Solutions: Acid/Base

This wiki covers a very vast amount of material. The many different types of solutions are explained throughout the couple of chapters that are outlined. We start off learning about the many unique properties of water. This leads us into its properties regarding solutions and aqueous systems. There is extensive explanations of the many properties of solutions. Some examples of the properties that are discussed are solubility, concentrations, and colligative properties. It gets into acids and bases and their properties. There are examples of strong and weak acids and bases. There are many subjects discussed within this one wiki but all relate to one another and are as important as one another.

Editor: Olivia Newton John

Group 1

Even Sommerich (pg 445-449)

Water and its Properties

• It can be a solid, liquid or gas.
• Its molecules are polar.

• It has many unique qualities. Such as its surface tension, its ability to change state and its low vapor pressure
• Surface tension-inward force that tends to minimize the surface area of a liquid.
• Surfactant- any substance that interferes with the hydrogen bonding process between water molecules, causing an effect on its water tension
• Vapor pressure-a measure of the force exerted by a gas above a liquid in a sealed container.

Water in a solid state

• Water freezes at 0 degrees Celsius and 32 degrees Fahrenheit.
• Its structure forms that which looks like a honeycomb.

Group 2

Sweet Caroline (pg 450-458)

Solvents and Solutes

• Aqueous solution: Water that contains dissolved substances.
• In a solution: The dissolving medium is the solvent, and the dissolved particles are the solute.
• A solvent dissolves the solute. The solute becomes dispersecd in the solvent.
• Solutions are homogenous mixtures. They are stable mixtures.
• Substances that dissolve most readily in water include ionic compounds and polar covalent molecules.

The Solution Process

• Water molecules are in continuous motion because of their kinetic energy,
• When a crystal of sodium chloride is placed in water, the water molecules collide with it.
• The polar solvent molecules (H2O) attract the solute ions (Na+ Cl-)
• As individual solute ions break away from the crystal, the negatively and positively charged ions become surrounded by solvent molecules and the ionic crystal dissolves.
• The process when the positive and negative ions of an ionic solid become surrounded by solvent molecules is called solvation.

_

• Water and oil do not mix. But what about iol in gasoline?
• Both oil and gaoline are made up of nonpolar molecules.

- This meaning, oil molecules can easily separate and replace gasoline molecules to form a solution.

This picture displays how water and oil DO NOT mix.

Electrolytes and Nonelectrolytes

• Electrolyte: is a compound that conducts an electric current when it is in a aqueous solution or in the molten state. 
• To conduct electricity it is necessary that the ions are mobil and able to carry an electrical current.
• All ionic compounds are electrolytes because they dissociate into ions.
• Nonelectrolyte: a compound that does not conduct an electric current in either aqueous solution or the molten states.
• Most compounds are nonelectrolytes becasue they are not composed of ions.
• Storong electrolyte: a solution in which a large protion of the solute exists as ions.
• Example of strong electrolyte: sodium chloride
• Weak electrolyte: a solution that conducts electricity poorly because only a fraction of the solute exists as ions.
• Example of weak electrolyte: ammonia, organic acids and bases
• 
  

Hydrates

• The water contained in a crystal is called the water of hydration or water of crystallization.
• Hydrate: a compound that contains water of hydration
• In writing the formula of a hydrate, use a dot to connect the formula of the compound and the number of water molecules per formula unit.
• 
  

Group 3

Big Dan's Restaurant (pg 459-463)

Heterogeneous Aqueous Systems:

• Suspentions:

  • homogeneous mixtures are soloutions but heterogenneous mixtures are not

  • a suspention is a mixture from which particles settle out upon standing

  • a suspention differs from a soloution because the suspentions doesnt stay indefinitely

  • suspentions are heterogeneous because at least two substances can be identified

• Colloids

  • gelatian is a type of mixture called a colloid

  • a colloid is a heterogeneous mixture containing particles that range in size from 1nm to 100nm.

