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Acids and Bases

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Chloe H

on 1 June 2013

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Transcript of Acids and Bases

Acids, Bases and Salts Arrhenius' Theory Acids: Substance that
releases aqueous Hydrogen ions into water Bases Substance that
releases aqueous hydroxide ions into water Salts Product of a neutralization reaction (Acid/Base).
It is an ionic compound, but not an acid or a base. HF An ionic species with a formula that begins with "H" Hydrofluoric Acid HCL Hydrochloric Acid HNO3 Nitric Acid HI Hydroiodic Acid NaOH An ionic species with a formula that ends with "OH" Sodium Hydroxide LiOH Lithium Hydroxide Ba(OH)2 Barium Hydroxide Al(OH)3 Hydroiodic Acid NaCl Any ionic species that doesn't begin "H" or end with "OH" Sodium Chloride LiF Lithium Fluoride HgS Mercury Sulfide CuSO4 Copper Sulfate Characteristics of Acids -Electrolytes
-React with certain metals and produce gaseous H2
-Turn Litmus paper RED
-Sour flavour Characteristics of Bases -Electrolytes
-Slippery to touch (soapy)
-Turns Litmus paper BLUE
-Bitter Flavour Common Acids and Bases H2SO4 Sulphuric Acid HCl Hydrochloric
Acid HNO3 Nitric Acid CH3COOH Acetic Acid NaOH Sodium
Hydroxide KOH Potassium
Hydroxide NH3 Ammonia Dehydrant
(removes water from substances) Good electrical conducter Slow reaction with some metals Blackens some organic materials, as it dehydrates them Reacts exothermically with water Properties H2SO4 in concentrated form is 98% H2SO4 and 2% water Uses Produces sulphates Used to make
fertilizers, explosives, dyes, insecticides, detergents and
plastics used as the conductive substance in car batteries Used to absorb water and keep non-
aqueous solutions water free Properties Produces sulphates Good conducter In concentrated form chocking Used to absorb water and keep non-
aqueous solutions water free Uses Cleaning metals and bricks Catalyst Stomach acid is a diluted version of HCl. Acts as enzyme, which is a type of catalyst Removes calcium and magnesium carbonate Producing chlorides Properties When concentrated (16M), is at 69% in water Changes protein
colour - yellow
(turns skin
yellow) Very reactive, attacks a lot of metals Uses Used in the production of fertilizers, explosives and dyes Used to make nitrates Uses Used to create fabrics and plastics Preserves food Used in the production of acetates Properties Only conducts electricity when diluted Only influences very reactive metals Properties Absorbs H2O from air Absorbs gaseous CO2 from the air, forms a carbonate Corrosive on plant and animal tissue Mixing with water results in exothermic reaction Uses Create sodium based salts Makes soap and other cleaning products Used to make glass, paper products, plastic and aluminum Properties Same as NaOH, but has a lower melting temperature Uses Liquid soap Remove CO2 through absorption Conducter in alkaline batteries Make potassium based salts Properties When dissolved in water, creates an exothermic reaction Very soluble in water Clear, toxic, smells bad, corrosive Uses Create Nitric Acid Used to create explosives, fertilizers and synthetic fibers Refrigeration gas The Proton + H+ is called a Proton When an electron is removed, it creates and H+ ion, which has a very potent positive charge and therefore a very strong attraction to any negative charges H3O+ is called a hydronium ion,
or hydrated proton Example: HF + H2O -> F- + H3O+ Brønsted-Lowry Theory Acid Substance which donates a proton to another substance Base Substance which accepts a proton from another substance (proton donor) (proton acceptor) More Definitions MONOPROTIC ACID
Can only supply one proton
DIPROTIC ACID
Can supply two or less protons
TRIPROTIC ACID
Can supply three or less protons
POLYPROTIC ACID
General term. Can supply multiple protons Example: H2C2O4 + H2O -> HC2O4- + H30+ Example: HS- + H20 -> H2S + OH- base acid acid base Notice that H2O reacts as a base in the first example, but reacts as an acid in the second. Because of this, water is called an AMPHIPROTIC substance, or a substance that can act as either acid or base.
Amphiprotic substances come from polyprotic acids which have lost at least one proton
To identify an amphiprotic, check if the substance has a negative charge, or an easily removed hydrogen (usually situated at the end of the formula) Amphiprotic Acids H2C2O4 + H2O -> HC2O4- + H30+ The Bronsted-Lowry reaction has both an acid and a base on either side of the equation. An easy way to find this is to connect the similar substances and assume that they are opposite on either side of the equation, making: Acid Base Base Acid Following the same Bronsted-Lowry rules from before, the acids and bases are sorted as follows: Base Acid Acid Base HS- + H20 -> H2S + OH- Conjugate Acids and Bases Conjugate Pairs
