Chapter 7. Reactions and Solution

Chapter 7. Reactions and Solution

Writing Chemical Reactions

Chemical reactions may be classified as Combination, Decomposition, or Replacement.

Combination Reactions

The combination of a metal and a non-metal to form a salt, for example,

Cu(s) + Cl2(g)  CuCl2(s)

4Li(s) + O2(g)  2Li2O(s)

3Ca(s) + N2(g)Ca3N2(s)

2Al(s) + 3S(s)  Al2S3(s)

Decomposition Reactions

Essentially, decomposition reaction are the opposite of combination reactions. A compound decomposes (i.e.,"splits-up") into two or more compounds and/or elements. For example mercury(II) oxide will, upon heating, decompose into mercury metal and oxygen:

2HgO(s) 2Hg(s) + O2(g)

2CuO(s) 2Cu(s) + O2(g)

CaCO3(s) CaO(s) + CO2(g)

Replacement Reactions

Replacement reactions are subclassified as either :

SingleReplacement reactions
DoubleReplacement (Substitution) reactions

Single replacement reaction:

Cationic (metal) single replacement: When a piece of copper placed into a silver nitrate solution, the solution begins to turn blue and the copper seems to disappear. Instead, a silvery-white material appears.

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

Anionic (non-metal) single displacement reaction: When chlorine gas is bubbled into a potassium iodide solution, the chlorine is used up and the solution turns purple-brown from the iodine. This is an example of an anionic single replacement reaction.

2 KI (aq) + Cl2(g)  2 KCl (aq) + I2(aq)

Types of Chemical Reactions

Reactions that produce products with similar characteristics are often classified as a single group. For example, the formation of a precipitate denotes precipitation reactions. Chemical reactions that have a common reactant may be grouped together. Reactions involving oxygen - combustion reactions - are such a class.

Another approach to the classification of chemical reactions is based on charge transfer. Acid-base reactions involve proton (H+) transfer and oxidation-reduction reactions involve the transfer of one or more electrons

Precipitation Reactions

Reaction of two soluble substances to form an insoluble precipitate, for example,

HCl(aq) + AgNO3(aq)  AgCl(s) + HNO3(aq)

Precipitation reactions are Double Replacement Reactions.

A typical double replacement reaction can occur when two ionic compounds are mixed together.In water these ionic compounds split apart into their respective anions and cations. The cations now have an opportunity to swap anions. A reaction occurs if, by swapping anions, a product is formed that cannot split apart into anions and cations.

For example: the reaction of silver nitrate and potassium chloride. The silver, nitrate, potassium and chloride ions all begin in solution. When a silver ion combines with a chloride ion, it leaves the solution and becomes a solid. This drives the reaction to completion!

AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)

The reaction must be driven by one of three driving forces: 1)Formation of a solid,2) formation of a gas or 3) formation of a weaklyionizing compound such as water.For example, when you mix NaNO3(aq)and KCl(aq), if there was no insoluble salt (NaCl(aq) + KNO3(aq) ) formed, there would not be a precipitation reaction.

NaNO3(aq) + KCl(aq)  NaCl(aq)+ KNO3(aq)

Solubility of ionic compound in water

There are a series of guidelines in your book. On the next page, I provide you with another presentation of these same rules:

Solubility Rules

1. All compounds containing Na+, K+, or NH4+ ions are soluble in water.

2. All nitrates (NO3) are soluble in water.

3. Most chlorides (Cl), and sulfates (SO42) are soluble. Some important exceptions are silver chloride (AgCl), barium sulfate (BaSO4), and lead sulfate (PbSO4) which are insoluble.

4. Most carbonates (CO32), phosphates (PO43), sulfides (S2), and hydroxides (OH) are insoluble in water. Important exceptions are those of Na+, K+,? and NH4+, as well as barium hydroxide, Ba(OH)2.

5. How do you know if a gas will form? If H2CO3 is formed it will form a gas by decomposing to H2O and CO2. You will see that in this. Other compounds that release water or a gas areNH4OH and H2SO3.

E.g.

i) AgNO3 (aq) + NaCl (aq) = AgCl (s) + NaNO3 (aq) Precipitation Reaction occurs.

ii) NaCl (aq) + KNO3 (aq) = NaNO3 (aq) + KCl (aq) No precipitation Reaction.

Gas Forming Reactions

Gas is removed from the reaction mixture completing reaction to products. Some salts react with acids to form gases. Carbonates and bicarbonates produce CO2.

It works with any carbonate or bicarbonate. It is actually a two-part reaction.

CaCO3(s) + 2 HCl(aq) ŕ CaCl2(aq) + H2CO3(aq)

H2CO3(aq) ŕ CO2(g) + H2O(l)

Sulfites and bisulfites produce SO2.

Sulfides produce H2S.