  • some examples of colloids are glue, gelatin, paint, aresol sprays and smoke

  • colloids have particles smaller that suspentions and larger than a soloution

• The Tyndall Effect

  • the tyndall effect is when a light is shone through a liquid being a slolution, colloid or a suspention and when the light passes through the solution there is no beam and when the light passes through the colloid and the suspention the light beam is visible

• The Brownian Motion

  • it is when the scintillations of light in a colloid are shown; when light hits the particles in the colloid the light gets reflected and the light becomse visibleC

• Coagulation

  • a colloid has particle of all the same charge so that it stays dispersed, it can be destroyed by adding particle of the opposite charge into the mixture

• Emulsions

  • an Emulsion is the a colloidal dispersion of a liquid in a liquid.

Group 4

Golf Tee (Alex) (pg 471-479)

Properties of Solutions:

Solution formation-

• Solutions = Homogeneous mixtures that can be solid, liquid or gas

• The compositions of the solvent and the solute determine whether a substance will dissolve. Stirring/agitation, temperature, and the surface area of the dissolving particles determine how fast the substance will dissolve.

• Stirring and solution formation-
  • Stirring speeds up the rate at which the solute dissolves.
  • Does not effect the amount of solute that will dissolve.

• Temperature and solution formation-
  • Higher temperature = faster dissolving rate
    • Greater frequency and force of collisions between solute and solvent molecules

• Particle size and solution formation-
  • Bigger surface area = faster dissolving rate
    • More area exposed to colliding molecules

http://www.youtube.com/watch?v=UaFRYYQe6-c

Solubility-

• Some particles will never dissolve no matter because of an “exchange process”
  • New particles from whatever is being dissolved are solvated and enter the solution
  • An equal number of already dissolved particles crystalize, come out of the solution, and deposit as a solid
    • Mass of undissolved crystals remains constant

• Saturated solution-solution containing the maximum amount of solute for a given amount of solvent at a constant temperature and pressure; an equilibrium exists between undissolved solute and ions in solution:
• Solubility-the amount of a substance that dissolves in a given quantity of solvent at specified conditions of temperature and pressure to produce a saturated solution:
  • Often expressed in grams of solute per 100 g of solvent
• Unsaturated solution-a solution that contains less solute than a saturated solution at a given temperature and pressure:
• Miscible-describes liquids that dissolve in one another in all proportions:
  • Liquid that has larger amount = solvent
  • Liquids that are slightly soluble = partially miscible
• Immiscible- describes liquids that are insoluble in one another; ex-oil & water:

Factors affecting solubility-

Temperature:

• Temperature affects the solubility of solid, liquid, and gaseous solutes in a solvent- both temperature and pressure affect the solubility of gaseous solutes.

• Solubility of most solids increases as temperature increases
  • But, for some exceptions solubility decreases with temperature (ex. ytterbium)

• Supersaturated solution-solution that contains more solute than it can theoretically hold at a given temperature; excess solute precipitates if a seed crystal is added:
  • Seed crystal = very small crystal of solute
  • Example- rock candy

• Effect of temperature on the solubility of gases in liquid solvents is opposite that of solids- solubilities of most gases are greater in cold water then in hot

Temperature & Water Solubility in Grams of Solute per 100 ml of H2O
Temperature(Celsius)	Solid Citric Acid	Solid Potassium Phosphate	Gaseous Nitrogen	Gaseous Oxygen
0	49	44	0.0030	0.0070
20	59	50	0.0020	0.0050
30	64	x	x	x
50	71	62	x	0.0031
70	76	x	x	x
100	84	x	x	0.0029

(x = not available)

Pressure:

• Changes in pressure have little affect on solubility of solids and liquids
• Strongly influences solubility of gases
  • Gas solubility increases as the partial pressure of the gas above the solution increases
  • Example- carbonated drinks

• Henry’s law-at a given temperature the solubility of a gas in a liquid is directly proportional to the pressure of the gas above the liquid:
  • S/P = S/P
  • This is how you find the solubility of a gas

Group 5

Stevie Wonder (pg 480-485)

