Solutions Acids and Bases
Editor: Nikki Sheehan

This unit deals with three major topics, water/aqueous systems, solutions, and acid-base theories. Aqueous refers to anything dealing with water and aqueous systems can be homogeneous (meaning uniform throughout) or heterogeneous (meaning components are not evenly distributed/easily distinguished). A solution is a homogeneous mixture consisting of solutes dissolved in a solvent. An acid is is a compound that produces hydrogen ions in solutions and a base is a compound that produces hydroxide ions in solutions. These basic definitions will be used as the foundation of our lessons in these chapters and knowing them will help in our growing knowledge of solutions, acids, and bases.
Key Equations
Percent H2O = (mass of water) / (mass of hydrate) x 100%
Henry's Law: (S1) / (P1) = (S2) / (P2)
Molarity (M) = (moles of solute) / (liters of solution)
Titration: M1 x V1 = M2 x V2
Percent by volume = (volume of solute) / (volume of solution) x 100%
Percent by mass = (mass of solute) / (mass of solution) x 100%
Molality (m): (moles of solute) / (kilograms of solvent)
Mole fractions: Xa = (na) / (na + nb) Xb = (nb) / (na + nb)

Kw = [H+] x [OH-] = 1.0 x 10-14
pH = -log[H+]
pOH = -log[OH-]
Ka = ([H+][A-]) / ([HA])
external image chemical+mixing.jpgHydrochloric_acid_05.jpgexternal image 20110415114023_Caustic-Soda-Flake.jpg
Nikki Sheehan

Group 1 (445-449) Water and Its Properties

Kayla Anghinetti

I. Water in the Liquid State
A. Water - simple triatomic molecule - H2O (oxygen atom forms covalent bond with each hydrogen atom)
B. oxygen attracts electron pair of covalent O-H bond to a greater extent because of its greater electronegativity.
1. oxygen atom acquires partially neg charge
2. less electronegative hydrogen gets partial pos charge
3. as a result, O-H bonds are very polar
C. Polar molecules are attracted to one another by dipole interactions (negative end attracts positive end)
D. Many unique and important properties of water- including its high surface tension and low vapor pressure- result from hydrogen bonding.

SURFACE TENSION
- defined as the inward force that tends to minimize the surface area of a liquid
- all liquids have surface tension, water is higher than most
- you can decrease surface tension of water by adding a surfactant (any substance that interferes with the hydrogen bonding between water molecules)
- soaps and detergents are surfactants

VAPOR PRESSURE
- result of molecules escaping from the surface of the liquid and entering the vapor phase
- because hydrogen bonds hold water molecules to one another, the tendency of these molecules to escape is low and evaporation is slow.

II. Water in the Solid State
- increasing density means that the molecules of a liquid move closer together so that a given volume of a liquid contains more molecules and more mass
- if cooling continues, liquid solidifies with a density greater than that of its liquid phase
- as water cools it initially behaves like a typical liquid
- structure of ice is a regular open framework of water molecules arranged like a honeycomb

external image water.jpg

Group 2 (450-458)

Lindsey Chou

Homogeneous Aqueous Systems
Solvents and Solutes
  • Aqueous solution = water that contains dissolved substances
  • Solvent = the dissolving medium
  • Solute = the dissolved particles
  • A solvent dissolves the solute, the solute becomes dispersed in the solvent
  • Can be any state of matter
  • Solutions = homogeneous and stable mixtures
  • Ionic compounds and polar covalent molecules dissolve most readily in water
  • Non-polar covalent molecules don’t dissolve in water
solvent_and_solute.jpg

