Acids And Bases

Acids Acids AndAnd BasesBasesAcids Acids AndAnd BasesBases

Section 19.1Acid-Base Theories

• OBJECTIVES:

–Define the properties of acids and bases.

Section 19.1Acid-Base Theories

• OBJECTIVES:

–Compare and contrast acids and bases as defined by the theories of: a) Arrhenius, b) Brønsted-Lowry, and c) Lewis.

Class question• Where can acids

be found?– Sodas– Stomach– Vinegar– Citrus fruits

• Where can bases be found?– Soap– Drano– Antacid tablets– Windex– detergent

Properties of Acids• Taste sour• React with bases • Litmus paper test – turn blue

litmus paper red• Electrolytic – conduct electricity

–Can be strong or weak electrolytes in aqueous solution

Properties of Acids• They have a pH of less than 7 (more

on this concept of pH in a later lesson)

• How do you know if a chemical is an acid?– It usually starts with Hydrogen.– HCl, H2SO4, HNO3, etc. (but not water!)

Acids Affect Indicators, by changing their color

Blue litmus paper turns red in contact with an acid (and red paper stays red).

Acids have a

pH less

than 7

Properties of Bases• Taste bitter• Feels slippery• React with acids• Litmus paper test – turn red litmus

paper blue• electrolytic

Bases Affect Indicators

Red litmus paper turns blue in contact with a base (and blue paper stays blue).

Phenolphthalein turns purple in a base.

Bases have a

pH greater than 7

Acid NomenclatureAcid NomenclatureAcid NomenclatureAcid Nomenclature

Nomenclature of Acids• Acids are composed of a(n)

________________ followed by a(n) _______

Hydrogen ion (H+)

anion

Ex: H+ + Cl1-

H+ + SO42-

HCl

H2SO4

Binary AcidsH+ + anion

• H+ + anion with –ide ending acid name is __________________

HCl anion? _______

acid name ________________

Hydro _____ic acid

Hydrochloric acid

chloride

Binary Acids H+ + anion

• H+ + anion with –ide ending acid name is __________________

HF anion? _______

acid name ________________

Hydro _____ic acid

Hydrofluoric acid

fluoride

OxyacidsH+ + anion

• H+ + anion with –ate ending acid name is __________________

HNO3 anion? _______

acid name ________________

_____ic acid

nitric acid

nitrate

Oxyacids H+ + anion

• H+ + anion with –ate ending acid name is __________________

H2SO4 anion? _______

acid name ________________

_____ic acid

sulfuric acid

sulfate

Oxyacids H+ + anion

• H+ + anion with –ite ending acid name is __________________

HNO2 anion? _______

acid name ________________

_____ous acid

nitrous acid

nitrite

Oxyacids H+ + anion

• H+ + anion with –ite ending acid name is __________________

HClO2 anion? _______

acid name ________________

_____ous acid

chlorous acid

chlorite

Writing acid formulas• Hydrobromic acid

anion? ___________ formula ______• Acetic acid

anion? ___________ formula ______• Nitrous acid

anion? ___________ formula ______

HBr

HC2H3O2

HNO2

Bromide (Br1-)

acetate(C2H3O21-)

nitrite (NO21-)

Base NomenclatureBase NomenclatureBase NomenclatureBase Nomenclature

Nomenclature of Bases• Bases are composed of a(n)

_______ followed by

a(n) ________________

cation

hydroxide (OH1-)

Writing Base Names• Rule: name the cation and add

“hydroxide”

• NaOH

• Mg(OH)2

• Fe(OH)3

sodium hydroxide

magnesium hydroxide

Iron (III) hydroxide

Memorize: NH3 = ammonia

Writing base formulas• potassium hydroxide

cation? ______ formula ______• Calcium hydroxide

cation? ______ formula ______• Aluminum hydroxide

cation? ______ formula ______

KOH

Ca(OH)2

Al(OH)3

K+

Ca2+

Al3+

Ions In Solution• Why are some solutions acidic,

basic, or neutral?