Concentrations of Solutions
-The concentration of a solution is a measure of the amount of solute that is dissolved in a given quantity of solvent
-Dilute solution is one that contains a small amount of solute
-Concentrated solution is one that contains large amount of solute
-Molarity is number of moles dissolved in one liter of solution
Molarity=moles of solute/liters of solution
-Diluting a solution reduces the number of moles of solute per unit volume, but the total number of moles of solute in solution does not change
Moles of Solute=M1 x V1=M2 x V2
The concentration of a solution in percant can be expressed in two ways:
-ratio of volume of solute to volume of solution
-ratio of mass of solute to mass of solution
Percent by volume (%(v/v))=volume of solute/volume of solution x 100%
http://www.youtube.com/watch?v=RCbhk3yyM88

Group 6

North/South (East) Dakota (pg 487-490)

Colligative Properties of Solutions
– Dakota Pimentel

• Colligative properties: a property that is dependant on the number of solute particles rather than the identity of the solute
  • Freezing point --> Depression
  • Boiling point --> Raising
  • Vapor pressure --> Dropping

• Vapor pressure: (dropping) “pressure exerted by a vapor that is in dynamic equilibrium with its liquid in a closed system.”
  • Non-volatile solute: a solute that is not easily vaporized creates a solution with a lower vapor pressure than its formerly pure solvent.
    • This is because the water, the solvent in the solution, surrounds the solute. This restrains some water molecules from vaporizing, causing a lower vapor pressure.

Key Concept: The decrease in a solution's vapor pressure is proportional to the number of particles the solute makes in a solution

• Freezing Point:(Depression) The temperature at which a solution turns from a liquid to a solid.
  • Freezing-point drop: The difference, in temperature, between the freezing point of a solution and the freezing point of its pure solvent.
    • When a substance freezes, the particles arrange themselves in an orderly pattern.
    • When a solute freezes, it is more difficult to make this pattern, making the freezing point drop lower than the pure solvant.
• 1 mol of any solute added to a solution drops the freezing point of 1000 g of water 1.86°C.

Key Concept: The magnitude of the freezing-point depression is proportional to the number of solute particles dissolved in the solvent and does not depend upon their identity.

• Boiling-Point:(Elevation) the difference in temperature between the boiling point of a solution and the boiling point of the pure solvent.
  • A decrease in vapor pressure causes additional energy to be needed in order for the solution to boil.
• 1 mol of solute added to a solution increases the boiling point of 1000 g of water by 0.512°C.

Key Concept: The boiling-point elevation is proportional to the number of solute particles dissolved in the solvent.




Pictures: Found by Dakota Pimentel

Group 7

Hole Punch (Mitch) (pg 491-493)

Molality

Molality: (m) The number of moles of solute dissolve in 1 kilogram of solvent.
Molality (m)= moles of solute
kilogram of solvent

Note: Molality is not the same as molarity.

This guy is pretty boring but it sounds like he knows what he is talking about
http://www.youtube.com/watch?v=AxgF6oMXJwg

Brandooski (494-496)

Group 8

The Comet (pg 587-589)

Properties of Acids and Bases

Acids:

• Acidic compounds give foods a tart or sour taste
• Aqueous solutions of acids are electrolytes
  • electrolytes conduct electricity
  • some are stronger electrolytes then others
• Acids Contain chemical dyes, called indicators, which change colors
• Acids react with compounds containing hydroxide ions to form water and a salt

Bases:

• Have a bitter taste, but most bases are hazardous to eat
• The slippery feeling of soap is another property of bases
• Aqueous solutions of bases are electrolytes, and will cause an indicator to change color
  • Can be strong or weak electrolytes

Arrhenius Acids and Bases

• In 1887, the Swedish chemist Svante Arrhenius proposed a revolutionary way of defining and thinking about acids and bases.
• He said that acids are hydogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solutions
• He also said that bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solutions

Arrhenius Acids:

• Acids that contain one ionizable hydrogen,, such as nitric acid (HNO3), are called monoprotic acids
• Acids that contain two ionizable hydogens, such as sulfuric acid (H2SO4) are called diprotic acids
• Acids that contain three ionizable hydrogens, such as phosphoric acid (H3PO4), are called triprotic acids
• Only the hydrogen in very polar bonds are ionizable
  • In such bonds, hydrogen is joined to a very electronegative element
• Only the hydrogen bonded to the highly electronegative element can be ionized

Arrhenius Bases:

• The base you are probably most familar with is Sodium Hydroxide (NaOH).
  • Sodium Hydroxide is an ionic solid
  • It ionizes to form a sodium ion and a hydroxide ion in aqueous solution
• The elements in Group 1A, the alkali metals, react with water to produce so

Nate the Rake (pg 590-593)

Brønsted-Lowry Concept

In this section we will consider the Brønsted-Lowry concept. This concept focuses on what an acid or base does .

