Acids
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Acids. H +. Acid – a compound that produces ions when dissolved in Examples: Vinegar – Lemon juice – Tea – Ant venom –. water. Acetic acid. Citric acid. Tannic acid. Formic acid. Acids and Bases. Properties of Acids. Sour. taste

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Acids

Acids

H+

  • Acid – a compound that produces ions when dissolved in

  • Examples:

    • Vinegar –

    • Lemon juice –

    • Tea –

    • Ant venom –

water

Acetic acid

Citric acid

Tannic acid

Formic acid


Acids and bases

Acids and Bases


Properties of acids

Properties of Acids

Sour

  • taste

  • Turns litmus paper

  • Reacts with metals to form gas

  • solutions of acids are

    (must be mixed with water!)

  • Reacts with to form and

Corrosive

red

H2

Aqueous

electrolytes

bases

H2O

salt


Properties of acids1

Properties of Acids

  • Sugar, corn syrup, modified corn starch, citric acid, tartaric acid, natural and artificial flavors, yellow 5, yellow 6, red 40, blue 1

  • What ingredients make these…

    so sour?


Naming binary acids

Naming Binary Acids

  • H + one other element

  • Begin with

  • Use the of the element name

  • Add the suffix

  • HCl

“hydro-”

root

“-ic”

chlor

ic

acid

hydro


Naming binary acids1

Naming Binary Acids

Hydrobromic acid

  • HBr 

  • HI 

  • HF 

Hydroiodic acid

Hydrofluoric acid


Naming ternary acids

Naming Ternary Acids

  • H + polyatomic ion

  • Begin with ion without

    the

  • Add suffix if there was an

  • Add suffix if there was an

  • HNO3

polyatomic

ending

-ate

-ic

-ous

-ite

Nitrate

Nitric acid


Naming ternary acids1

Naming Ternary Acids

Chloric acid

  • HClO3 

  • H3PO3 

  • H2CO3 

Phosphorous acid

Carbonic acid


Strength of acids

Strength of Acids

Strong

ionize

  • acids – completely

    in water (create a lot of )

  • 3 binary acids

  • Ternary acids

    Strong if # of atoms - # of atoms ≥

    H2SO4HNO3

H+

HBr

HCl

HI

2

O

H

3 – 1 = 2

4 – 2 = 2


Strength of acids1

Strength of Acids

Weak

slightly

  • acids – ionize only in

    solution

  • Binary acids – all others not listed above

  • Ternary acids

    Weak if # of atoms - # of atoms ≥

    H3PO3HNO2

aqueous

1

O

H

2 – 1 = 1

3 – 3 = 0


Bases

Bases

OH-

  • Base – a compound that produces ions

    when dissolved in

  • Examples:

    • Milk of Magnesia – neutralizes stomach acid

    • Drain cleaner–

(hydroxide)

water

Magnesium hydroxide

Sodium hydroxide


Properties of bases

Properties of Bases

Bitter

  • taste

  • Turns litmus paper

  • solutions of bases are

    (must be mixed with water!)

  • Reacts with to form and

Slippery

blue

Aqueous

electrolytes

acids

H2O

salt


Naming bases

Naming Bases

polyatomic

  • Use the same rules as for ions (name the cation, then name the anion)

  • NaOH 

  • Ca(OH)2

  • KOH 

Sodium hydroxide

Calcium hydroxide

Potassium hydroxide


Strength of bases

Strength of Bases

Strong

ionize

  • bases – completely in water (create a lot of ions).

  • All hydroxides with groups and metals (except ).

  • bases - only

  • All bases not listed above as strong.

OH-

1

2

Be

ionize

Weak

slightly


Arrhenius theory

Arrhenius Theory

Arrhenius

H

  • An acid must contain a

    and ionize in water to produce

  • An base must contain a

    and dissociates in water to produce

H+ / H3O+

HCl + H2O  H3O+ + Cl-

Arrhenius

OH

OH-

NaOH Na+ + OH-


Disadvantages of arrhenius theory

Disadvantages of Arrhenius Theory

OH-

  • Only compounds with can be classified as a base. What about ammonia, ?

  • Can only be applied to reactions that occur in

  • Would classify some compounds as acids, such as

NH3

water

incorrectly

CH4


Arrhenius acids and bases

Arrhenius Acids and Bases

  • Classify each of the following as an

    Arrhenius acid (A – acid) or base (A – base).

  • Ca(OH)2

  • HBr 

  • H2SO4 

  • LiOH 

A - base

A - acid

A - acid

A - base


Bronsted lowry theory

Bronsted-Lowry Theory

acid

  • A Bronsted – Lowry is any substance that can a

  • A Bronsted – Lowry is any substance that can a

  • HCl + H2O  H3O+ + Cl-

donate

H+

base

accept

H+

B-L base

conjugate base

B-L acid

conjugate acid

Accepts H+

Donates H+


Bronsted lowry theory1

Bronsted-Lowry Theory

  • Let’s look at the reverse reaction.