It depends on number of H+ and OH- ions present.

Ions In Solution• Acidic solution – contain more H+

ions than OH- ions

4000 H+ and 0 OH- is acidic1000 H+ and 500 OH- is acidic5 H+ and 3 OH- is acidic

Ions In Solution• Basic Solution – contain more OH-

ions than H+ ions

4000 OH- and 0 H+ is basic1000 OH- and 500 H+ is basic5 OH- and 3 H+ is basic

Ions In Solution• Neutral Solution – equal amounts

of H+ and OH- ions

4000 OH- and 4000 H+ is neutral1000 OH- and 1000 H+ is neutral5 OH- and 5 H+ is neutral

Self Ionization of Water• Proper ionization

H2O + H2O hydronium ion

O

H H

H

H

O

H H O

H H

O++

H3O+ + OH-

Self Ionization of Water• simplified version

H2O H+ + OH-

Acid-Base TheoriesAcid-Base TheoriesAcid-Base TheoriesAcid-Base Theories

Types of Acids/Bases• Arrhenius Model

• Bronsted-Lowry Model

• Lewis Model

Svante Arrhenius• He was a Swedish chemist (1859-1927),

and a Nobel prize winner in chemistry (1903)

• one of the first chemists to explain the chemical theory of the behavior of acids and bases

• Dr. Hubert Alyea (professor emeritus at Princeton University) was the last graduate student of Arrhenius.

Hubert N. Alyea (1903-1996)

Svante Arrhenius (1859-1927)

Arrhenius Model of Acids and Bases

• Arrhenius Acids– Defn: contain H+ and ionizes to form H+

– ExamplesHCl

HNO3

H+ + Cl-

H+ + NO3-

makes solutionACIDIC

Arrhenius Model of Acids and Bases

• Arrhenius Bases– Defn: – contain OH- and ionizes to produce OH- ions– Examples

NaOH

Ca(OH)2

Na+ + OH-

Ca2+ + 2 OH-

makessolutionBASIC

Flaw with Arrhenius model

• Not all bases contain hydroxide• Ex: ammonia (NH3) is basic

– According to Arrhenius, since ammonia can NOT produce OH- it is NOT a base

• Therefore a new type of acid/base must be determined

Gilbert Lewis (1875-1946)

Lewis Acids and Bases

• Gilbert Lewis focused on the donation or acceptance of a pair of electrons during a reaction

• Most general of all 3 definitions; acids don’t even need hydrogen!

Lewis Model• Lewis acid

– an atom, ion, or molecule that accepts an electron pair to form a covalent bond

• Lewis base– An atom, ion, or molecule that donates

an electron pair to form a covalent bond

Lewis Model• Lewis acid-base reaction

– The formation of one or more covalent bonds between an electron-pair donor and an electron-pair acceptor

Example– AcidsAcids are electron pair acceptors. – BasesBases are electron pair donors.

Lewis base

Lewis acid

Johannes Brønsted Thomas Lowry (1879-1947) (1874-1936) Denmark England

Bronsted-Lowry Model

• Bronsted-Lowry Acid– Defn: proton/H+ donor

• can give H+ to another species

• Bronsted-Lowry Base– Defn: proton/H+ acceptor

• can take H+ from another species

**Acids and bases always come in **Acids and bases always come in pairs.**pairs.**

Bronsted-Lowry Model• REMEMBER!!!! REMEMBER!!!!

acids donate, bases accept protons

Ashley does boys always

Bronsted-Lowry Model

• Examples

HCl + H2O Cl- + H3O+

What is happening here?

Acid (donate

s proton)

Base (accepts proton)

Which is the acid? base?

HCl is an acidHCl is an acid — — wwhen it dissolves in water, it gives it’s proton to water.

Water is a base—when the HCl gives up the proton, water accepts it to form the hydronium ion

Bronsted-Lowry Model• Examples

NH3 + H2O NH4+ + OH-

What is happening here?