Acids

With the Brønsted-Lowry concept we usually refer to a hydrogen ion as a proton . That is because a proton is all that is left when a hydrogen atom loses an electron to become anion .
Brønsted and Lowry independently came up with the idea that an acid is an acid because it provides or donates a proton to something else . When an acid reacts, the proton istransferred from one chemical to another. As will be noted later, the chemical which accepts the proton is a base.

When an acid dissolves and dissociates in water it gives a proton to the water. Equations to represent this are shown here ( and in example 16 in your workbook). The Brønsted-Lowry view is that the acid (HCl) gives a proton to water to make two ions, one of which is H3O+. H3O+ is called hydronium ion . (By the way a hydronium ion is sometimes called an oxonium ion.)		HCl + H2O	rtarrow.gif (850 bytes)
rtarrow.gif (850 bytes)

H3O+ + Cl-

These equations show a different acid (H2SO4) giving a proton to water. In this case, the product HSO4- still has a proton that can be donated to another water molecule.		H2SO4 + H2O	rtarrow.gif (850 bytes)
rtarrow.gif (850 bytes)

H3O+ + HSO4- ||

HSO4- + H2O	rtarrow.gif (850 bytes)
rtarrow.gif (850 bytes)

H3O+ + SO42- || ||

This equation shows HCl giving a proton to a hydroxide ion (OH-) rather than water.		HCl + OH-	rtarrow.gif (850 bytes)
rtarrow.gif (850 bytes)

H2O + Cl-

The first chemical in each of these equations is an acid because they are each giving a proton to something else.

Bases

Note that in order for an acid to act like an acid, there needs to be something for it to react with. There needs to be something to take the proton. There needs to be a base . A base is a proton acceptor . Compare this to the definition that an acid is a proton donor.

Bases are the opposite of acids. Bases are basic because they take or accept protons . Hydroxide ion, for example can accept a proton to form water. Brønsted and Lowry realized that not all bases had to have a hydroxide ion. As long as something can accept a proton it is a base.		OH- + H+	rtarrow.gif (850 bytes)
rtarrow.gif (850 bytes)

H2O

So anything, hydroxide or not, that can accept a proton is a base under the Brønsted-Lowry definition. The water molecules that accept protons when HCl dissolves in water are acting as bases.

external image Ammonia-3D-balls-A.png

Acid	Formula
Acetic	CH3COOH
Hydrochloric	HCl
Sulfuric	H2SO4
Nitric	HNO3
Carbonic	H2CO3

Base	Formula
Ammonia water	NH4OH
Sodium hydroxide	NaOH
sodium carbonate	Na2CO3
soaps	Varies

Lewis acids and bases

The most general definition of acids and bases, which encompasses the Arrhenius and Bronsted-Lowry definitions is due to our old friend, Lewis and his dot structures. A Lewis acid is defined to be any species that accepts lone pair electrons. A Lewis base is any species that donates lone pair electrons. Thus,

${rm H}^+$

${rm H}^+$

is a Lewis acid, since it can accept a lone pair, while

${rm OH}^-$

${rm OH}^-$

and NH

$_3$

$_3$

are Lewis bases, both of which donate a lone pair:

begin{displaymath}{rm H}^++ stackrel{..}{stackrel{:{rm O}:}{..}}{rm H}^- longrightarrow {rm H}_2{rm O}end{displaymath}

begin{displaymath}{rm H}^++ stackrel{..}{stackrel{:{rm O}:}{..}}{rm H}^- longrightarrow {rm H}_2{rm O}end{displaymath}

