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 existswhen 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.
  • 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:
  1. Is the substance a METAL?
/ If so, it conducts
  1. Is the phase a SOLID?
/ If so, it doesn’t conduct
The following assume that the substance is a liquid or in aqueous solution
  1. Is the substance an ACID/BASE?
/ If so, it conducts
  1. Is the substance IONIC?
/ If so, it conducts
  1. 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,
  4. Since Fe+2 is not listed in the rows, it will be in the “all others” category, the all others category has low solubility.

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:

  1. KNO3 + AlBr3 →______

Net ionic equation is:

  1. 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