  • Cl- + H3O+ H2O + HCl

B-L base

conjugate acid

B-L acid

Donates H+

conjugate base

Accepts H+


Bronsted lowry theory2

Bronsted-Lowry Theory

Conjugate

  • acid – formed when a

    accepts a H+ from an acid.

  • base – a that remains after an acid gives up a H+.

  • Conjugate acid – base pair – 2 substances related to each other by the

    of a single H+.

base

Conjugate

particle

accepting/

donating


Types of acids

Types of Acids

  • Defined by how many H+ they can donate.


Bronsted lowry theory3

Bronsted-Lowry Theory

  • Identify the acid, base, conjugate acid, and conjugate base.

conjugate acid

B-L base

B-L acid

HNO3 + H2O  H3O+ + NO3-

Donates H+

conjugate base

Accepts H+


Bronsted lowry theory4

Bronsted-Lowry Theory

  • Give the formula and name of the conjugate base of the following B-L acids.

    (After the B-L acid donates a H+)

  • HI 

  • HCO3-

I- Iodide ion

Since we take away a +, make the ion more –

CO32- carbonate ion


Bronsted lowry theory5

Bronsted-Lowry Theory

  • Give the formula and name of the conjugate acids of the following B-L bases.

    (After the B-L base accepts a H+)

  • H2PO4- 

  • ClO3-

H3PO4 phosphoric acid

Since we add a +,

make the ion more +

HClO3chloric acid


Acidity basicity

Acidity/Basicity

  • Water can sometimes act as a B-L acid and sometimes as a B-L base.

  • The of water:

  • H2O + H2O  H3O+ + OH-

self-ionization

conjugate acid

B-L base

B-L acid

conjugate base


Acidity basicity1

Acidity/Basicity

  • This reaction occurs to a very small extent:

    = = 1 x 10-7 M

    [ ] means

    [H+] x [OH-] =

    relationship

[H+]

[OH-]

concentration

1 x 10-14

Inverse


Acidity basicity2

Acidity/Basicity

pH = 0

pH = 7

pH = 14

[H+]

[OH-]

acid

base

neutral


Acidity basicity3

Acidity/Basicity


Acids

pH

  • [H+] are often small, so the pH scale is easier to use to represent acidity and basicity.

  • pH range is from to

  • log 102 =

  • log 10-3 =

0

14

pH = -log [H+]

2

-3


Acids

pH

neutral

  • In water, a solution,

    ==

    pH =

[H+]

[OH-]

1 x 10-7

pH = negative log [H+]

So… take the exponent

and change the sign!

- (-7) = 7


Acids

pH


Acids

pH

  • If [H3O+] = 1.0 x 10 –5 M, what is the pH?

  • Is the solution basic, neutral, or acidic?

Same as [H+]

pH = - (exponent) = -(-5) = 5

Because < 7


Acids

pH

  • If [H3O+] = 1.0 x 10 –12 M, what is the pH?

  • Is the solution basic, neutral, or acidic?

pH = - (exponent) = -(-12) = 12

Because > 7


Acids

pH

  • Given that a solution has a pH of 2.0, determine the [H3O+].

pH = - (exponent)

2 = - (exponent)

2 = - (-2)

[H3O+] = 1 x 10-2


Acids

pOH

  • Similar to pH, there is also pOH.

  • Because [H+] x [OH-] =

pOH = -log [OH-]

1 x 10-14

pH + pOH = 14


Ph poh

pH/pOH

10-6

10-2

10-4

4

6

12

8

10-12

10-10


Ph poh1

pH/pOH

10-8

10-10

10-14

12

10

2

0

6

10-6

10-4

10-2

1


Ph poh2

pH/pOH

  • If [OH-] = 1.0 x 10 –10 M, what is the pOH?

  • What is the pH?

  • Is the solution basic, neutral, or acidic?

pOH = - (exponent) = -(-10) = 10

pH + pOH = 14

pH + 10 = 14

pH = 4


Ph poh3

pH/pOH

  • What is the pH and the pOH for 1.0 x 10 –6 M HF?

  • pH

  • pOH

pH = - (exponent) = -(-6) = 6

pH + pOH = 14

6 + pOH = 14

pOH = 8


Ph poh4

pH/pOH

  • Given that a solution has a pH of 9.0, determine the [OH -] and the pOH.

  • pOH

  • [OH -]

pH + pOH = 14

9 + pOH = 14

pOH = 5

pOH = - (exponent)

5 = - (exponent)

5 = - (-5)

[OH-] = 1 x 10-5


Review acids conjugate bases

Review: Acids  Conjugate Bases

  • Acids LOSE H+ to become conjugate bases.

  • This is a H atom.

