acid base powerpoint.ppt

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Properties of Acids

 Electrolytes, or break up into ions when _{Electrolytes, or break up into ions when}

aqueous, conduct electricity in aqueous

solutions

 Strong acid: good conductor_{Strong acid: good conductor}

 _{Weak acid: poor conductor}_{Weak acid: poor conductor}

 Sour Taste; citric acid and acetic acid_{Sour Taste; citric acid and acetic acid}

 _{React with most metals to produce H}_{React with most metals to produce H}₂₂_(g)_(g)  _{React with bases to form H}_{React with bases to form H}₂₂_{O and salt}_{O and salt}  React with indicators to change color_{React with indicators to change color}

– turn litmus _{turn litmus}redred

– phenolphthalein turns colorless_{phenolphthalein turns colorless} – pH values are 0.0-6.9_{pH values are 0.0-6.9}

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Properties of Bases

 Electrolytes conduct electricity in aqueous _{Electrolytes conduct electricity in aqueous}

solutions

 Strong base: good conductor_{Strong base: good conductor}

 Bitter Taste; caustic, feel slippery (soap)_{Bitter Taste; caustic, feel slippery (soap)}

 _{React with acids to form H}_{React with acids to form H}₂₂_{O and salt}_{O and salt}  React with indicators to change color_{React with indicators to change color}

– litmus turns _{litmus turns}blueblue

– phenolphthalein turns _{phenolphthalein turns}pinkpink --pH values are 7.1-14

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Arrhenius’ Theory on

Acids and Bases

Arrhenius Acid



Definition: a substance that ionizes in

water to give

hydrogen ions (proton) or

a Hydronium ion:

H

+ +

or H

3

O+



Ex: Hydrochloric acid, HCl (a strong

acid).



HCl (g)



H

+ +

(aq)

+ Cl

- -

(aq)



ACIDIC

solutions are formed when an

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Hydronium Ion



The H

++

interacts strongly with a lone pair of

electrons on the oxygen of a water

molecule. The resulting ion, H

₃₃

O

++

is called

the hydronium ion.



Ex: HCl + H

₂₂

O H

₃₃

O

+1+1

+ Cl

-1-1

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Arrhenius’ Bases



Definition: a substance that ionizes in water

to give

hydroxide ions: OH-

hydroxide ions:

OH-

Example: NaOH, NH

₃₃

NaOH (s)



Na+(aq) +

OH

--

(aq)

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-SALTS

SALTS



A salt is an ionic compound that does not

produce an H+ ion or OH- ion when

dissolved in water.



Salts are not acids or bases they are neutral

substances.



Ex.

 NaCl_NaCl

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Binary acids –two elements - H+ *

 Hydrogen _____ ide_{Hydrogen _____ ide}

 Hydrogen chloride_{Hydrogen chloride}

 Hydrogen fluoride_{Hydrogen fluoride}

 _{Hydrogen sulfide}_{Hydrogen sulfide}  Hydrogen bromide_{Hydrogen bromide}

 Hydro_____ic acid_{Hydro_____ic acid}

 Hydro_Hydrochlorchloric acidic acid  Hydro_Hydrofluorfluoric acidic acid  _Hydro_Hydro_sulfur_sulfur_{ic acid}_{ic acid}  Hydro_Hydrobrombromic acidic acid

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Ternary acids

3 elements: H+ and a

polyatomic ion

 Hydrogen _____ate_{Hydrogen _____ate}

 Hydrogen sulfate_{Hydrogen sulfate}

 Hydrogen chlorate_{Hydrogen chlorate}

 _{Hydrogen nitrate}_{Hydrogen nitrate}

 Hydrogen phosphate_{Hydrogen phosphate}

 _________ic acid_{_________ic acid}

 Sulfur_Sulfuric acidic acid  chlor_chloric acidic acid  _nitr_nitr_{ic acid}_{ic acid}

 Phosphor_Phosphoric acidic acid

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Hydrogen ________ite



_________ ous acid

 Hydrogen _____ite_{Hydrogen _____ite}

 Hydrogen _Hydrogennitrnitriteite  Hydrogen sulfite_{Hydrogen sulfite}

 _{Hydrogen chlorite}_{Hydrogen chlorite}

 _______ous acid_{_______ous acid}

 Nitr_Nitrous acidous acid  Sulfurous acid_{Sulfurous acid}

 Chlorous acid _{Chlorous acid}

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Table K and Table L



These two tables list the most common

acids and bases

.