(Lindsey Chou)
The Solution Process
  • Solvation = the process by which positive and negative ions of an ionic solid become surrounded by solvent molecules
  • As solute ions break away from the crystal:
  1. The negatively/positively charged ions become surrounded by solvent molecules
  2. The ionic crystal dissolves
Solvation.gif
“Like dissolves like” – polar solvents dissolve ionic compounds and polar compounds, nonpolar solvents dissolve nonpolar compounds
Electrolytes and Nonelectrolytes
  • Electrolyte = compounds that conducts electric current when in an aqueous solution/molten state
  • All ionic compounds are electrolytes!
  • Nonelectrolyte = does not conduct electric current in either the aqueous solution/molten state
  • Some polar molecules become electrolytes when dissolved in water
  • Strong electrolyte = conducts electric current better because most of the solution exists as ions
  • Weak electrolyte = conducts electricity poorly because only some of the solute exists as ions
Hydrates
  • Hydrate = compound that contains water of hydration
  • To write the formula of a hydrate, connect the formula of the compound and the # of water molecules per formula unit with a dot
  • Efflorescent hydrates = when the vapor pressure of the hydrate is higher than the pressure of water vapor in the air, the hydrate loses its water (effloresces)
  • Hygroscopic hydrates = hydrates with low vapor pressure remove water from moist air to form higher hydrates
  • Deliquescent compounds = compounds that are so hygroscopic that they become wet when exposed to normally moist air, they remove water from air to dissolve completely and form solutions
  • Percent H2O = mass of water/mass of hydrate x 100%

Group 3 (459-463)

Elena Conroy



15.3 Heterogeneous Aqueous Solutions

By: Elena Conroy

Suspensions

  • Heterogeneous mixtures are not solutions
  • A suspension is a mixture from which particles settle out upon standing
  • A suspension differs from a solution because the particles of a suspension are much larger and do not suspended indefinitely
  • Particles in a suspension are usually greater than 1000nm in diameter
  • The particles in a solution are usually 1nm

Colloids

  • A colloid is a heterogeneous mixture containing particles that range in size from 1nm to 1000nm
  • Examples of colloids are glues, gelatin, paint, aerosol sprays, and smoke
  • Colloids have particles smaller than those in suspensions and larger than those in solutions
external image images?q=tbn:ANd9GcTizRjBVsdqu79XkGxPqFV-lEiZXDUw1105LCsiS_fO4PTwX3i5iQ&t=1

The Tyndall Effect

  • The Tyndall Effect is the scattering of visible light by colloid particles
  • Suspensions also exhibit the Tyndall Effect, but solutions do not
external image tyndall.jpg

Brownian Motion

  • Brownian Motion is the chaotic movement of colloidal particles
  • It was first observed by the Scottish botanist Robert Brown
  • It is caused by collisions of the molecules of the dispersion medium with the small, dispersed colloidal particles
  • These collisions help the colloidal particles from settling

Coagulation

  • Colloidal particles also tend to stay suspended because they become charged by absorbing ions from the dispersion medium onto the surface
  • All the particles in a particular colloid system will have the same charge
  • The repulsion between the like-charged particles prevents the particles from forming heavier aggregated that would have a greater tendency to settle out
  • Addition of ion that have an opposite charge can cause particles to neutralize and clump together, disturbing the colloid

Emulsions

  • An emulsion is a colloidal dispersion of a liquid in a liquid
  • An emulsifying agent is essential for the formation of an emulsion and for maintain an emulsions stability
  • An example of an emulsion is mayonnaise, which egg yolk acts as a emulsifying agent
external image mayonnaise.jpg
external image table22-3.jpg

Group 4 (471-479)

Bryan Dextradeur


16.1 Properties of Solutions

Solution Formation

  • Solutions are homogeneous mixtures that may be solid, liquid, or gaseous.
  • The compositions of the solvent and the solute determine whether a substance will dissolve. Stirring (agitation), temperature, and surface area of the dissolving particles determine how fast the substance will dissolve.
    • Example: granulated sugar dissolves faster than sugar cubes in a beverage and sugar dissolves faster in tea than in iced tea.

Stirring and Solution Formation

  • When a solution is stirred (agitated), it causes the solute to dissolve faster because fresh solvent is continually brought into contact with the surface of the solute.
  • Stirring only increases the rate at which the solute will dissolve, not how much of the solute will dissolve. An insoluble substance will still remain undissolved regardless of how long the solution is shaken.
    • An example of this is how mixing sugar into a drink such as tea increases the rate of the sugar dissolving.

Temperature and Solution Formation

  • At higher temperatures, the Kinetic Energy of water molecules is greater than at lower temperatures, so they move faster.
  • The more rapid motion of the solvent molecules leads to an increase in the frequency and the force of collisions between water molecules and the surface of the solute's molecules.
    • An example of this is how sugar dissolves faster in hot tea than it does in iced tea.