  • When a H+ is lost from an acid, this (-) electron remains.

  • The (+) proton is taken with the H.

-

+

o


Review acids conjugate bases1

Review: Acids  Conjugate Bases

  • What is the conjugate base for the acid HBr?

  • HBr  H+ + Br-

H

Br

Proton is kept

by H.

Electron is left

by H.

+

-

H

+

Br

Conjugate base


Review acids conjugate bases2

Review: Acids  Conjugate Bases

  • What is the conjugate base for the acid HNO2?

  • HNO2 H+ + NO2-

NO2

H

Proton is kept

by H.

Electron is left

by H.

+

-

NO2

H

+

Conjugate base


Review acids conjugate bases3

Review: Acids  Conjugate Bases

  • What is the conjugate base for the acid HSO3-?

  • HSO3- H+ + SO32-

-

SO3

H

Electron is left by H

(added to the one that

was there already).

Proton is kept

by H.

+

2-

SO3

H

+

Conjugate base


Review bases conjugate acids

Review: Bases  Conjugate Acids

  • Bases GAIN H+ to become conjugate bases.

  • What is the conjugate acid for CN-?

  • CN- + H+ HCN

-

+

CN

H

+

The + and the – cancel

out in the final molecule.

CN

H


Review bases conjugate acids1

Review: Bases  Conjugate Acids

  • What is the conjugate acid for NH3?

  • NH3 + H+ NH4+

+

NH3

H

+

There is no – on the NH3 to cancel

the + from H, so the final molecule

is positive.

+

NH3

H


Ph poh cont

pH/pOH cont.

  • Review:

  • pH = [H+] =

  • pOH = [OH-] =

  • pH + pOH =

-log[H+]

10-pH

-log[OH-]

10-pOH

14


Ph poh cont1

pH/pOH cont.

  • Example 1: Determine the pH of a 0.01 M HCl solution.

  • Example 2: Determine the pH of a 0.0010 M NaOH solution.

pH = -log[H+]

pH = -log[0.01]

pH = 2

pOH = -log[OH-]

pOH = -log[0.001]

pOH = 3

pH = 11


Ph poh cont2

pH/pOH cont.

  • Example 3: Determine the pH of a 0.150 M KOH solution.

  • Example 4: Find [H3O+] for a solution that has a pH of 3.0.

pOH = -log[OH-]

pOH = -log[0.15]

pOH = 0.82

pH = 13.18

[H+] = 10-pH

[H+] = 10-3

[H+] = 0.001 M


Ph poh cont3

pH/pOH cont.

  • Example 5: Find [H3O+] for a solution that has a pH of 8.2.

  • Example 6: Find [H3O+] and pOH for a solution that has a pH of 4.85.

[H+] = 10-8.2

[H+] = 10-pH

[H+] = 6.31 x 10-9 M

pOH = 14 - pH

pOH = 9.15

[H+] = 10-4.85

[H+] = 1.41 x 10-5M


Ph poh cont4

pH/pOH cont.

  • Example 7: Find [OH-] for a solution that has a pH of 11.2.

pOH = 14 – 11.2 = 2.8

[OH-] = 10-pOH = 10-2.8

[OH-] = 1.58 x 10-3 M


Neutralization reactions

What happens with you mix an acid with a base? A ____________________________________________ reaction

Neutralization Reactions

  • What happens with you mix an acid with a base? A reaction

  • HCl + NaOH  +

  • Products are always a (

    and ) and

  • This is called a reaction

double replacement

H2O

NaCl

salt

nonmetal

metal

water

neutralization


Neutralization reactions1

Neutralization Reactions

  • Write the balanced chemical equation for the neutralization reaction between:

    nitric acid and potassium hydroxide

+

+

H2O

HNO3

KOH

KNO3

Must use criss-cross to make salt

Must balance equation

This equation is already balanced


Neutralization reactions2

Neutralization Reactions

  • Write the balanced chemical equation for the neutralization reaction between:

    sulfuric acid and magnesium hydroxide

2

+

+

Mg(OH)2

H2O

H2SO4

MgSO4

Must balance equation


Neutralization reactions3

Neutralization Reactions

  • Neutralization reaction:

acid + base  salt + water


Titration

Titration

  • Titration – a process in which an acid-base neutralization reaction is used to determine the of a solution.

concentration

MV = MV

mol mol


Titration1

Titration

  • 1. How many mL of 0.45 M HCl acid must be added to 25.0 mL of 1.00 M KOH to make a neutral solution?


Titration2

Titration

  • 2. What is the molarity of nitric acid if 15.0 mL of the solution is completely neutralized by 38.5 mL of 0.150 M NaOH?


Titration3

Titration

  • 3. A 25.0 mL solution of sulfuric acid is completely neutralized by 18 mL of 1.0 M LiOH. What is the concentration of the H2SO4 solution?


Titration4

Titration


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