Acid (donates proton)

Base (accepts proton)

Which is the acid? base?

Why Ammonia is a Base• Ammonia can be explained as a

base by using Brønsted-Lowry:

NH3(aq) + H2O(l) ↔ NH41+

(aq) + OH1-(aq)

Ammonia is the hydrogen ion acceptor (base), and water is the hydrogen ion donor (acid).

Conjugate Acid/Base Pairs

• Conjugate acid – new species produced when base gains H+ ion

• Conjugate base – new species produced when acid donates H+ ion

Thus, a conjugate acid-base Thus, a conjugate acid-base pair is related by the pair is related by the loss or loss or gaingain of a of a single hydrogen ionsingle hydrogen ion..


• general Bronsted-Lowry reaction

acid + base conj. acid + conj. base

conj. acid/base pair


Every acid has a conjugate base.Every base has a conjugate acid.


• Examples

HNO3 + H2O H3O+ + NO3-

What is the acid? base?

acid base What is the conjugate acid/base?

C.A. C.B.




• Examples


What is the acid? base?

acidbaseWhat is the conjugate acid/base?

C.B.C.A.



Conjugate acid-base pairs

• What is the conjugate base of:H2SO4 _________ H3O+ ________

• What is the conjugate acid of:HPO4

2- _________ OH1- ________

HSO41-

H2PO41-

H2O

H2O

How can H2O be both acid and base?

What is the conjugate base?

Acid Conjugate BaseH2SO4

HPO42-

NH41+

H3O1+

H2O

HSO41-

PO43-

NH3

H2O

OH-

How can H2O be both acid and base?

Amphoteric• Defn – substance that can act as

both acids and bases

HNO3 + H2O H3O+ + NO3-


base

acid

Water is amphoteric b/c it acts as a base in one reaction and acts as an acid in the second

Is H2O abase or acid?

Is H2O abase or acid?

Mono-, Di-, Triprotic Acids• Defns

– monoprotic (HA) – one ionizable protonex: HF, HCl, HBr (= normality is 1)

– diprotic (H2A) – two ionizable protons

ex: H2SO4, H2CO3 (= normality is 2)

– triprotic (H3A) – three ionizable protons ex: H3PO4, H3BO3 (= normality is 3)

Polyprotic Acid Ionization

• Always forms ONE H+

H3PO4

H2PO41-

HPO42-

H2PO41- + H+

HPO42- + H+

PO43- + H+

Strength of Acids and Bases

• Acid/base strength is based on the degree to which they ionize

1) strong ()

2) weak ()

Strong Acid/Base• Defn – acid or base that

completely ionizes

HA H+ + A-

XOH X+ + OH-

100%

ionization

every single HA molecule ionizes into H+ and A-

100%

ionization

Strength• Strong Acid/BaseStrong Acid/Base

– 100% ionized in water– strong electrolyte

- +

HCl

HNO3

H2SO4

HBr

HI

HClO4

NaOH

KOH

Ca(OH)2

Ba(OH)2

Strong Acid• Illustration

H A

H A

H A

H A

H A

H A

+ -

+ -

+ -

All break into ions

+

+

+

6 Strong Acids• HCl – hydrochloric acid• HBr – hydrobromic acid• HI – hydroiodic acid

• HClO4 – perchloric acid

• H2SO4 – sulfuric acid

• HNO3 – nitric acid

Strong Bases• Group I and II metal hydroxides

LiOH

NaOH

KOH

RbOH

Mg(OH)2

Ca(OH)2

Sr(OH)2

Ba(OH)2

No need tomemorizeexact ones

Weak Acid/Base• Defn – acid or base that partially

ionizes

HA H+ + A-

XOH X+ + OH-

partial

ionization

not all will ionize; the weaker it isthe less it ionizes

Strength

• Weak Acid/BaseWeak Acid/Base– does not ionize completely– weak electrolyte– Dissociates into both ions and molecules

- +

HF

CH3COOH

H3PO4

H2CO3

HCN

NH3

Weak Acid• Illustration

H A

H A

H A

H A

H A

H A

+ -

Only some break into ions

+

What are the weak acids and bases?