Table K

lists the common acids, these acids

are the major acids you will need to be

familiar with.



The top 4 acids on table K are all

strong

acids

while the last 2 are

weak acids

.



Table L

lists common bases, these bases

are the major bases you will need to be

familiar with.



The top 3 bases on table L are

strong

bases

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Strengths of Acids



Strong acids: ionize completely which

means if 100 molecules dissolve in water all

100 will break up into ions



Examples:

HCl



H

++

+ Cl

-

-H

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Strengths of Acids



Weak acids:

ionize slightly or only a small

percentage will break up into ions



Example: acetic acid (vinegar)

CH

CH33COOH COOH  CH CH33COOCOO-- + H + H++

****note

****note if an organic compound ends in COOH it’s a _{if an organic compound ends in COOH it’s a}

weak acid!!!

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Strengths of Bases



Strong bases:

ionize completely



Example:

NaOH

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-Strengths of Bases

Strengths of Bases



Weak base: ionizes slightly



Example: ammonia

NH

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-Acid and Base Reactions

Acid and Base Reactions



You will need to be familiar with 2 reactions

that involve acids and bases.



The first reaction only involves an acid with

a metal.



The second reaction deals with reacting an

acid with a base to produce a salt and water

or neutralization.

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Acids with metals



Some metals will replace the hydrogen in

the acid and form hydrogen gas



Single Replacement Reaction (SRR)



Use Table J to decide if the metal is more

reactive or to see if the reaction will work.

Mg + HCl

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Neutralization Reaction

 In a neutralization reaction an Acid will react with _{In a neutralization reaction an Acid will react with}

a base to produce salt and water.

 Three ways we can express neutralization _{Three ways we can express neutralization}

reactions:

1.

1. Word equationWord equation 2.

2. Formula equationFormula equation 3.

3. Net ionic equationNet ionic equation

 A _Anet ionic equationnet ionic equation has only the ions that have has only the ions that have taken part in the reaction. In the reaction the ions

taken part in the reaction. In the reaction the ions

that do not change are omitted, these ions are

called

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Neutralization

HCl + NaOH

acid + base water + a salt

HOH + NaCl

Hydrochloric Acid + Sodium Hydroxide yields Water + Sodium Chloride

H+

(aq) + Cl-(aq) + Na+(aq) + OH-(aq) H2O + Na+(aq) + Cl-(aq)

H+

(aq) + OH-(aq) H2O

This is the net ionic equation for all neutralization reactions this

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Try This One

H

₂

SO

₄

+ 2NaOH

2H

₂

O + Na

₂

SO

₄

2H+ + SO

4-2 + 2Na+ + 2OH- 2H2O + 2Na

+ + SO 4-2

H+ + OH- H 2O

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Concentration of H

++

in Acids



A. 1.0 M HCl (monoprotic acid)

– HCl _HCl H H+ + + Cl_{+ Cl}-

-– Produces 1 H_{Produces 1 H}++ so the concentration of H_{so the concentration of H}++ is 1.0 _{is 1.0}

M

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Concentration of H

++

in Acids



B. 1.0 M H

₂₂

SO

₄₄

(diprotic acid)

– HH22SOSO44  2H 2H+ + + SO+ SO 4 4-2-2

– Produces 2 H_{Produces 2 H}++ so the concentration of H_{so the concentration of H}++ is _is