Halides produce hydrogen halides (with nonvolatile acids such as H2SO4 and H3PO4).

Reactions with Oxygen (Combustion Reactions)

When a substance combines with oxygen releasing a large amount of energy in the form of light and heat, it is a combustion reaction. An example of a combustion reaction is:

C3H8 + 5O2 3CO2 + 4H2O

Acid-Base Reactions

(Formation of a weaklyionizing compound such as water)

When an acid and a base are placed together, they react to neutralize the acid and base properties, producing a salt. The H+ cation of the acid combines with the OH- anion of the base to form water. The compound formed by the cation of the base and the anion of the acid is called a salt. The combination of hydrochloric acid and sodium hydroxide produces common table salt, NaCl:

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

The word salt is a general term which applies to the products of all such acid-base reactions.

Salts: AgNO3, NaCl, NaNO3, KNO3

Names formulas of acids / Names and formulas of bases
HCl hydrochloric acid
HNO3 -nitric acid
HNO2 -nitrous acid
H2SO4 - sulfuric acid
H2SO3 -sulfurous acid
H3PO4-phosphoric acid
HC2H3O2 acetic acid / HClO4 perchloric acid.
HClO3 chloric acid
HClO2 chlorous acid
HClO hypochlorous acid.
H3PO3-phosphorous acid
H3BO3 -boric acid / NaOHsodium hydroxide
Ba(OH)2 barium hydroxide
KOHpotassium hydroxide
NH4OH ammonium hydroxide
Ca (OH)2 calcium hydroxide

Oxidation-Reduction (Redox) Reactions

Redox reactions are chemical reactions where electrons are transferred from one compound to another resulting in a chemical change.

E.g. Zn + H+Cl-(aq) --->Zn2+Cl2-1 (aq) + H2

Zn loses two electrons to H+ to form H2 gas. The oxidation state (charge on the metal or non-metal) of atoms on the left side of the chemical equation is changed when they react to form products during redox reactions. Single replacement reactions are redox reactions.

Which of the following reactions are, precipitation, acid-base, gas forming, and redox?

a) NaCl + AgNO3 ----> AgCl + NaNO3

b) NaOH + HCl ----> NaCl + H2O

c) Zn + 2HCl ----> ZnCl2 + H2

d) CaSO3(s) + 2HCl(aq) ŕ CaCl2(aq) + H2O(l) + SO2(g)

Properties of Solutions

Solutions occur commonly in nature and many chemical reactions occur in solutions. By learning about the nature of solutes, solvents and solutions, students can better understand both these natural phenomena and the chemistry of solutions.

A solution is defined as a homogeneous mixture.

Many materials exist as homogeneous mixtures

: air (g), lake water (l), gasoline (l), stainless steel (s).

Solution Components: Solvent and solute

Solution - homogeneous mixture of of two or more substances; often comprised of solvent and solute, e.g.,when gases or solids are dissolved in a liquid

Solvent- component with same phase as solution; substance present in excess in liquid-liquid mixtures

Solute- minor component mixed with solvent

Properties of Solutions

The majorities of chemical reactions, and virtually all important organic and biochemical processes, occur in solution. A solution is composed of one or more solutes dissolved in a solvent. In aqueous solutions, the solvent is water.

Transparent to Light

Liquid solutions are clear and transparent with no visible particles of solute. They may be colored or colorless, depending on the properties of the solute and solvent.

Electrical Conductivity of solutions

In solutions of electrolytes, the solutes are ionic compounds that have dissociated; the solution conducts electricity. Solutions of nonelectrolytes are nonconducting.

Solubility

The degree of solubility depends on the difference between the polarity of solute and solvent ("like dissolves like"), the temperature, and, for solutions of gases, the pressure.

A saturated solution contains all the solute that can be dissolved at a particular temperature. A supersaturated solution is an unstable condition that occurs when more than the theoretical amount of solute is temporarily held in solution.

Solubility and Equilibrium

Insoluble salts which barely dissolve and always have solid precipitate present. A typical example is

AgCl(s) + H2O Ag+(aq) + Cl-(aq)

This describesAgCl(s) dissolved in solution tosilver (Ag+) and chloride (Cl-) ions the aqueous phase (aq), in equilibrium with the sparingly soluble solid (s) salt AgCl. This is anexample of solubility equilibrium.

Solubility of Gases - Henry's Law

Henry's law describes the solubility of gases in liquids. At a given temperature the solubility of a gas is proportional to the partial pressure of the gas.

Sg = kHPg

where Sg is the solubility, kH is the Henry’s the Law constant, Pg is partial pressure of gas

Solute particle sizes

Colloidal suspensions have particle sizes between those of true solutions and precipitates. A suspension is a heterogeneous mixture that contains particles much larger than a colloidal suspension. Over time, these particles may settle, forming a second phase.