Interestingly, however, is that species which have no hydrogen to donate (a la the Bronsted-Lowry scheme) can still be acids according to the lewis scheme. As an example, consider the molecule BF

$_3$

$_3$

. If we determine Lewis structure of BF

$_3$

$_3$

, we find that B is octet deficient and can accept a lone pair. Thus it can act as a Lewis acid. Thus, when reacting with ammonia, the reaction would look like:

begin{figure}begin{center}leavevmodeepsfbox{lec21_fig1.ps}{small}end{center}end{figure}

begin{figure}begin{center}leavevmodeepsfbox{lec21_fig1.ps}{small}end{center}end{figure}

Group 9

Christo the Count (pg 594-596)

Hydrogen Ions from Water
-water molecules are highly polar and in continuous motion
when a water molecule loses a hydrogen ion it becomes a negatively charged ion- hydroxide (OH-)
-adds one- becomes positively charged hydronium ion (H3O+)
the reaction in which water molecules produce ions is called the self-ionization of water
Any aqueous solution where hyudroxide and hydronium are equal is a neutral solution
For an aqueous solution, the product of the hydrogen-ion concentration and hydroxide ion concentration equals 1.0 x 10^ -14
Product of the concentrations of the hydrogenn ions and hydroxide ions in water is called the ion product constant for water
acidic= H+ > OH-

basic= OH-> H+

basic solution is also known as alkaline solutions

Angry-a (pg 596-604)

The pH Concept

• The pH scale was established by Soren Sorensen in 1909.

• The pH scale ranges from 0 to 14.
  • A pH of zero is acidic.
  • A pH of 7 is neutral.
  • A pH of 14 is basic.
  • 
• The pH of a solution is the negative logarithm of the hydrogen-ion concentration.

• You must use the log function key on a calculator to calculate the pH of a solution.
• A solution in which H+ is greater than 1 x 10^-7M has a pH less than 7.0 and is acidic.
• The pH of pure water or a neutral aqueous solution is 7.0.
• A solution with a pH greater than 7 is basic and has a H+ of less than 1 x 10^-7M.
  • In an acidic solution, there are more H+ ions.
  • In a basic solution, there are more OH- ions.

• The pOH of a solution equals the negative logarithm of the hydroxide-ion concentration.
  • A pOH of 7 is neutral.
  • A pOH of 0 is basic.
  • ApOH of 14 is acidic.
  • pOH + pH = 14
  • Most pH values are not whole numbers.
• You can calculate the hydrogen-ion concentration of a solution if you know the pH.
• If the pH is not a whole number, use the y^x function on a calculator to calculate the hydrogen-ion concentration.

• Measuring pH
  • The pH scale is used to create a balance of acid-bases in a swimming pool, to create soil conditions ideal for plant growth, and to medical diagnoses.
  • If you know the OH- of a solution, you can find its pH.
  • The ion-product for water defines the relationship between H+ and OH-.
• Acid-Base Indicators
  • An indicator is an acid or base that undergoes disociation in a known pH range. An indicator is a valuable tool for measuring pH because its acid form and base form have different colors in a solution.
  • The acid form dominates at a low pH.
  • The base form has a high pH.

• On an indicator, the color of the solution is a mixture of the acid and base forms. This estimated color of the pH solution tells if the solution is an acid or base.
• Indicators are usually tested at 25 degrees Celcius. If they are tested at other temperatures, the color indicator may be distorted.
• Dissolved salts may also effect the indicator;s solution.
• An indicator strip is a piece of paper or plastic impregnated with an indicator.

• The pH Meters
  • A pH meter connected to a computer can be used to make a continuous recording of pH changes.
  • A pH meter is easier and more accurate than indicator strips.
  • Color nor cloudiness of the solution effect the pH.

Group 10

SHANOOSK (pg 605-607)

19.3 – Strengths of Acids and Bases (Shannon Lamy)

Strong and Weak Acids and Bases

*Whether an acid is strong or weak depends on its ability to ionize in water.

• Strong Acids – completely ionize in water

Ex. Hydrochloric acid

• Weak Acids – slightly ionize in water

Notice how the particles break apart in a strong acid, but not in a weak acid.