• The ones that are NOT strong

Strong or weak, concentrated or diluted• For acids and bases, it is important

to distinguish between concentrated and dilute from strong and weak. The words _________ and __________have different meanings. Similarly, ___________ and ___________ are not the same either.

strong

weak

concentrated

dilute

Strong or weak, concentrated or diluted• Strong and weak refer to

____________________________

• Concentrated and dilute refer to____________________________

how much substance ionizes

how much solute is present

Example• 1 M HCl

• 12 M HCl

• 1 M H2CO3

• 12 M H2CO3

Strong and dilute

Strong and concentrated

weak and dilute

weak and concentrated

Ion Product Constant for Water (Kw)

• Defn: equilibrium value for self ionization of water (H2O H+ + OH-)

• Formula Kw = [H+][OH-] = 1 x 10-14

ALWAYSALWAYSALWAYS

Ion Product Constant for Water (Kw)

• in pure water [H+] = [OH-] = 1 x 10-7

• in non pure water (acidic/basic conditions), value of [H+] and [OH-] differ

Remember pure water is neutral

But still [H+][OH-] = 1 x 10-14

Is solution acidic, basic, or neutral?

• Acidic

• Basic

• neutral

[H+] > [OH-]

[H+] < [OH-]

[H+] = [OH-]

pH = -log[ H+]

The pH Scale

0

7INCREASING

ACIDITY NEUTRALINCREASING

BASICITY

14

pouvoir hydrogène (Fr.)“hydrogen power”

The pH Scale

pH of Common SubstancespH of Common SubstancespH of Common SubstancespH of Common Substances

Relation of pH and pOH

• pH + pOH = 14

If given one variable, subtract to find the other

The pH Scale

pH = -log[H+]

pOH = -log[OH-]

pH + pOH = 14

pHc) change one pH unit

represents a ten fold change in strength- ex: pH = 3 vs pH = 4

pH 3 is 101 or 10 times more acidic- ex: pH = 7 vs pH = 10

pH 7 is 103 or 1000 times more acidic

Overall Relationship

pH pOH

[H+] [OH-]

Overall Relationship

The pH Scale

• What is the pH of 0.050 M HNO3?

pH = -log[H+]

pH = -log[0.050]

pH = 1.3

Acidic or basic?Acidic

Calculating [H+] and [OH-]

• Find the hydroxide ion concentration of 3.0 10-2 M HCl.

[H+][OH-] = 1.0 10-14

[3.0 10-2][OH-] = 1.0 10-14

[OH-] = 3.3 10-13 M

Acidic or basic?Acidic

• What is the molarity of HBr in a solution that has a pOH of 9.6?

pH + pOH = 14

pH + 9.6 = 14

pH = 4.4

Acidic

pH = -log[H+]

4.4 = -log[H+]

-4.4 = log[H+]

[H+] = 4.0 10-5 M HBr

Calculating [H+] and [OH-]Calculating [H+] and [OH-]

• A Ca(OH)2 solution has a pH of 8.0. Determine the [H+], [OH-], and [Ca(OH)2] for the solution.

pH = - log [H+]

log [H+] = -pH

[H+] = antilog (-pH) = antilog (-8.0) = 1 x 10-8 M H3O+


Antilog is the same thing as 10^(x)

[H+] [OH-] = 1 x 10-14 M2

[OH-] = 1 x 10-14 M2 = 1 x 10-14 M2

[H+] 1 x 10-8 M

[OH-] = 1 x 10-6 M


• Ca(OH)2 Ca2+ + 2OH-

( 1 x 10-6 M)

1 x 10-6 mol OH- 1 mol Ca(OH)2 = 5 x 10-7 mol/L Ca(OH)2

Liter 2 mol OH-

[Ca(OH)2] = 5 x 10-7 M


Sample problem #1• Calculate the pH of a solution with

[H+] = 3.0 x 10-6 M.

pH = -log [H+] = - log [3.0 x 10-6]

= 5.52

Sample problem #2• Calculate the pH of a solution with [OH-] =

8.2 x 10-6.

pOH = -log[OH-]= -log [8.2 x 10-6]= 5.09

pH + pOH = 14pH + 5.09 = 14

pH = 8.91

[OH-] pOH pH

Sample problem #2• Calculate the pH of a solution with

[OH-] = 8.2 x 10-6.

