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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