Accelerated Chemistry

Chapter 13

Accelerated ChemistryChapter 13 notes

Chapter 13 Review: 4, 8-12, 20, 22, 23, 27, 28, 29, 30, 34, 39

13.1 Compounds in Aqueous Solution

A. Aqueous solution = A solution in which is the solvent ( ).

  1. Example: A solution of water (the ) and NaCl (the )
  2. Aqueous solutions can be electrolytes or non electrolytes.
  1. Electrolytes and non-electrolytes are solutes of solutions.
  2. Do electrolytes conduct electricity?
  3. Are non-electrolytes conductors?
  4. Are all electrolytes conductors?
  5. Are all conductors electrolytes?
CONDUCTORS
Pure Substance / Mixtures
Elements / Compounds / Alloys / Electrolytic Sol’n
All metals: Cu, Ag, Fe / All ionic compounds in liquid state: NaBr (I), KNO3 (I) / Stainless steel, Sterling silver / Water solutions of NaBr,HCl, NH3
(aqueous solution)
NONCONDUCTORS
Pure Substance / Mixtures
Elements / Compounds / Non-electrolytic Sol’n
All nonmetals: I2, P4 / All covalent compounds in liquid state: HBr (I), Al2Cl6 (I),
All solid compounds: Sucrose, NaBr (s), AlBr (s) / Water solutions of sucrose, isopropyl alcohol, ethyl alcohol, glycerin (aqueoussolutions)

B. Theory of Ionization

Theory of Ionization - Some water solutions conduct . These solutions are called . The formation of from solute molecules by the action of the .

Strong electrolytes:

NaCl(s) + H2O(l) 

HCl(g) + H2O(l) 

Weak electrolyte:

HC2H3O2(l) + H2O(l) 

A arrow shows a strong electrolyte fully producing . A yield sign indicates that the weak electrolyte does not fully .

In 1887, Svante Arrhenius (Sweden) proposed the theory of .

Some substances He based his ideas on observations of changes in and points with different molal concentrations

Note: Some reactions get a arrow ( ) and some get a arrow ( ).

Arrhenius proposed that when some chemicals are dissolved in , they produce particles with .

ionic compounds

NaCl(s) + H2O 

MgCl2(s) + H2O 

acids (covalent dissociation)

HCl(g) + H2O 

H2SO4(g) + H2O 

Substances that are not , , or do not dissociate/ionize.

C. Dissociation

The separation of ions that occurs when an ionic compound dissolves.

  1. A 1.0 M solution of sodium chloride contains:

1 mole of Na+ ions and 1 mole of Cl- ions.

NaCl(s) ------>

  1. A 1.0 M solution of calcium chloride contains:

1 mole of Ca+2 ions and 2 moles of Cl- ions – a total of 3 moles of ions.

CaCl2(s) ------>

  1. (a) Dissolve Al2(SO4)3 in water. (b) How many moles of aluminum ions and sulfate ions are produced by dissolving 1 mol of Al2(SO4)3. (c) What is the total number of moles of ions produced by dissolving 1 mol of Al2(SO4)3?

(a) Al2(SO4)3(s) ------>

(b) 1 mole ------>

(c) 2 moles Al3+ + 3moles SO42- = of solute ions

D. Solubility Equilibria:

No ionic compound has solubility.

No ionic compound has solubililty.

Rough rules of solubililty (using the solubility tables):

  1. If more than 1 gram per 100g H20 before saturation =
  1. If .1 gram to 1 gram per 100g H20 =
  2. If less than .1 gram per 100g H20 =
  1. Very slightly soluble ionic compounds – when placed in water, an equilibrium is established between the solid compound and its ions in solution:

Example:

AgCl(s)

Fe(OH)3(s)

Ag2S(s)

  1. Precipitation Reactions = Soluble compounds form products.

Type of rxn: - remember – reactants are soluble in water

  1. Net ionic equations – double replacement reactions and other reactions of ions in aqueous solutions are represented as ‘net ionic equations.’
Steps:

1. write an equation (molecular equation)

2. write total ionic equation

3. write a net ionic equation – only those compounds and ions that undergo a chemical change in a reaction in an aqueous sol’n and does not include spectator ions (ions found on the reactants and products side).