[H+][OH-] = 1 x 10-14

[H+][8.2 x 10-6] = 1 x 10-14

[H+] = 1.22 x 10-9

pH = -log[1.22 x 10-9] = 8.91

[OH-] [H+] pH

Sample problem #3• What is the [H+] of a solution with

pH = 2?pH = -log[H+][H+] = 10-pH

= 10-2

= 0.01 M

Sample problem #4• (i) What is the [OH-] of a solution

with pOH = 3.7?

[OH-] = 10-pOH

[OH-] = 10-3.7

= 2 x 10-4 M

Sample problem #4

pH + pOH = 14pH + 3.7 = 14

pH = 10.3

•(ii) What is the pH and the [H+] if the pOH is 3.7?

pOH pH [H+]

[H+] = 10-pH

= 10-10.3

= 5 x 10-11 M

Reaction between acids and bases

• Neutralization (defn) – reaction of acid and base to form a salt and water– The reaction is a double replacement

• Salt (defn) – ionic compound made of – cation from base and – anion from acid

Reaction between acids and bases

• Ex reaction

Mg(OH)2 + HCl MgCl2 + H2Obase acid salt water

Mg2+ + OH- H+ + Cl-Mg2+ - cation from base

Cl- - anion from acid

NeutralizationACID + BASE SALT + WATER

HCl + NaOH NaCl + H2O

HC2H3O2 + NaOH NaC2H3O2 + H2O

– Salts can be neutral, acidic, or basic.

– Neutralization does not mean pH = 7.

weak

strong strong

strong

neutral

basic

Ex problems• i) What is the salt formed from

sulfuric acid (H2SO4) and potassium hydroxide (KOH)?

base cation?acid anion?

K+

SO42-

What is salt? K2SO4

Ex problems• ii) What is the salt formed when

Al(OH)3 and HBr react?

base cation?acid anion?

Al3+

Br-

What is salt? AlBr3

Is salt solution acidic, basic, or neutral?

• a) strong acid + strong base

• b) strong acid + weak base

• c) weak acid + strong base

Neutral salt

acidic salt

basic salt

Ex problem• Determine if salt solution is

acidic, basic, or neutral.

a) LiBracid?base?

HBr

LiOH

(strong acid)(strong base)

Salt is ____________NEUTRAL

Li+ + Br-

Ex problem• Determine if salt solution is

acidic, basic, or neutral.

b) Fe(NO3)3

acid?base?

HNO3

Fe(OH)3

(strong acid)(weak base)

Salt is ____________ACIDIC

Fe3+ + NO3-

B. Titration• Titration

– Analytical method in which a standard solution is used to determine the concentration of an unknown solution.

standard solution

unknown solution

• Equivalence point (endpoint)– Point at which equal amounts

of H3O+ and OH- have been added.

– Determined by…• indicator color change

B. Titration

• dramatic change in pH

Titration

•dramatic change in pH

B. Titration

moles H3O+ = moles OH-

MV n = MV n

M: MolarityV: volumen: # of H+ ions in the acid

or OH- ions in the base

Titration• 42.5 mL of 1.3M KOH are required to

neutralize 50.0 mL of H2SO4. Find the molarity of H2SO4.

H3O+

M = ?V = 50.0 mLn = 2

OH-

M = 1.3MV = 42.5 mLn = 1

MV# = MV#M(50.0mL)(2)

=(1.3M)(42.5mL)(1)

M = 0.55M H2SO4

Subscript of H or OH from formulas!

Acids And Bases

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