Particle Size and Solution Formation

  • The more surface of the solute that is exposed, the faster the solute will dissolve.
  • A much greater surface area of a solute area leaves more surface area exposed to collisions with water molecules.
    • A example of this is how granulated sugar dissolves faster than sugar cubes in tea because more surface is exposed to the tea.

Solubility

  • According to Kinetic Theory, water molecules are in constant motion.
  • This could lead to the assumption that water can dissolve any amount of a substance. However, when a solute in placed in water, opposite processes are occurring. While the solute is being dissolved, the dissolved particles crystallize to form a solid. At equilibrium, these two processes occur at an equal rate, and the substance remains completely dissolved. However, if there is more of the substance, no matter how long it is left in the water, there will still be some solid left over because crystallization is occurring at a faster rate than dissolving is.
    • An example of this is that if you place 36g of sodium chloride into 100g of water at 25oC, it will dissolve completely, but is you add just one gram more of sodium chloride, not all of it will dissolve.
  • A Saturated Solutioncontains the maximum amount of solute for a given quantity of solvent at a constant temperature and pressure.
    • If any more of the solute is added, then it will not dissolve completely.
  • The Solubility of a substance is is the amount of solute that dissolves in a given quantity of solvent at a specified temperature and pressure to produce a saturated solution.
  • Solubility is often expressed in in grams of solute per 100g of solvent.
    • Sometimes the solubility of a gas is expressed in grams per liter of the solution (g/L).
  • An Unsaturated Solutionis a solution that contains less solute than a saturated solution at a given temperature and pressure.
    • If additional solute is added to an unsaturated solution, it will dissolve until the solution is saturated.
  • Two liquids are Miscibleif they dissolve in each other in all proportions; that is, if any amount of one liquid will dissolve completely in any amount of the other liquid.
    • In miscible solutions, the liquid present in the larger amount is considered the solvent and the liquid present in the lesser amount is considered the solute.
    • An example of miscible liquids is water and ethanol
    • Liquids that are slightly soluble in each other, such as water and diethyl ether, are known as Partially Miscible.
  • Immiscible liquids are liquids that are insoluble in one another, such as oil and water.

Factors Affecting Solubility

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

Temperature

  • The solubility of most substances increases as the temperature of the solvent increases.
    • There are a few exceptions to this, such as in ytterbium sulfate, where its solubility decreases as the temperature of the solvent increases.
  • A Supersaturated Solutioncontains more more solute than it can theoretically hold at a given temperature.
    • This occurs when a saturated solution is prepared at certain temperature, and left undisturbed while it is allowed to cool to a lower temperature. Theoretically, the solubility of the solute is less at this temperature than it was at the first temperature, yet no crystallization occurs.
    • Crystallization can be initiated if seed crystal (very small crystal) of the solute is added to the solution.
    • An example of this is Rock Candy. During the production of Rock Candy, a Supersaturated Solution is prepared and seed crystals are added to initiate crystallization
  • The effect of temperature on the solubility of gases in liquid solvents is opposite that of solids.
  • The solubilities of most gases, such as oxygen, are greater in cold water than they are in hot water.
    • A consequence of this is that when industrial plants use lake water for cooling and then return it to the lake at a hotter temperature, it increases the temperature of the entire lake and can harm aquatic animal and plant life because the increase in temperature lowers the concentration of dissolved oxygen in the lake water. This is known as thermal pollution.

Pressure

  • Changes in pressure have little affect on the solubility of liquids and solids, but pressure strongly influences the solubility of gases.
  • Gas solubility increases as the partial pressure of the gas above the solution increases.
    • An example of this in carbonated drinks. The drinks are bottled/ canned under a high pressure of carbon dioxide gas, which forces large amounts of the gas into the solution. When the beverage is opened, the partial pressure of the gas above the liquid decreases, and the carbon dioxide gas immediately begins to escape from the bottle, lowering its concentration in the solution. This causes the fizzle that results when the drink is opened and the "flat" taste it develops when left open for a long period of time from lack of carbon dioxide in the solution.
  • Henry's Law states that at a given temperature, the solubility (S) of a gas in a liquid is directly proportional to the pressure (P) of the gas above the liquid.
    • In other words, as the the pressure of the gas above the liquid increases, the solubility of the gas increases, and as the pressure of the gas above the liquid decreases, the solubility of the gas decreases.
    • Equation: S1 / P1 = S2 / P2
      C12F11.GIFimages.jpeg