2(1.0 M)

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Concentration of OH

--

in Bases



A. 1.0 M KOH

– KOH KOH  K K++ + OH + OH-

-– Produces 1 OH_{Produces 1 OH}-- so the concentration of OH_{so the concentration of OH}-- is _is

1.0 M.

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Concentration of OH

--

in Bases



B. 1.0 M NaOH

– NaOH _NaOH Na Na++ + OH_{+ OH}-

1.0 M.

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Concentration of OH

--

in Bases



C. 1.0 M Ca(OH)

₂₂

– Ca(OH)Ca(OH)₂₂  Ca Ca+2+2 + 2OH_{+ 2OH}-

2(1.0 M).

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Concentration of OH

--

in Bases



D. 1.0 M Mg(OH)

₂₂

– Mg(OH)Mg(OH)₂₂  Mg Mg+2+2 + 2OH_{+ 2OH}-

2(1.0 M).

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Titration

 An Acid-Base titration is a lab technique used find the unknown concentration of a specific volume for an acid or a base by neutralizing it with a specific volume of an acid or base of a known concentration.

 During a titration you deliver known volumes of a base to a known volume of an acid until it is neutralized or just slightly basic.

 This is also known as the endpoint of a titration.

 End point_{End point}: point at which the solution : point at which the solution changes from acidic to “just” basic.

changes from acidic to “just” basic.

 Using the Acid-Base titration formula listed on _{Using the Acid-Base titration formula listed on} Table T

Table T you can solve for your unknown you can solve for your unknown concentration.

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Acid- Base Titrations

 Acids and bases neutralize each other when the concentration of the HAcids and bases neutralize each other when the concentration of the H++_{equals the concentration of the OH}_{equals the concentration of the OH}--_._.

M

M_a_aVV_a_a = M = M_b_bVV_b_b(Table T)(Table T)

Ma = molarity of H+

V_a = volume of acid M_b = molarity of OH

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Titrations (Neutralization)

Problems

 Ex. What volume of 0.50M HCl is required to neutralize Ex. What volume of 0.50M HCl is required to neutralize 100mL of 2.0M NaOH?

100mL of 2.0M NaOH?

M

M_aVV_a_a = M = M_b_bVV_b_b M

M_a_a = .5M = .5M V

V_a_a = ?? = ?? M

M_b_b = 2.0M = 2.0M V

V_b_b = 100mL = 100mL

Plug in and solve

0.5(x) = 2.0(100)

X = 400mL HCl

Try the rest on your own!!!

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Acidity and Alkalinity of Solutions



Although water is a covalently bonded

substance it does

ionize

to a very small

extent as shown by the equation below.

H

₂₂

O H

++

+ OH

-

-

It can be seen that in pure water the

[H+] = [OH-].



When one of the ion concentrations

increases the other will decrease.

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Acidity and Alkalinity of Solutions

(continued)

 So if you dissolve HCl in water the concentration _{So if you dissolve HCl in water the concentration}

of [H+] will

of [H+] will increaseincrease and the [OH-] will and the [OH-] will decreasedecrease

making the

making the solution acidicsolution acidic. .

 When NaOH is dissolved in water the _{When NaOH is dissolved in water the}

concentration of [OH-]

concentration of [OH-] increasesincreases and the [H+] and the [H+]

decreases

decreases making the making the solution basicsolution basic. .

 So when _{So when}[H+] > [OH-][H+] > [OH-] the solution is the solution is acidicacidic and and when

when [H+] < [OH-][H+] < [OH-] the solution is the solution is basicbasic..

 The terms acidity and alkalinity (or basicity) refer _{The terms acidity and alkalinity (or basicity) refer}

to the strength of the acid or base in terms of H+

and OH- concentrations.

