AP Chemistry



AP Chemistry Chapter 10 Notes

(Student’s edition)

Chapter 10 problem set: 10, 13, 15, 29, 30, 35, 41, 44-46, 52, 58-60, 63, 64, 67, 68, 73, 74, 81, 87, 89

(optional ( 5, 8, 9, 16, 17, 20-22, 95, 98)

Students should read the introduction to the chapter – it’s important to understand why we learn about the topics that are selected for study in this course.

10-1 Properties of Aqueous Solutions of Acids and Bases

Students should read this section to learn about the properties of acids and bases.

10-2 The Arrhenius Theory

Some historical developments:

1680 – Robert Boyle (Irish) noted acids substances, they affect the color of natural

, and lose their properties when reacted with .

1814 – J. Gay-Lussac (French)noted acids bases and that stated that acids and bases

can only be defined in terms of .

1884 – Svante Arrhenius (Swedish) presents his theory of .

Some new definitions:

Acid – compound that contains and produces in solution

Base – compound that contains and produce in solution

Neutralization – combination of H+1 and OH-1 to form

Arrhenius has ideas that are helpful, but limited in scope as we soon shall see….

3. The Hydrated Hydrogen Ion (Hydronium)

Arrhenius thought of H+1 as a “bare proton.” – . H+1 combines with water to form the

___________________________.

[pic]

This is written and implied in many ways - H+, H+(H2O)n , H+(aq) , H3O+(aq) , etc….

Of course, this involves a

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10-4 The Bronsted-Lowry Theory

In 1923, Bronsted and Lowry independently presented a theory. This new theory was an extension of the original Arrhenius Theory.

Acid – proton

Base – proton

Neutralization – transfer of a proton from an acid to a base

Ionization of HCl, HF, NH3 (in NH3 – ammonia is Base1):

Label: (Base1), (Acid2), (Acid1), (Base2), and show → vs. ↔

Strong Acid Example:

H2O(l) + HCl(g) ( H3O+1(aq) + Cl-1(aq)

(Base1) (Acid2) (Acid1) (Base2)

H2O and are conjugate acid/base pairs. is the conjugate acid of , the base.

HCl and are conjugate acid/base pairs. is the conjugate base of , the acid.

Weak Acid Example:

H2O(l) + HF(g) ( H3O+1(aq) + F-1(aq)

(Base1) (Acid2) (Acid1) (Base2)

H2O and are conjugate acid/base pairs. is the conjugate acid of , the base.

HF and are conjugate acid/base pairs. is the conjugate base of , the acid.

Weak Base Example:

NH3(g) + H2O(l)

(Base1)

NH3 and are conjugate acid/base pairs. is the conjugate acid of , the base.

H2O and are conjugate acid/base pairs. is the conjugate base of , the acid.

Water is . It can act as an acid or base.

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10-5 Autoionization of Water

Tap water conducts electricity – why? – many preset – examples:

Distilled water appears to not conduct electricity, but it does – just a little, tiny bit

H2O + H2O ( H3O+1 + OH-1

but the reaction happens: 0.0000002% in this direction and 99.9999998% in this direction

Draw with Lewis Structures:

10-6 Amphoterism

Amphiprotic (accepts or donates vs. Amphoteric (acts like an or a )

A substance that is amphoteric can act in an amphiprotic way if it is behaving like an acid or a

base by accepting or donating protons.

Al(OH)3 acting as a Base:

Al(OH)3(s) + HCl(aq) →

Al(OH)3 acting as an Acid (accepting H from NaOH…an excess of a strong base):

Al(OH)3(s) + NaOH(aq) → NaAl(OH)4(aq) this is a complex ion

Sn(OH)4(s) + 2 NaOH → Na2Sn(OH)6(aq) see table 10-1 on p376

Most famous amphoteric substance:

Elements of electronegativity form amphoteric hydroxides.

10-10 The Lewis Theory

This is a more encompassing (general) acid/base theory:

Acid – an electron pair (a substance that has room for a pair of electrons)

Base – an electron pair

Neutralization – the formation of a

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(Base1) (Acid2) (Acid1) (Base2)

Example #1: draw lewis structures and geometric shapes for the reactants and products of:

BCl3(g) + NH3(g) → Cl3B:NH3

Acid Base

[pic][pic] [pic] [pic]

[pic] [pic] [pic]

Example #2: draw geometric shapes for the reactants and products of:

SnCl4(1) + 2Cl-1(aq) → SnCl6-2(aq)

Acid Base

[pic] [pic]

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Example #3: draw the lewis structures and geometric shapes for the reactants and products of:

AlCl3(s) + Cl-1(aq) →

Acid Base

11. The Preparation of Acids

Students should go over this section extensively to start to appreciate the richness of chemistry, but….

Should definitely know the following:

a) Direct Combination of Hydrogen and Halogens:

H2(g) + Cl2(g) (

b) Nonvolatile Oxyacids + Salts ( Acid Salts + Weaker Hydrohalic Acids

H2SO4(l) + NaF(s) ( NaHSO4(s) + HF(g)

H2SO4(l) + NaCl(s) (

Note: Acids that are strong oxidizers cannot be used to prepare hydrogen bromide or hydrogen

iodide. Free hydrogen would be produced.