• http://www.youtube.com/watch?v=dv9UK2m9c7w

^^Here is a strange little video to help understand the concept of dissociation of strong and weak acids.

Acid Dissociation Constant

*Dissociation is when an acid breaks up into smaller particles

• Strong Acids – completely dissociate in water; H3O+ is high

• Weak Acids – undissociate in water; H3O+ is low

*Equilibrium-constant equation can be formed from balanced chemical equations

*The concentration of water in a dilute solution is constant

• Acid dissociation constant (Ka) - ratio of the concentration of the dissociated form of an acid to the concentration of the undissociated form.

• Dissociation constants are sometimes called ionization constants

Key Point – Weak acids have small Acid dissociation constants, stronger acids have stronger Acid dissociation constants.
Stronger Acids have more complete ionization

Peter Rabbit IS THE MAN (pg 608-611)

There happens to be both strong and weak bases as well as acids.

• strong bases - dissociate completely into metal ions and hydroxide ions in aqueous solutions
  • Examples of strong bases include:
    • calcium & magnesium - hydroxide
  • Strong bases do not react well with water
• Weak Bases - React with water to form hydroxide ion and the conjugate acid base
  • Example of a weak base is:
    • Ammonium

PJ

• Base Dissociation Constant - (K subscriptb) - is the ratio of the concentration of the conjugate acid times the concentration of the hydroxide ion to the concentration of the base

PJ

PJ

Peter Rabbit IS THE MAN (pg 608-611)

Group 11

The Zoe (pg 612-613)

Acid-Base Reactions:

Zoey Killion

If you mix a solution of strong acid containing hydronium (hydrogen) ions with a solution of a strong base that has an equal number of hydroxide ions, a neutral solution results. The solution is neither acidic or basic, but neutral.

neutralization reactions - reactions in which an acid and a base react in an aquueous solution to produce a salt and water
In general, the reaction of an acid with a base produces water and one of a class of compounds called salts.
Some Salts and Their Applications:

Name	Formula	Applications
Ammonium Sulfate	(NH4)2SO4	Fertilizer
Barium Sulfate	BaSO4	Gastrointestinal studies; white pigment
Calcium Chloride	CaCl2	De-icing Roadways and sidewalks
Calcium Sulfate Dihydrate (Gypsum)	CaSO4 x 2H20	Plasterboard
Copper Sulfate Pentahydrate (Blue Vitriol)	CuSOS4 x 5H20	Dyeing, Fungicide
Calcium Sulfate Hemihydrate	CaSO4 x 1/2H20	Plaster Casts
Potassium Chloride	KCl	Sodium-free Salt Substitute
Potassium Permanganate	KMnO4	Disinfectant and Fungicide
Silver Nitrate	AgNO3	Cauterizing Agent
Silver Bromide	AgBr	Photographic Emulsions
Sodium Hydrogen Carbonate (Baking Soda)	NaHCO3	Antacid
Sodium Carbonate Decahydrate (Washing Soda)	Na2CO3 x 10H20	Glass Manufacture; Water Softner
Sodium Chloride (Table Salt)	NaCl	Body Electrolyte; Chlorine Manufacture
Sodium Thiosulfate (Hypo)	Na2S2O3	Fixing Agent in Photographic Process

Titration (Page 613):
Zoey Killion
Acids and bases sometimes, but not always, react in a 1:1 mole ratio. When sulfuric acid reacts with sodium hydroxide, however, the ratio is 1:2. Two moles of the bse sodium hydroxide are required to neutralize one mole of H2SO4. Similarly, hydrochloric acid and calcium hydroxide react in a 2:1 ratio.
equivalence point - of a titration is when the numnber of moles of hydrogen ions equals the number of moles of hydroxide ions.