Above Images: Bryan Dextradeur

Group 5 (480-482, 483-486)

Lauren Murphy & Hannah Kumlin


Lauren Murphy
Molarity:
  • A substance can dissolve in some extent in a solvent to form a solution
  • Concentration-measure of an amount of a solute that is dissolved in solvent
  • Dilute solution-contains a small amount of solute
  • Concentrated solution-contains a large amount of solute
  • Molarity (M)- number of moles of solute that can be dissolved in a liter of solution

  • M= moles of solute/liters of solution
  • See the below link to do some practice molarity problems with some guy with a chill accent! J
http://www.youtube.com/watch?v=8oTqwBAvbnY


external image volumetric_soln_conc.jpg


Hannah Kumlin

Making Dilutions

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 before dilution = moles of solute after dilution.
Recall the equation used to find molarity: M=moles of solute / liters of solution.
- By rearranging this equation we get the formula for moles of solute: Moles of solute = Molarity (M) x liters of solution (V) and because the total number of moles of solute remains the same upon dilution, you can write this equation:
Moles of solute = M1 x V1 = M2 x V2

*M1 and V1 are the molarity and volume of the inital solution, and M2 and V2 are the molarity and volume of the diluted solution. Be sure to use the same unit measurements for both V1 and V2.
Example Problem (pg 484)-How many milliliters of aqueous 2.00M MgSO4 solution must be diluted with water to prepare 100.0 mL of aqueous 0.400M MgSO4?


Percent Solutions

Another way to describe the concentration of a solution is by the percent of a solute in the solvent.
The concentration of a solution in percent can be expressed in two ways: as the ratio of the volume of the solute to th volume of the solution or as the ratio of the mass of the solute to the mass of the solution.
Concentration in Percent (Volume/Volume)
If both the solute and the solvent are liquids, a convienent way to make a solution is to measure the volumes of the solute and the solution. The concentration of the solute is then expressed as a percent of the solution by volume.
The relationship between percent by volume and the volumes of solute and solution is:
Percent by volume (%(v/v)) = volume of solute/ volume of solution x 100%

Sample Problem (pg 485)What is the percent by volume of ethanol (C2H6O) in the final solution when 85 mL of ethanol is diluted to a volume of 250mL with water?

Concentration in Percent (Mass/Mass)

Another way to express the concentration of a solution is as a percent (mass/mass), which is the number of grams of solute in 100 grams of solution.

Percent by mass (%(m/m)) = mass of solute/ mass of solution x 100%
Information is often expressed as percent composition on food labels like the one below:
Food_label.jpgHannah Kumlin

Group 6 (487-490)

Chris Kelly


Colligative Properties of Solutions


Vapor-Pressure Lowering

-colligative property: a property that depends only upon the number of solute particles, and not upon their identity

3 important colligative properties
-vapor-pressure lowering
-boiling-point elevation
-freezing-point elevation

-the decrease in a solution’s vapor pressure is proportional to the number of particles the solute makes in solution

external image erlenmeyers.gif

Freezing-Point Depression


-freezing-point depression: the difference in temperature between the freezing point of a solution and the freezing point of the pure solvent
-The magnitude of 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


-boiling-point elevation: the difference in temperature between the boiling point of a solution and the boiling point of the pure solvent
-the magnitude of the boiling-point elevation is proportional to the number of solute particles dissolved in the solvent


Group 7 (491-496)

Ellie Kawa

Calculations involving Colligative Properties
Molality and Mole Fraction
Ellie Kawa pgs 491-496
  • the unit molality and mole fractions are two additional ways in which chemists express the concentration of a solution
  • molality- the number of moles of solute dissolved in one kilogram (1000g) of solvent
    • also known as molal concentration
    • molality= moles of solute/kilogram of solvent
Sample Problem
Using Solution Molality

How many grams of potassium iodide must be dissolved in 500.0 g of water to produce a 0.060 molal KI solution?

