Chapter 7

Reactions and Solutions

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7.1 Writing Chemical Reactions

• We will learn to identify the following patterns of chemical reactions:– combination

– decomposition

– single replacement

– double replacement

• Recognizing the pattern will help you write and understand reactions.

Combination Reactions

• The joining of two or more elements or compounds.

A + B AB

• Example: MgO(s) + CO2(g) MgCO3(s)

Decomposition Reactions

• produce two or more products from a single reactant

AB A + B

• Example: CaCO3(s) CaO(s) + CO2(g)

Replacement Reactions (2 types)

1. Single replacement

A + BC B + AC

• Example:

Cu(s)+2AgNO3(aq) 2Ag(s)+Cu(NO3)2(aq)

2. Double Replacement

AB + CD AD + CB

• Example:

HCl(aq)+NaOH(aq) NaCl(aq)+H2O(l)

7.2 Types of Chemical ReactionsPrecipitation Reactions

• Chemical change in a solution that results in one or more insoluble products.

• To predict if a precipitation reaction can occur it is helpful to know the solubilitie of ionic compounds (Table 7.1).

2

Predicting Whether Precipitation Will Occur

•Have the ionic compounds exchange partners.

•Look at the new compounds found and determine if any are insoluble according to the rules (Table 7.1)

•The insoluble salt will be the precipitate.

•Example: Pb(NO3)2(aq) + NaCl(aq) PbCl2(s) + NaNO3(aq)

Reactions with Oxygen

•Reactions with oxygen generally release energy.

• Example: combustion of natural gas

•Organic compounds (carbon containing cp.) CO2 and H2O usually the products.

• Example: CH4+2O2CO2+2H2O

• Rusting (corrosion of iron) is an inorganic example.

4Fe + 3O2 2Fe2O3

Acid-Base Reactions

• These reactions involve the transfer of a hydrogen ion (H+) from one reactant to another.

• Example:

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

The hydrogen ion on HCl was transferred to the oxygen in OH-, giving H2O

Oxidation-Reduction Reactions

• Reaction involves the transfer of one or more electrons from one reactant to another.

• Example:

Zn(s) + Cu2+(aq) Cu(s) + Zn2+(aq)Two electrons are transferred from Zn to Cu2+

7.3 Properties of Solutions

• Solution - homogeneous mixture• Solute – that which is dissolved, the substance

present in the lesser quantity• Solvent – that which does the dissolving, the

substance present in the largest quantity• Aqueous Solution - solution where the solvent is

water• Solutions can be liquids as well as solids and gases.

General Properties of Liquid Solutions

• Clear, though may have color

• Remember that electrolytes are compounds that dissociate when they dissolve in water.

)(-Cl )(Na )NaCl( OH2 aqaqs

• Nonelectrolytes do not dissociate.

• Volumes are not additive.

–1 L of ethanol + 1 L water does not give 2 L of solution.

• Solutions vs. Colloids vs. Suspensions

4

• Colloidal suspension - contains solute particles which are not uniformly distributed.

– Due to larger size of particles (1nm - 200 nm)

– Smaller than 1 nm, have solution.

– Larger than 200 nm, have a precipitate.

• Tyndall effect - the ability of a colloidal suspension to scatter light.

– See this as a haze when shining light through the mixture.

– Solutions: light passes right through without scattering.

• Suspension:– mixture contains particles much larger than a colloidal

suspension

– particles easily settle out and form a precipitate

Degree of Solubility

• Solubility - how much solute can dissolve in a solvent.

• Factors which affect solubility:

1 Polarity of solute and solvent

• The more they are different, the less soluble.

2 Temperature

• increase in temp. usually increases solubility.

3 Pressure

• for gas solutes, increased pressure of gas increases solubility

• Unsaturated solution - does not contain all the solute that can be dissolved at a particular temperature.

• Saturated solution - contains all the solute that can be dissolved at a particular temperature.

• Supersaturated solution - contains more solute than can be dissolved at that temperature.

• How is this done? • Heat solvent, saturate it with solute then cool slowly.

• Sometimes the excess will precipitate out.

• If it doesn’t precipitate, the solution will be supersaturated.

Solubility of Gases: Henry’s Law

•Henry’s Law - the number of moles of a gas dissolved in a liquid is proportional to the partial pressure of the gas above the liquid. (At constant temperature)

•Solubility and Equilibrium – a saturated solution is an example of a dynamic equilibrium.

7.4 Concentration of Solutions: Percentage

• Concentration - amount of solute dissolved in a given amount of solution. 6

solution ofamount

solute ofamount ion concentrat

• We will learn to express concentration in 3 ways:

Weight/Volume Percent, Weight/Weight Percent and Molarity

Weight/Volume Percent

%100solution of smilliliter

solute of grams

V

W%

Weight/Weight Percent

%100solutions grams

solute grams

W

W%

Used most often with solid solutions.

Note: there is an error with the above formula on page 186 of your book.