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Summary of Aqueous Solutions

 Autoionization of water_{Autoionization of water}

HOH H

HOH H++ + OH + OH-

- In any aqueous [H+] x [OH-]_{In any aqueous [H+] x [OH-]}= 1.0x10= 1.0x10-14-14

 Neutral Solution _{Neutral Solution}

[H+] = [OH-][H+] = [OH-] 1.0x10

1.0x10-7-7M = 1.0x10_{M = 1.0x10}-7-7M_M

 Acidic Solution_{Acidic Solution}

[H+] > [OH-][H+] > [OH-] 1.0x10

1.0x10-1-1M > 1.0x10_{M > 1.0x10}-13-13M_M

 Basic Solution_{Basic Solution}

[H+] < [OH-]

1.0x10

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pH Scale



A scale, called the pH scale, has been

developed to express

[H+]

as a number

from 0 to 14.



A pH of 0 is

strongly acidic

, a pH of 7 is

neutral

, and a pH of 14 is

strongly basic



The pH scale is logarithmic which means

each change of a single pH unit signifies a

tenfold

change in the concentration of the

[H+]

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pH Scale

(continued)



Thus a solution with a pH of 5 is ten times

greate

r in [H+] than a solution with a pH of 6



Because the [H+] and [OH-] are directly

related, a pH change of one unit represents

a

tenfold increase or decrease

of both

[H+] and [OH-]



As the concentration of the

hydrogen ion

increases

the

hydroxide ion

concentration decreases

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pH Scale

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pH (continued)

additional info about pH and [H+]

 To calculate pH you can take the _{To calculate pH you can take the}– log([H+])– log([H+]) for for example if the [H+] = 1x10

example if the [H+] = 1x10-2 -2 the ph will he equal to _{the ph will he equal to}

– log (1x10

– log (1x10-2-2) which calculates to a pH of 2._{) which calculates to a pH of 2.}

 Calculate the pH of the following solutions _{Calculate the pH of the following solutions} when given the [H+]

when given the [H+]

 1. [H+] = 1x10_{1. [H+] = 1x10}-11-11

 Ans: pH =11_{Ans: pH =11}

 2. [H+] = 1x10_{2. [H+] = 1x10}-4-4

 Ans: pH =4_{Ans: pH =4}

 3. [H+] = 1x10_{3. [H+] = 1x10}-9-9

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Acid – Base Indicators

 An indicator is a substance that _{An indicator is a substance that}changes its colorchanges its color

when it gains or loses a proton or an H+ ion.

 There are several different indicators and they _{There are several different indicators and they}

change differently when exposed to different

change differently when exposed to different pH pH values

values. .

 _{On your reference tables there is a list of common}_{On your reference tables there is a list of common} acid base indicators that show the color changes

acid base indicators that show the color changes

at varying pH values

See Table M See Table M

 You can use more than one indicator to estimate _{You can use more than one indicator to estimate}

the approximate

the approximate pHpH of an unknown substance of an unknown substance from the color changes

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An Additional Theory on the

Definition of Acids and Bases

 Another theory on acids and bases defines an acid as any _{Another theory on acids and bases defines an acid as any}

substance that

substance that donates a proton or an H+ ion when donates a proton or an H+ ion when dissolved in water.

dissolved in water.

 For example_{For example} HClHCl is an acid because it donates an H+ ion is an acid because it donates an H+ ion to H

to H₂₂O when it dissolves.O when it dissolves.

 The theory also states that a base is any substance that _{The theory also states that a base is any substance that}

accepts a proton or an H+ ion

accepts a proton or an H+ ion. .

 _{So in the same situation when HCl is dissolved in H}_{So in the same situation when HCl is dissolved in H}₂₂_{O the}_{O the}

H

H₂₂O acts as a O acts as a base because it base because it accepts an H+ ion and accepts an H+ ion and

becomes H

becomes H₃₃OO++

 So in summary an acid can be any substance that So in summary an acid can be any substance that donates donates

an H+ ion or a proton and a base can be any substance

that