(conjugate Acid-Base pair) A pair of substances that vary by only one proton Conjugate Acid The substance of the pair that has the excess proton Conjugate Base The substance that is missing the excess proton Conjugate Pair Conjugate Acid Conjugate Base Example: H2SO3 + H2O -> HSO3- + H3O+ H2SO3, HSO3- H2O, H3O+ H2SO3 HSO3- H30+ H2O To find the conjugate acid of a given substance, add a proton (H+) to the assumed base. To find the conjugate base of a given substance, remove a proton (H+) from the assumed acid Hydronium Ion: Remember, organic acids ending with a COOH group lose the proton from the end Strong Acids and Bases versus Weak Acids and Bases Strong Acid A strong acid is 100% ionized, that is completely dissociated Strong Base Same as the strong acid, a strong base is 100% dissociated Weak Acid A weak acid is less that 100% dissociated Weak Base Again, same as the weak acid, a weak base is less than 100% ionized Remember, equilibrium equations can only exist between weak acids and bases. The six strong acids are located at the top of the "Relative strengths of Bronsted-Lowry Acids and Bases" table which should be handed out. Notice that all reactions are one way, there are no reverse reactions, meaning that there will never be an equilibrium established. All of its ions are transferred to water.
Every strong acid has the same strength in an aqueous solution The last two reactions on the table are strong base reactions. They too also only react in one direction, backwards. The forward reaction will never occur.
Both bases have the same strength in aqueous solutions Weak acids are the substances on the left between HIO3 and H2O. These are reactions that lead to partially ionized solutions. The previously mentioned strong bases will never act as acids. The higher an acid is on the table, the stronger it is.
Weak acids are also weak electrolytes and the weaker the acid, the stronger the conjugate base.
Substances acting as acids in water reactions will always produce H3O+ Weak bases are located on the right side of the table from H2O to PO4^3-. The six strong acids are not included as they will never react as bases. The lower a base is on the table, the stronger it is.
Much like weak acids, weak bases are also weak electrolytes, and a very weak base will have a stronger conjugate acid.
Substances acting as bases in water reactions will always produce OH- Used to describe why all strong acids in aqueous solutions will dissociate to equivalent solutions of H3O+ and all strong bases will dissociate to equivalent solutions of OH- The Levelling Effect Equilibrium Constants for the Ionization of water Acidic Solution: [H3O+] > [OH-] Basic Solution: [H3O+] < [OH-] Neutral Solution: [H3O+] = [OH-] When a strong acid strong base react, a lot of heat is produced (59J) The equation for the ionization of water can be written as
H2O + 59J <-> H+(aq) + OH-(aq) Kw = [H3O+][OH-] = 1.00 * 10^-14 Ka and Kb Ka: Acid Ionization constant Kb: Base Ionization constant Ka appears on the table in the last column. The higher the Ka, the stronger the acid Kb however, does not appear on the table, but can be found using Ka, as will later be explained Ka and Kb equations are written by multiplying the concentration of the two products, then dividing it by the reactant. Pure water is not included example: HNO2 + H2O <-> NO- + H3O+ Ka = = 4.6 * 10^-4 [H3O+][NO-] [HNO2] Ka's and Kb's are related, multiplying them by each other results in Kw, which is constant. It is through this formula (Ka [conjugate acid] * Kb [conjugate base] = Kw) that Kb can be found. When two substances that can both act as either acid or base are mixed, a "proton competition" occurs. The side with the higher Ka value will win, meaning that more protons will be donated by it. Therefore:
The side with the weaker acid is favoured When working with weak acids, Ka can be used to find other facts about the acid Ka [H30] [HA] If given two out of the three of these, the third one can be found. R HNO2 + H2O <-> NO2- + H30+ I C E ecation nital hange nd 0 0 + X + X X X Find the [H3O+] of a 0.5M solution of HNO2. (Ka = 4.6 *10^-4) 0.5 - X 0.5 - X Ka= = = 4.6*10^-4 X= 1.52*10^-3 [NO2-][H3O+] HNO2 X^2 0.5 (to avoid quadratics, it is assumed that 0.5 - X = 0.5, as the number is so small) Working with Kb calculations and weak bases is similar to the Ka, but the Kb value must first be found, using previously shown steps. Also, instead of [H3O+], a [OH-] is used. pH and pOH Indicators Titrations Hydrolysis Buffers Buffers keep the pH of a relatively constant when small amounts of acids or bases are added pOH = -log[OH-] pH = -log[H3O+] [OH-]=antilog(-pOH) [H3O+]= antilog(-pH) pH + pOH = 14 All these calculations can now look like this: [H3O+] pH pOH [OH-] pH + pOH = 14 1.00*1-^14 = [H3O+][OH-] pH = -log[H3O+] pOH = -log[OH-] Sig Figs: In pH/pOH, only the sigfigs that fall after the decimal are counted. When the pH scale increase by one, the [H3O+] decreases by 10 Mixing strong acids and bases: This can result in a basic, a neutral or an acidic solution, depending on the amounts and strengths of the reactants Step 1