Colloids and suspensions

Mixtures include solutions, colloids and suspensions. Only solutions are considered true solutions.

Describe the following terms:

a) True solutions b) Colloids (Tyndall effect) c) Suspensions.

a) True solutions

Normally light passes through true solutions or they transparent to visible light. Because of small solute particle sizes they do not scatter light.. True solutions have solute particles with diameters less than 1 x103 pm. Colloids, on the other hand have solute particles with diameters in the range 1 x103 to 1 x105 pm. Suspensions have larger particle diameters which are greater than 1 x105 pm.

b) Colloids (Tyndall effect)

Colloids have solute particle diameters in the rage 1 x103 to 1 x105 pm. Colloids scatter light and the solution become opaque and relative degree of opacity depends on the sizes and amount of the particles. The scattering of light by colloidal particles is called Tyndell effect which has been used to distinguish between true solutions and colloids. This effect will increase with increasing solute particle diameter of a colloid. For example, milk is opaue because of the higher diameters of the solute particles such as proteins in the milk.

c) Suspensions.
Suspensions have particle diameters greater than 1 x105 pm. The particles in a suspension are lage enough to be affected by gravity and settle to the bottom of the mixture with time. For example, muddy water.

Surfactants

Surfactants constitute the most important group of detergent components. Generally, these are water-soluble surface-active agents comprised of a hydrophobic portion, usually a long alkyl chain, attached to hydrophilic or water solubility enhancing functional groups. The hydrophilic end, which is either polar or ionic, dissolves readily in water. The hydrophobic, or non-polar, end, however, does not dissolve in water. In fact, the hydrophobic, or "water-hating" end will move as far away from water as possible. This phenomenon is called the hydrophobic effect. Because one end of a surfactant resists water and the other end embraces it, surfactants have very unique characteristics.

Emulsifiers

A surfactant is also called an emulsifier. Emulsifiers do help oil and water remain in stable emulsions. Aggregates of oil and emulsifiers form micelles and stay dispersed in water solution. Examples of emulsifiers include sodium dodecyl sulfate.

Concentration of Solutions: Percentage

The amount of solute dissolved in a given amount of solution is the solution concentration. The more widely used percentage-based concentration units are:

weight/volume percent

weight/weight percent.

Weight/Volume Percent

Calculating weight/volume percent we use following equation:.

What is the % (W/V) of a 476 L Ar in CCl4 solution prepared by dissolving 50.0 g.

How many grams of NaOH is required prepare 2.0 L of 1.0 %(W/V) solution?

Solving for mass,

Substituting,

ow HWhat is the volume of the solution if 10.0g of NaCl is dissolved to obtain 25 %(W/V) solution?

Solving for volume,

Substituting,

Weight/weight percent.

What is the %(W/W) concentration of Ar in a gas solution prepared by mixing 50.0 g Ar with 80.0 g of He.

Volume/volume percent.

What is the %(V/V) concentration of a aqueous acetic acid solution prepared by adding 20.0 mL of acetic acid to water to a final volume of 2.50 L?

Concentration of Solutions: Moles and Equivalents

Molarity (M)is the number of moles of solute per liter of solution. Dilution calculations require both concentrations (V1 and V2) to be in the same units.

When discussing solutions of ionic compounds, molarity emphasizes the number of individual ions. In contrast, equivalents per liter emphasizes charge. One equivalent of an ion is the number of grams of the ion corresponding to a mole of electrical charges. Changing from molarity to equivalents per liter (or the reverse) is done using conversion factors.

Molarity

What is the molarity of a solution prepared by dissolving 0.700 g of KCl in water to a solution with 1.00 mL solution?

How many grams of AgNO3 are in 2.00 L of 0.500 M AgNO3solution?

Solving for moles of AgNO3,

mol AgNO3 = (MAgNO3)(L solution)

mol AgNO3 = (0.500 M)(2.00 L) = 1.00 mol AgNO3

Then, convert mol AgNO3 g AgNO3

Dilution

Concentrated solutions can be mixed with solvent to make weaker or dilute solutions. This is the kind of thing people do everyday with consumer products like fruit juice. Some concentrated solutions are used as "stock" solutions. Weaker solutions are typically used but the concentrated solutions require less storage space. In recent years accidents have occurred in the health care professions when dilutions were done incorrectly. Some of these errors have resulted in deaths or serious injuries. A number of health care facilities have abandoned the practice of "diluting" stock solutions because dilution instructions were too hard to follow. They do not want to take the risks associated with errors in preparing diluted solutions.

Mi = intial molarity; Vi = initial volume; Mf = final molarity; Vf = final volume

What would be the volume of a 0.002 M solution prepared staring with 50.0 mL of 0.400 M solution?