Ex1: Potassium Chloride + Silver nitrate

Molecular Equation:

Total Ionic Equation:

Net Equation:

Ex2:

Molecular Equation:

Total Ionic Equation:

Net Equation:

15.2 Colligative Properties of Solution

Molecular Electrolytes

A. Molecular solutes can form electrolytic solutions if they are highly polar.

Ionization versus dissociation
  1. Dissociation = The separation of that occurs
  1. Ionization = The formation of that occurs when a compound dissolves in water (water rips aparts molecules and turns them into ions

Ionization example:

H2O + HCl 

When a hydrogen chloride molecule ionizes in water, its hydrogen ion bonds covalently to a water molecule. A ion and a ion are formed.

Hydronium Ion:

The H+ ion attracts other molecules or ions so strongly that it does not normally exist, so the H+ ion becomes covalently bonded to oxygen.

Substances which form electrolytic solutions are:
Acids / HX / HCl, HNO3
Bases / MOH / NaOH
Salts / MX / NaCl, KBr, CaCO3

Why?

Which of the following form electrolytic solutions?

MgBr2C8H18KOHC12H22O11HNO3

Strong vs. weak electrolytes

Some compounds ionize/dissociate completely, while others don’t. ( )

Strong electrolyte – a compound that when dissolved/ionized, yields % ions.

Distinguishing factor of strong electrolytes – to whatever extent they dissolve in water, they yield only : HCl, HBr and HIare 100% ionized in dilute aqueous solutions.

Weak electrolyte – a solute that yields a relatively concentration of ions in an aqueous solution.

HF(aq) + H2O(l)

In an aqueous solution, the majority of HF molecules are present as dissolved HF s.

In general, the extent to which a solute ionizes in solution depends on the bonds within the molecules of the and the strength of attraction to molecules.

Note: If the strength of bonds in molecules < the attractive forces of the dipoles, then the covalent bonds break and the molecule separates into .

Properties of Electrolyte Solutions

Conductivity of Solutions

To compare the conductivities of strong and weak electrolytes, the conductivities of solutions of equal concentration must be compared.

Ionization of pure water:

H2O(l) + H2O(l)

So why does water that comes out of the tap conduct electricity?

It contains a high enough concentration of dissolved to make it a better conductor than pure water.

Colligative Properties of Electrolytic Solutions

Properties that depend on the concentration of the solute particles. Freezing point and boiling point are properties.

Freezing point depression – the difference between the freezing points of a pure solvent and a nonelectrolyte solution in it.

Solutions that conduct electricity contain .

Ionic compounds :

NaCl(s) + H2O(l) yields

MgCl2(s) + H2O(l) yields

Acids (dissociation of a covalent compound):

HCl(g) + H2O(l) yields

H2SO4(l) + H2O(l) yields

Substances that are not acids, bases, and salts do not .

When solutes dissolve in liquids, they the freezing point.

Two factors affect the degree of change in the temperature: the amount of the and the nature of the .

∆tf = kf (m)(x)x = # of produced when the solute dissolves

kf water = -1.86 oC/m

As the number of solute particles increase, the freezing point .

Ex1: Calculate the freezing point of 10.00 grams of NaCl in 200.0 grams of water.

Boiling point elevation – when solutes dissolve in liquids, they the boiling points.

Same concept as freezing point depression except boiling point .

kb water = 0.512 oC/m

Why does boiling point elevation occur?

The solute takes up space on the of a liquid. This decreases the ability of the liquid to . Thus, the vapor pressure . Boiling occurs when the atmospheric pressure the vapor pressure. So, an in energy is needed to increase the vapor pressure to reach the atmospheric pressure.

= solvent

versus

= solute

“A”“B”

Which would produce more vapor?

Which would have a higher vapor pressure?

Which would take less energy to raise the vapor pressure to atmospheric pressure?

Which would have a higher boiling point?

Ex1: Calculate the boiling point of a solution of 10.00 grams of NaCl in 200.0 grams of water.

Ch3 Notes - S Page 1 of 9