Step 2 Use dilution calculations to find the starting concentrations: [new] = M * old volume new volume If OH- is in excess, use the equation:


If H3O+ is in excess, use: [OH-]xs = [OH-]st - [H3O+]st [H3O+]xs = [H3O+]st - [OH-]st Weak organic base or acid that change colour between its conjugate acid and base When in a basic solution, the indicator is in its conjugate base form When in an acidic solution, the indicator is in its conjugate acid form When the solution changes colour, it signifies that it has reached the 'end point' or the 'transition point'.
This is not the same as the equivalence point. The equivalence point is where the stoichiometry of the the reaction is satisfied Used to find the concentration of an unknown solution, with known volume, by reacting a solution with known concentration and known volume. When the desired equivalence point is reached (as mentioned in the indicators section) the titration is over Percent Purity Some solutions are not 100% pure, and calculations can be performed to find its purity Actual concentration
Expected concentration * 100% When working with actual calculations, there are some experimental errors that can be made. Every experiment should be performed multiple times to leave room for error. If the mL added are not close to each other, discard the farthest one out. Then find the average of the remainder and use this in your calculations A 25 ml solution of 0.5 M NaOH is titrated until neutralized into a 50 ml sample of HCl. What was the concentration of the HCl?
Step 1 - Find [OH-]: 1:1 ratio, therefore [OH-] = 0.5 M
.Step 2 - Find the number of moles of OH-:
Molarity = # of moles/volume# of moles = Molarity x Volume
# of moles OH- = (0.5 M)(.025 L)
# of moles OH- = 0.0125 mol

Step 3 - Determine the number of moles of H+

When the base neutralizes the acid, moles H+ = moles OH-.
= 0.0125 moles.
Step 4 - Determine concentration of HCl

HCl -> Cl- + H+
moles of H+ = moles of HCl
Molarity = # of moles/volume
= (0.0125 mol)/(0.050 L)
= 0.25 M Example Hydrolysis is the reaction between a salt and a cation and/or anion. It will produce an acidic of basic solution.

All salts are considered to be 100% oxidized, so the calculations occur with the ions of the salt Spectator Ions Do not participate in the reactions. They are the conjugates of strong acids and bases
Spectator Cations: Alkali Metals, Alkaline Earth Metals
Spectator Anions: The conjugate bases of the first five strong acids (HSO4- acts as an acid elsewhere) Steps to find the reactions of a salt in water: 1. Find the dissociated ions of the salt
2. Discard the spectator ions
3. The remaining ions are acidic if found on the left side of the table, and basic if found on the right NaF -> Na+ + F-
Na is a spectator ion.
F- + H2O <-> HF- + OH-
OH- is produced, so the solution is BASIC Example 2 NH4Cl -> NH4+ + Cl-
Cl is a spectator ion.
NH4+ + H2O <-> NH3- + H3O+
H3O+ is produced, so the solution is ACIDIC Example 1
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