7.5 Concentration of Solutions: Moles and Equivalents

Molarity:

solution L

solute molesM

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• A dilution is required if..

– you have a solution in stock and wish to make a solution of less concentration.

– Simply add the correct amount of water.

• Since:

• Then: M x V = moles of solutesolution L

solute molesM

• In a dilution will the moles of solute change? No.

M1V1 = M2V2

This is the dilution formula, DO NOT USE IT FOR ANYTHING ELSE!

• You may use %W/V concentration units as well

• So, since M x V=moles

Representation of Concentration of Ions in Solution 9

• Two common ways of expressing concentration of ions in solution:

1. Moles per liter (molarity)

2. Equivalents per liter (emphasizes charge) We will not cover this.

• Let’s look at Na3PO4.

• If we dissolved 1 mol of sodium phosphate into 1 L of water, what would be the concentration (M) of the solution? 1 M

• Now let’s consider the individual ions. What would the concentration of Na+ be? 3 M

• What would the concentration of PO43- ions be? 1 M

• Equivalent is defined by the charge. One way of looking at it: it is the mole of charges:

• 1 mol Na+ = 1 equivalent Na+

• 1 mol PO43- = 3 equivalents PO4

3-

• Okay, back to the 1 M Na3PO4 solution. What are the equivalents of Na+?

• 3 mol Na+ = 3 equivalents of Na+.

• What are the equivalents of PO43-?

• 1 mol PO43- = 3 equivalents of PO4

3-.

• Do you see where the word equivalent might come form?

• So the following are things we can say about a 1 M solution of Na3PO4.

3 mol/L Na+1 mol/L PO43-

3 equivalents/L Na+3 equivalents/L PO43-

• One Equivalent of an ion is the number of grams of the ion corresponding to Avogadro’s number of electrical charges.

ionon charges ofnumber

(g)ion of massmolar ionan of equivalent One

7.6 Concentration-Dependent Solution Properties

• Colligative properties- properties of solutions that depend on the concentration of the solute particles, rather than the identity of the solute.

• There are four:– vapor pressure lowering

– boiling point elevation

– freezing point depression

– osmotic pressure

Vapor Pressure Lowering

• Raoult’s Law - when a nonvolatile solute is added to a solvent, the vapor pressure of the solvent decreases in proportion to the concentration of the solute.

Freezing Point Depression and Boiling Point Elevation

• Before we can discuss the above colligative properties we must learn a new concentration unit.

• Molality (m) = moles of solute per kg of solvent.

solvent kg

solute moles Molality

• Note: the denominator is in kg solvent, not in kg solution.

Freezing point depression (Tf) is proportional to the

number of solute particles

Tf=kf m

freezing point depression constantEach solvent has it’s own value.

• Note, it is solute particles not just solute.• If the solute is an electrolyte, what will it do? It will

break apart into ions.• Therefore the same concentration of NaCl will affect the

freezing point twice as much as glucose (a nonelectrolyte)

Boiling point elevation (Tb) is proportional to the

number of solute particles

Tb=kb m

boiling point elevation constantEach solvent has it’s own value.

• Again, an electrolyte will affect the boiling point to a greater degree than a nonelectrolyte with the same concentration.

Osmotic Pressure

• Semipermeable membranes- allow solvent but not solute to diffuse from one side to another.

• Osmosis - the movement of solvent from a dilute solution to a more concentrated solution through the membrane.

• Osmotic pressure (Π) - the amount of pressure required to stop the flow.

• Again, M is molarity of particles. Called osmolarity.

=MRT

Isotonic solution = both solutions have the same concentration of solutes and are in equilibrium

Hypotonic solution = a solution has a lesser concentration of solutes than another across a membrane

Hypertonic solution = a solution has a greater concentration of solutes than another across a semipermeable membrane

7.7 Water as a Solvent• Water often referred to as the “universal solvent”.

• Most abundant liquid on earth.

• ~ 80% of the human body is water.– transports ions, nutrients and waste into and out of cells.

– solvent for biochemical reactions in cells and digestive tract

– reactant or product in some biochemical processes.

• Active transport: Cellular energy must be expended to make concentration of ions different on each side of the cell membrane.– This is accomplished via large protein molecules embedded in cell

membranes.

CATIONS IN THE BLOOD and CELLS

• Na+ and K+ two most important cations.

• Danger to body occurs when Na+ and K+ in blood and in cells becomes too high or low.

- Na+ too low: - Decrease of urine output

- Drymouth, flushed skin, fever

- Na+ too high:- Confusion, stupor or coma

- K+ too high or too low:- Death by heart failure

ANIONS IN THE BLOOD

• Cl- - acid/base balance- maintenance of osmotic pressure- oxygen transport by hemoglobin

• HCO3-

- Form in which most waste CO2 is carried out of the body.

PROTEINS IN THE BLOOD

• Blood clotting factors

• antibodies

• albumins (carriers of non-polar substances which cannot dissolve in water)

• Proteins are transported as a colloidal suspension.

• The blood also transports nutrients and waste products.