Big Mikey-Mike (IN YOUR FACE) (pg 614-616)

Titration (Continued)

• You can determine the concentration of an acid or base in a solution by performing a neutralization reaction.
• An appropriate acid-base indicator is usually needed in order to show when neutralization has occurred.
• Some examples of acid-base indicators include the juice of a red cabbage and phenolphthalein

A sample of several test tube in red cabbage juice indicators


An example of an acid changed through
titration using phenolphthalein

• In neutralization reactions, the following are taken (This is the same as the Titration lab we did in class)
  1. A measured volume of an acid solution of unknown concentration is added to a flask
  2. Several drops of the indicator are added to the solution while the flask is gently swirled.
  3. Measured volumes of a base of known concentration are mixed into the acid until the indicator just barely changes color.

This image shows someone performing
a titration by slowly releasing the indicator
into the flask until the acid changes color.

• Titration is the process of adding a known amount of solution of known concentration to determine the concentration of another solution.
• A Standard Solution is the solution which has a known concentration.
• The End Point is the point at which the indicator changes color at the conclusion of the titration.
• Key concept: The point of neutralization is the end point of the titration.
• This link gives a very good demonstration of a titration experiment:
• http://www.youtube.com/watch?v=g8jdCWC10vQ
  
  

Group 12

Rebecca From Sunny Brook (pg 618-620)

Salt Hydrolysis

• a salt consists of an anion from an acid and a cation from a base
• although solutions of many salts are neutral, some are acidic and others are basic
• CH3COOH(aq) + NaOH(aq) = CH3COONa(aq) + H2O(l)

Ethanoic Acid + Sodium Hydroxide = Sodium ethanoate + Water

• in salt hydrolysis, the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ions to water
• hydrolyzing salts are usually derived from a strong acid and a weak base, or from a weak acid and a strong base
• in general, salts that produce acidic solutions contain positive ions that release protons to water; salts that produce basic solutions contain negative ions that attract protons from water
• sodium ethanoate is the salt of a weak acid- ethanoic acid, and a strong base- sodium hydroxide:
  • CH3COONa(aq) = CH3COO-(aq) + NA+(aq)
• ethanoate ion forms electrically neutral ethanoic acid and negative hydrocixide ions (equilibrium):
  • CH3COO-(aq) + H2O(l) = CH3COOH(aq) + OH-(aq)
• ammonium chloride is the salt of a strong acid- hydrochloric acid, and a weak base- ammonia
  • NH4Cl(aq) = NH4+(aq) + Cl-(aq)
• ammonium ion is a strong enough acid to donate a hydrogen ion to a water molecule
  • NH4+(aq) + H2O(l) = NH3(aq) + H3O+(aq)
• how to determine if a solution is acidic or basic:
  • strong acid + strong base = neutral solution
  • strong acid + weak base = acidic solution
  • weak acid + strong base = basic solution

*for more help with salt hydrolysis, see this video demonstration: http://www.youtube.com/watch?v=cZftblLcXbk

Elizabeth Newton John (620-621)

Buffers

• Buffer: A solution in which the pH remains constant when small amounts of acid or base are added; a buffer can be either a solution of a weak acid and the salt of a weak acid or a solution of a weak base and the salt of a weak base.

• A buffer's pH does not greatly change when a small amount of strong acid or base is added to it
• Buffer solutions are better able to resist drastic changes in pH than pure water
• Buffers are used to keep pH at a nearly constant value in chemicals

Titration of an acified solution of a weak acid

• An example of a buffer solution in our daily lives is blood

Human blood must be buffered to remain within the necessary pH level

• Buffers are used to keep correct pH in enzymes for organisms to work
  • Buffers are important because when things such as enzymes are not in the proper pH range, they may have difficulty functioning
• Example of how a buffer works:
  • Ethanoic acid (CH3COOH) and its anion (CH3COO-) act as reservoirs of neutralizing power and react with hydroxide ions added to the solution
  • When the acid is added, the ethanoate ions act as a sponge and absorb the hydroxide ions, creating ethanoic acid, which does not ionize extensively in water, allowing the pH to stay pretty stable
  • When a base is added to the solution, the ethanoic acid and hydroxide ions react to produce ethanoate ion, which is not very strong and cannot accept many hydrogen ions, causing the pH not to change much.
• Buffer Capacity: A measure of the amount of acid or base that may be added to a buffer solution before a significant change in pH occurs.

Diagram of Buffer Capacity