Analyze
knowns
  • mass of water= 500.0 g= 0.5000 kg
  • solution concentration= 0.060m
  • molar mass KI= 166.0 g/mol
Unknown
  • mass of solute= ? g KI
Calculate

0.5000 kg H2O*0.060 mol KI/1.000 kg H2O*166.0 g KI/1 mol KI= 5.0 g KI

  • mole fraction- the ratio of the moles of that solute to the total number of moles of solvent and solute
  • XA= nA/nA+nB
  • XB= nB/nA+nB

Sample Problem
Calculating Mole Fractions

Ethylene glycol (C2H6O2) is added to automobile cooling systems to protect against cold weather. What is the mole fraction of each component in a solution containing 1.25 mol of ethylene glycol (EG) and 4.00 mol of water?

Analyze
Knowns
  • moles of ethylene glycol (nEG)= 1.25 mol EG
  • moles of water (nH2O)= 4.00 mole H2O
Unknowns
  • mole fraction (XEG)= ?
  • mole fraction H2O (XH2O)= ?
Calculate
XEG= nEG/nEG+nH2O= 1.25 mol/1.25 mol+4.00 mol= 0.238
XH2O= nH2O/nEG+nH2O= 4.00 mol/1.25 mol+4.00 mol= 0.762
Freezing-Point Depression and Boiling-Point Elevation
  • the magnitudes of the freezing-point depression and the boiling point elevation of a solution are directly proportional to the molal concentration, when the solute is molecular, not ionic
    • change in the freezing temperature is the difference between the freezing point of the solution and the freezing point of the pure solvent
    • change in the boiling point is the difference between the boiling point of the solution and the boiling point of the pure solvent
    • ∆Tt=Kf*m
    • molal freezing-point depression constant (Kf)- equal to the change in freezing point for a 1-molal solution of a nonvolatile molecular solute
      • ∆Tb=Kb*m
      • molal boiling-point elevation constant-equal to the change in boiling point for 1-molal solution of a nonvolatile molecular solute
Sample Problem
Calculating the Freezing-Point Depression of a Solution

Antifreeze protects a car from freezing. It also protects it from over-heating. Calculate the freezing-point depression and the freezing point of a solution containing 100 g of ethylene glycol (C2H6O2) antifreeze in 0.500 kg of water.

Analyze
Knowns
  • mass of solute= 100 g C2H6O2
  • mass of solvent= 0.500 kg H2O
  • Kf for H2O= 1.86°C/m
  • ∆Tt=Kf*m
Unknowns
  • ∆Tt= ?°C
  • freezing point= ?°C
Calculate
moles of C2H6O2= 100 g C2H6O2* 1 mol/ 62.0 g C2H6O2= 1.61 mol
m= mol solute/kg solvent= 1.61 mol/0.500 kg= 3.22 m
Ellie Kawa
Ellie Kawa

∆Tt=Kt*m= 1.86°C/m*3.22 m= 5.99°C
Sample Problem
Calculating the Boiling Point of an Ionic Solution

What is the boiling point of a 1.50 m NaCl solution?

Analyze
Knowns
  • concentration= 1.50 m NaCl
  • K2 for H2O= 0.512°C/m
  • ∆Tb= Kb*m
Unknown
  • boiling point= ?°C
Calculate
∆Tb=Kb*m= 0.512 °C/m*3.00 m= 1.54°C


Group 8 (587-590, 590-593)

Emily Stewart & Jessen Foster


Properties of Acids and Bases page 587
Emily Stewart

Acids give foods - such as fruit or vinegar - a tart or sour flavor. Vinegar contains ethanoic, or acetic, acid. Lemons contain citrus acid.

Aqueous solutions of acids are called electrolytes, and they conduct electricity. Acids cause indicators (certain chemical dyes) to change color. Some electrolytes react with metals such as zinc or magnesium to produce hydrogen gas. When acids react with compound that contain hydroxide ions, water and a salt are produced.

Bases have a bitter taste, but it is generally dangerous to taste bases. Bases are also slippery.

Aqueous forms of bases are electrolytes that will cause indicators to change color. Water and salt are formed when a base containing hydroxide ions reacts with an acid.

