic or polar molecules**
A bond between atoms with different electronegativities gives rise to a dipole.
Examples:
Pages 199-202 #9-12
HYDROGEN BONDING
- a relatively strong type of dipole – dipole attraction
- exists where an H atom covalently bonds to N, O, or F.
Examples:
Page 203 #13 – 16
6.2 – Dissolving
POLAR AND NONPOLAR SOLVENTS:
Look at the common solvents on pg. 204 and label them as either polar or nonpolar
|Solvent |Polar or nonpolar |Solvent |Polar or nonpolar |Solvent |Polar or nonpolar |
|Water | |Ethoxyethane | |Carbon tetrachloride | |
|Methanol | |Acetone | |Heptane | |
|Ethanol | |Acetic acid | |Liquid ammonia | |
|Benzene | |Chloroform | | | |
After completing many experiments concerned with MIXING polar and nonpolar solvents with polar and nonpolar solutes, results point to the following conclusions:
________ or ________ solutes dissolve in ________________
____________ solutes dissolve in ________________________
The “short” reason why…
• ________ and ________solutes have ________ bonds holding the solid together.
• ____________ solvents have _______________________and cannot exert enough energy to overcome the strong bonds.
• Only ________ solvents have sufficient____________ to the solute to be able to “pull” the solute out of the crystal and into the solution.
Therefore ________________can dissolve ________________.
• Nonpolar species ____ ____ possess ________and ________ ends.
o Therefore there is ____ __________to polar or ionic species.
• Only ________ solvents can attract ________ solutes, because they ________have _______ ____________ _____________.
Therefore ________________can dissolve ________________.
Read pages 205 – 206
Page 207#18-22, 23-26
6.3 – Dissociation Equations & Conductivity
THE CONDUCTIVITY OF AQUEOUS SOLUTIONS:
|How to Decide if a Substance will conduct electricity: |
| Is the substance a METAL? |If so, it conducts |
| Is the phase a SOLID? |If so, it doesn’t conduct |
|The following assume that the substance is a liquid or in aqueous solution |
|Is the substance an ACID/BASE? |If so, it conducts |
|Is the substance IONIC? |If so, it conducts |
|If none of the above |It doesn’t conduct |
Page 198 #6-8
Solvation – the interaction between a solute and a solvent
Ionic Solid – a crystalline solid made up of ions
Molecular Solid – a crystalline solid made up of neutral molecules.
Dissociation – separating previously existing ions in an ionic solid.
Example: NaCl(s) ( Na+(aq) + Cl-(aq)
Ionization – breaking up of a neutral molecule into ions.
Example: CH3COOH(l) ( CH3COO-(aq) + H+(aq)
Both reactions appear identical and both produce electrically conducting solutions.
Examples:
Show the dissociation of FeBr3 (s)
Show the ionization of HCN(s)
Show the ionization of K3PO4(s)
Page 210 #28a,c,e,g
CALCULATING THE CONCENTRATIONS OF IONS IN SOLUTION:
What is the molar concentration of the chloride ions in 0.25 M AlCl3?
What is the concentration of each type of ion in a solution made by mixing 50.0 mL of 0.240 M AlBr3 and 25.0 mL of 0.300 M CaBr2?
Page 212 #30 – 38
Predicting the Solubility of Salts:
• If a substance is _____________, then it has the ability to ____________ ______________
• A substance is said to be _____________ if it ______________ dissolve in water
In theory nothing is insoluble, everything can dissolve to some extent in water but it will dissolve so little that the concentration is negligible.
• If a compound ___________ _____________ in water then it is said to have LOW SOLUBILITY.
A substance is said to have LOW SOLUBILITY if a saturated solution of the substance is less than 0.1M.
We can determine if a substance is soluble or has low solubility by using the table called – Solubility of Common Compounds in Water
Example:
1. Determine whether FeCO3(s) is soluble.
1. Get out the solubility table
2. Find the negative ion, CO3-2 in the column
3. Find the positive ion, Fe+2, in the row,
a. Since Fe+2 is not listed in the rows, it will be in the “all others” category, the all others category has low solubility.
[pic]
Practice:
Label the following as either Soluble (S) or Low solubility (LS)
1. NaCl _________
2. Na2SO4 _________
3. FeCl3 _________
4. Ba(OH)2 _________
5. ZrSO4 _________
6. HCl _________
7. CrS _________
8. CuI _________
9. NaNO3 _________
• A _______________is a solid that forms in a solution when to aqueous ions react.
• A precipitation reaction or ________________ REACTION shows the ions reacting to form the solid. This is a type of ________________ reaction.
Write out the net ionic reaction for the following low solubility compounds.
1. AgCl (s) __________________________________
2. PbI2 (s) _________________________________
3. Mg(OH)2 (s) __________________________________
4. Ca3 (PO4)2 (s) __________________________________
Predicting if a Precipitate will form:
Will a precipitate (solid) form when solutions of CaS and Na2SO4 are mixed?
• These reactions will always be a double replacement reaction.
Ca+2 and S-2 will react with Na+1 and SO4-2 → CaSO4 and Na2S will be formed
• use the table to find if these compound have low solubility
Na2S will have/be ______________ (soluble or low solubility)
CaSO4 will have/be _____________ (soluble or low solubility)
The balanced equation including subscripts is
_____________________________________________________
The net ionic equation is: ____________________________________
Practice Questions:
1. An aqueous solution of Pb(NO3)2 is mixed with an aqueous solution of KBr
a) Write a balanced formula equation for this reaction. (Include all subscripts.)
b) net ionic equation is:
2. KNO3 + AlBr3 →______________________________
Net ionic equation is:
3. CaI2 + Sr(OH)2 → ________________________________
Net ionic equation is:
6-4: Dilution Calculations
M1V1 = M2V2
M1 – the original molarity M2 – New molarity after mixing
V1 – original volume V2 – new volume after mixing
Example
If 300.0 mL of 0.15M CaCl2 is added to 100.0mL of H2O, what is the new [CaCl2]?
How many litres of a 15.4M HNO3 solution is required to make 2.50L of a 0.375M solution?
If 200.0mL of 0.500 M NaCl is added to 300.0 mL of water, what is the resulting [NaCl] in the mixture?
Mixing 2 solutions with common ions:
If 300.0 mL of 0.250M NaCl is added to 500.0mL of 0.100M NaCl, what is the new [NaCl]?
450.0mL of 0.125M CaCl2 is mixed with 175.0mL of 0.761M CaCl2. What is the new concentration of CaCl2
Class Work - pg 102 # 78-82, 84, 87, 88, 90, 91
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