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Parts of Solutions

• Solution- homogeneous mixture.• Solute- what gets dissolved.• Solvent- what does the dissolving.• Soluble- Can be dissolved.• Miscible- liquids dissolve in each other.

Figure 4.1 The Water Molecule

Hydration

• The process of breaking the ions of salts apart.

• Ions have charges and are attracted to the opposite charges on the water molecules.

Figure 4.2 Polar Water Molecules Interact with the Positive and Negative Ions of a Salt

Assisting in the Dissolving Process

How Ionic solids dissolve

Click here for Animation

Solubility

• How much of a substance will dissolve in a given amount of water.

• Usually g/100 mL• Varies greatly, but if they do dissolve the

ions are separated,• and they can move around.• Water can also dissolve non-ionic

compounds if they have polar bonds.

Figure 4.3a The Ethanol Molecule Contains a Polar O-H Bond Similar to Those in the

Water Molecule

Figure 4.3b The Polar Water Molecule Interacts Strongly with the Polar-O-H bond in

Ethanol

Electrolytes

• Electricity is moving charges.• The ions that are dissolved can move.• Solutions of ionic compounds can conduct

electricity.• Solutions are classified three ways.

Types of solutions

Strong electrolytes- completely dissociate (fall apart into ions).

a) Many ions- Conduct well.

Weak electrolytes- Partially fall apart into ions.

a) Few ions -Conduct electricity slightly.

Non-electrolytes- Don’t fall apart.a) No ions- Don’t conduct.

Figure 4.4a-c Electrical Conductivity of Aqueous Solutions

Figure 4.5 NaCl Dissolves

Figure 4.6 HCL is Completely Ionized

Figure 4.7 An Aqueous Solution of Sodium Hydroxide

Figure 4.8 Acetic Acid (HC2H3O2)

Figure 4.9 The Reaction of NH3 in Water

Measuring Solutions

• Concentration- how much is dissolved.• Molarity = Moles of solute

Liters of solution• abbreviated M• 1 M = 1 mol solute / 1 liter solution• Calculate the molarity of a solution with

34.6 g of NaCl dissolved in 125 mL of solution.

Figure 4.10a-c Steps Involved in the Preparation of a Standard Aqueous Solution

Figure 4.11a-b Measuring Pipets and Volumetric Pipets Measure Liquid Volume

Figure 4.12a-c A Measuring Pipet is Used to Add Acetic Solution to a Volumetric Flask

Figure 4.14 a&b Reactant Solutions

Figure 4.15 a&b The Reaction of K2CrO4 and Ba(NO3)2

Figure 4.17 Molecular-Level Representations Illustrating the Reaction of KCl (aq) with AgNO3 (aq) to Form AgCl (s)

Writing Net Ionic Equations

• Strong acids are all written in net ionic form-Binary acids – are all strong (except for HF(aq) ) -Oxyacids-If the number of oxygens exceeds the number of hydrogens by 2 or more they are considered strong.-Polyprotic acids ionize one (1) hydrogen at a time. All subsequent ionizations of acidic hydrogens are considered weak.(except for HSO4(aq) 1- )

H2SO4(aq) → H(aq) 1+ + HSO4(aq) 1-

• Strong bases are all written in ionic form.-Group IA and IIA metal hydroxides are strong bases.

• All weak acids and bases (those not mentioned above) are always written in molecular form.

• Ionic Salts-If soluble-written in ionic form -if insoluble written in molecular/undissociated form.

Writing Net Ionic Equations

• Oxides are always written in molecular/undissociated form.

• Gases are always written in molecular form.

• Molecular compounds are always written in molecular form.

Determining the Mass of Product Formed

Performing Calculations for Acid-Base Reactions

Neutralization Reactions I

Neutralization Reactions II

Neutralization Titration

Figure 4.19 The Reaction of Solid Sodium and Gaseous Chlorine to Form Solid Sodium

Chloride

Figure 4.20 A Summary of Oxidation-Reduction Process

The Half-Reaction Method (Acidic Solution)

The Half-Reaction Method (Basic Solution)

Figure 4.4a-c Electrical Conductivity of Aqueous Solutions

An Aqueous Solution of Co(NO3)2.

Figure 4.10 Steps Involved in the Preparation of a Standard Aqueous Solution

Figure 4.13 Yellow Aqueous Potassium

Figure 4.14a-b Reactant Solutions

Figure 4.15c Solution Post-Reaction

Figure 4.16 Addition of Silver Nitrate to Aqueous Solution of Potassium Chloride

Figure 4.17 Reaction of KCI(aq) with AgNO3(aq) to form AgCI(s).

Lead Sulfate

KOH and Fe(NO3)3 Mix to Create Solid Fe(OH)3.

Figure 4.18a-c The Titration of an Acid with a Base

Figure 4.19 The Reaction of Solid Sodium and Gaseous Chlorine to Form Solid Sodium

Chloride

Oxidation of Copper Metal by Nitric Acid

Magnetite

Aluminum and Iodine Mix to Form Aluminum Iodide

Chocolate

When Potassium Dichromate Reacts with Ethanol, the Solution Contains Cr3+.

Table 4.1 Simple Rules for the Solubility of Salts in Water

Table 4.2 Rules for Assigning Oxidation States

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