Examples of bases are soap and milk of magnesia.

Arrhenius Acids and Bases page 588
Emily Stewart

In 1887, Swedish chemist Svante Arrhenius said that acids are hydrogen-containing compounds that ionize to make hydrogen ions in aqueous solutions & that bases are compounds that ionize to make hydroxide ions in aqueous solutions.

Arrhenius acids
monoprotic - contain one ionizable hydrogen (ex: nitric acid)
diprotic - contain two ionizable hydrogens (ex: sulfuric acid)
triprotic - contain three ionizable hydrogens (ex: phosphoric acid)

note:
- not all compounds that contain hydrogen are acids
- not all the hydrogens in an acid may be released as hydrogen ions
- only the hydrogens in very polar bonds are ionizable
external image HCl.jpg
Hydrogen chloride is polar covalent. It ionizes to form an aqueous solution of hydronium ions and chloride ions.

Common Arrhenius bases
Sodium hydroxide - ionic solid, dissociates into sodium ions and hydroxide ions in aqueous solution
Potassium hydroxide - ionic solid, dissociates into potassium ions and hydroxide ions in aqueous solution
Calcium hydroxide - not very soluble in water, even very dilute when saturated, low concentration of hydroxide ions
Magnesium hydroxide - less soluble than calcium hydroxide, low concentrations of hydroxide ion in suspensions in water


Jessen Foster, pages 590-593
Bronsted -Lowry Acids and Bases
- Arrhenius definition of acids and bases is very narrow
- in 1923, Bronsted and Lowry proposed a new definition
- the Bronsted -Lowry Theory defines an acid as a hydrogen-ion donor and a bases as a hydrogen-ion acceptor
- why ammonia is a base: very soluble in water, so it accepts a hydrogen ion from water

external image sat117002_0602.gif
Conjugate Acids and Bases
  • conjugate acid: particle formed when a base gains a hydrogen atom
  • conjugate base: particle that remains when an acid has donated a hydrogen ion
  • conjugate acid-base pair: two substances related by the loss or gain of a single hydrogen ion
  • amphotetic: property where a substance can act as both an acid and base

Lewis Acids and Bases
  • Gilbert Lewis proposed that an acid accepts a pair of electrons during a reactions, while a base donates a pair of electrons
  • Lewis acids accept a pair of electrons in a covalent bond
  • Lewis bases donate a pair of electrons in a covalent bond


Group 9 (594-603)

Allie Chabot

19.2- Hydrogen Ions and Acidity
Hydrogen Ions from Water
  • Water molecules are highly polar and are in continuous motion, even at room temperature.
  • A reaction in which water molecules produces ions is called a self-ionization of water.
  • In an aqueous solution hydrogen ions are always joined to water molecules as hydronium ions.
  • A neutral solution is an aqueous solution in which hydrogen ions and hydroxide are equal.
Ion Product Constant for Water
  • For aqueous solutions, the product of the hydrogen ion concentration and the hydroxide ion concentration equals 1.0X10-14.
  • The product of the concentration of the hydrogen ions and hydroxide ions in water is called the ion product constant of water.
  • Not all solutions are neutral.
  • An acidic solution is one in which the hydrogen ions is greater than the hydroxide.
  • A basic solution is one in which the hydrogen ions is less than the hydroxide. Basic solutions are also known as alkaline solutions.
The pH Concept
  • A pH scale ranges from 0-14; 7 is considered neutral, 0 is considered very acidic, and 14 is considered very basic.
  • The pH of a solution is the negative logarithm of the hydrogen ion concentration.
  • The pOH of a solution equals the negative logarithm of the hydroxide ion concentration.
external image phscale.gif(Allie Chabot)
Measuring pH
  • An indicator is an acid or base that undergoes dissociation 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 solution.
  • The acid form dominates the dissociation equilibrium at low pH and the base form dominates the equilibrium at high pH.
  • Indicators have certain characteristics that limit their usefulness.
  • A pH meter makes a rapid yet accurate pH measurement.
  • A pH meter is often easier to use than liquid indicators or indicator strips.