M i = 0.400 M

V i = 50.0 mL x (1L/103 mL) = 5.00 x 10–2 L

M f = 0.200 M

V f = ? L

Solving for the final volume of 0.200 M sugar, V f,

Representation of Concentration of Ions in Solution

To calculate the concentration of ions of a solution containing salts that can be dissociated into ions, for example: NaCl and

NaCl  1 Na+ + 1 Cl-,

CaCl2 1 Ca2+ + 2 Cl-

If the concentration of CaCl2 was 3 M concentration of Cl- will be 3 x 2 = 6 M.

Two ions (particles) are generated through the complete dissociation of NaCl. The number of particles formed on the dissociation is or NaCl is equal called Vant Hoff factor.Vant Hoff factor is 2 for NaCl. The Vant Hoff factor or CaCl2 is equal to 3.

For weak electrolytes which dissociate in completely the Vant Hoff factor is always less than the one expected from complete dissociation.

Concentration-Dependent Solution Properties

Colligative properties are properties that depend on concentration rather than the identity of the solute. Four colligative properties are: vapor pressure lowering, freezing-point depression, boiling-point elevation and osmotic pressure. Each has a number of very practical ramifications.

Molality (m) a new concentration unit, is the number of moles of solute per kilogram of solvent in a solution, and is more commonly used in calculations involvingboiling pointelevation and freezing point depression.

Mole fractiona new concentration unit, is the number of moles of solute per total moles of solute ans solvent in a solution, and is more commonly used in calculations involving, vapor pressure lowering.

A = solute; B = solvent

Mole fraction of A [X(A)] = (1 – mole fraction of B [X(B])

Four colligative solution properties are:

•Vapor pressure lowering. Raoult's Law states that when a solute is added to a solvent, the vapor pressure of the solvent decreases in proportion to the concentration of the solute.

•Freezing-point depression. When a nonvolatile solute is added to a solvent, the freezing-point of the resulting solution decreases.

•Boiling-point elevation. When a nonvolatile solute is added to a solvent, the boiling point of the resulting solution increases.

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

Vapor Pressure Lowering

Vapor-pressure of a liquid is lowered by addition of a solute. According to Raoult’s law it vapor of pressure of solution is inversely proportional to mole fraction of the solvent : Raoult's law reduces to

P = P0X

P = vapor pressure of the solution; P0 = vapor pressure of the pure solvent; X mole fraction of the solvent.

The freezing depression:

The freezing (melting or fusion) point of a liquid is decreased by dissolving a solid solute in water. The boiling point decrease or the depression of a solution is given by the equation:

Tf = Kf msolute

Tf = boiling point elevation.

Kf = molal freezing point depression constant-molal means concentration is in given as molality(m).

msolute= concentration of solute expressed as molality(m).

Boiling point Elevation

The boiling point of a liquid is increased by dissolving a solid solute in water. The boiling point increase or the elevation of a solution is given by the equation:

Tb = Kb msolute

Tb = boiling point elevation.

Kb = molal boiling point elevation constant-molal means concentration is in given as molality(m).

msolute= concentration of solute expressed as molality(m).

Osmotic pressure

Osmolarity (osmol), iM, isrelated to the molarity of the solution multiplied by a factor (i) to obtained number of particles of particles in the solution. Osmotic pressure () is calculated from the equation:

= i MRT.

i = vant Hoff factor for dissociating acid, bases, salts, compounds.

M = osmolarity of the solution.

T = Kelvin temperature of the solution.

Hypertonic solutions cause the outflow of water from a cell, resulting in cell collapse. Hypotonic solutions cause water to flow into the cell, resulting in hemolysis (cell rupture). Isotonic solutions have identical osmotic pressures.

Calculate the osmolarity of a 5.0 x 10-3 M glucose solution.

C6H12O6 is a non-electrolyte; it does not dissociate.

i = 1 and , since C6H12O6 is a non-electrolyte; it does not dissociate.

and,

Osmolarity is =

What is the osmotic pressure in atm of a 5.0 x 10-3 M glucose solution?

We found that the solution was 5.0 x 10-3 osmol; substituting this value for M (M represents osmolarity) in the expression:

 = MRT

 = 0.12 atm

Using our definition of osmotic pressure,  = MRT

M must be represented as osmolarity.

Now, substituting in our osmotic pressure equation:

 = (0.50 mol particles/L) x (0.0821 Latm/Kmol particles) x 298 K = 12 atm

(Note, two significant figures for the answer; the sucrose concentration has two significant figures)

Water as a Solvent

The role of water in the solution process deserves special attention. Water is often referred to as the "universal solvent," and it is the principal biological solvent. These characteristics are a direct consequence of the molecular geometry and structure of water and its ability to undergo hydrogen bonding.

Ammonia is a polar substance, as is water. The rule “like dissolves like” predicts that ammonia would be water soluble. Methane, a nonpolar substance, would not be soluble in water.