Group 10 (605-608, 608-611)

Neal McGovern & Shannon Leavey


Strengths of Acids and Bases
By: Neal McGovern

Strong and Weak Acids and Bases
  • Acids are classified as strong or weak depending on how they dissociate (or ionize) in water
  • Strong acids are completely ionized in water
  • Weak acids ionize only slightly in aqueous solution
external image Strong-Acid-Weak-Acid.gif

Acid Dissociation Constant
  • Strong acids are completely ionized in aqueous solutions
  • Weak acids ionize only slightly in aqueous solutions
  • the equilibrium constant can be written as (Keq)
  • For dilute solutions, the concentration of water is a constant, and it can be combined with Keq to give an acid dissociation constant
  • An acid dissociation constant (Ka) is the ratio of the concentration of the ionized form of the acid to the concentration of the nonionized
  • Weak acids have smaller Ka value than stronger acids

Base Dissociation Constant & Calculating
By: Shannon Leavey

Strong bases: dissociate completely into metal ions and hydroxide ions in aqueous solution.
Weak bases: react with water to form the hydroxide ion and the conjugate acid of the base.
base dissociation constant: the ratio of the concentration of the conjugate acid times the concentration of the hydroxide ion in the concentration of the base.

Concentration and Strength:
-strong and weak refer to the extent of ionization of an acid or base
- a small number of HCl molecules are present in a given volume
-solutions of ammonia can be dilute or concentrated
-
external image t0185-01.jpg
Shannon Leavey

Calculating Dissociation Constants:
-to find the K of a weak acid or the K of a weak base, substitute the measured concentratons of all the substances present at equlilbrium into the expression for K

K= H+ * A-/ HA

Group 11 (612-616)

Teresa Lynch


19.4 Neutralization Reactions- p. 612-616
By: Teresa Lynch

Acid- Base Reactions:
-Neutralization Reactions are reactions in which an acid and a base react in an aqueous solution to produce a salt and water. (This is a way to prepare pure samples of salts.)
-In general, the reaction of an acid with a base produces water and one of a class of compounds called salts.
- Salts are compounds consisting of an anion (from an acid) and a cation (from a base).

table_salt.jpg
NaCl, or table salt is an example of a salt compound containing an acid and base.
Photo: Teresa Lynch
Titration:
-The equivalence point is when the number of moles of hydrogen ions equals the number of moles of hydroxide ions when an acid ad base are mixed.
-You can determine the concentration of acid (or base) by performing a neutralization reaction. In a neutralization reaction you must use the correct acid-base indicator (like in the pH lab the various colors formed in the cabbage juice or on the litmus paper and the titration lab the pink color the acid turned).
cabbage_juice_indicator.jpg
This is an example of a cabbage juice indicator.
Photo: Teresa Lynch
-Titration is the process of adding a known amount of a solution of a known concentration to determine the concentration of another solution (with a known amount).
titration_image.gif
This is a flask that is being used in a titration lab, similar to the lab we did.
Photo: Teresa Lynch
-Standard solution is the solution of known concentration.
-End point is the point at which the indicator changes color.
-The point of neutralization is the end point of the titration.

Group 12 (618-622)

David Monti

Salt Hydrolysis
  • A salt consists of an anion from an acid and a cation from a base
  • Some salt solutions can be acidic or basic, but they are primarily neutral
  • Salt Hydrolysis – a process of determining the acidity in salts in which the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ions to water.
  • The acidity is dependent on the direction of the hydrogen ion transfer
  • 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
  • Basically, hydrolysis is the process of determining the acidity levels of salts through the removal and transferring of hydrogen ions from water molecules

salt-773845.jpg ph-buffer-solution-377332.jpg



Strong acid + strong base = neutral solution
Strong acid + weak base = acidic solution
Weak acid + strong base = basic solution


Buffers
  • Buffer – a solution in which the pH remains relatively constant when small amounts of acid or base are added
  • A buffer is a solution of a weak acid and one of its salts, or a solution of a weak base and one of its salts
  • A buffer solution is better able to resist drastic pH changes better than water
  • It reacts to the changing pH to form either acidic or basic solutions
  • Ethanoate buffers react less with acidic pH which is 1 example of buffer capacity
  • Buffer capacity – the amount of acid or base that can be added to a buffer solution before a significant charge in pH occurs.
  • There are many different forms of buffers with different capacity levels