Instead, a non-oxidizing acid should be used:

H3PO4(l) + NaBr(s) ( NaH2PO4(s) + HBr(g)

H3PO4(l) + NaI(s) (

Note: Dissolving each of the gaseous hydrogen halides in water gives the corresponding

hydrohalic acids

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c) Acid Anhydrides (Nonmetal Oxides) + Water ( Oxyacids

Note: some are easy….just add the nonmetal oxide to the water to determine the product.

CO2(g) + H2O(l) ( H2CO3(aq)

SO3(g) + H2O(l) (

Note: some are more complicated…but there are no changes in oxidation state. The Cl below

maintains an oxidation # of

Cl2O7(l) + H2O(l) ( 2 HClO4(aq)

d) Nonmetal Halide + Water ( Oxyacids

Note: The halides and oxyhalides of some nonmetals hydrolyze to produce two acids (a

binary acid and a ternary acid). Also, there are no changes in the oxidation numbers.

PCl3(l) + 3 H2O(l) ( H3PO3(aq) + 3 HCl(aq)

PCl5(s) + 4 H2O(l) ( H3PO4(aq) + 5 HCl(aq)

e) Metal Oxide + Water ( Oxyacids

Note: Some high oxidation state metal oxides(normally thought of as base anhydrides) are actually acid anhydrides that react with water to form oxyacids (interesting point – they can only be made in solution form – no pure form has ever been isolate). Also, there are no changes in the oxidation numbers.

Note: The Mn below maintains an oxidation # of

Mn2O7(l) + H2O(l) ( 2 HMnO4(aq)

Note: The Cr below maintains an oxidation # of

2 CrO3(l) + H2O(l) ( H2Cr2O7(aq)

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NIB - The Preparation of Bases

a) Metal Oxide + Water ( Traditional Base

Note: This is a basic anhydride (a low oxidation state metal oxide)

Na2O(s) + H2O(l) ( 2 NaOH(aq)

MgO(s) + H2O(l) (

b) Metal + Water ( Traditional Base

Na(s) + H2O(l) ( NaOH(aq) + H2(g)

c) Metal Hydride + Water ( Traditional Base

CaH2 (s) + H2O(l) ( Ca(OH)2 (aq) + H2(g)

10-7 Strengths of Acids

Acid strength increases as we move down a family (binary acids)

Acid Strength: HI > HBr > HCl >> HF - why?

Bond Strength: HF>> HCl > HBr > HI – again why?

Electronegativity Difference: HF (1.9)>> HCl (0.9) > HBr (0.7) > HI (0.4)

Bond Length: HI > HBr > HCl >> HF

However, HCl, HBr, and HI are all considered in water, because they all

.

but

In a solvent (water is a solvent – all three ionize to the same degree even

though their strengths are different) there are different degrees of ionization.

A similar trend is found with acids made from members of group 16 (VIA).

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Strengths of Ternary Acids – Let’s compare two hydroxides:

1.0 3.5 2.1

Na------O-------H Thus, the Na+1 and OH-1

HNO3 ( NO2(OH) electronegativities are 3.0, 3.5, and 2.1

Thus, the NO3-1 and H+1

Water (after all we’re mixing these compounds with water) will split the molecules where their

bonds have the .

If two ternary acids have the same central atom, the one with the most oxygens (highest

oxidation # on the central atom) is the .

HClO4 > HClO3 > HClO2 > HClO >

HNO3 > HNO2

but…..

H3PO3 > H3PO4

Isn’t there always an exception in this class? why?

Another exception – H3PO2 > H3PO3 – why?

If two ternary acids have a different central atom in the same oxidation state,

the atom forms the stronger acid (it results in weaker bonds).

HClO4 > HBrO4

HNO3 > H3PO4

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8. Acid-Base Reactions in Aqueous Solutions

This section should be called “acid + base yields salt and water.” It details writing molecular, total, and net ionic equations for acid/base neutralizations.

Some points you need to know:

Strong acids – memorize table 10-2 p. 377

Strong bases – memorize – group 1 and group 2 from Ca↓ (based on solubility rules – know them – table 4-8 p. 135)

Solubility of salts – memorize – based on solubility rules – know them – table 4-8 p. 135

When writing ionic equations, we always write the predominant form of the compound in solution:

HCl is written as H+1(aq) + Cl-1(aq) HNO3 is H+1(aq) + NO3 -1(aq)

H2CO3 is written as itself (weak) Ba(OH)2 is written as Ba+2(aq) + 2 OH-1(aq)

Mg(OH)2 is written as itself (weak) NaCl is written as Na+1(aq) + Cl-1(aq)

AgCl is written as itself (not soluble)

Molecular, Total Ionic, and Net Ionic Equations:

Example #1:

Molecular Equation:

CH3COOH(aq) + NaOH(aq) ( NaCH3COO(aq) + H2O(l)

Total Ionic Equation:

Net Ionic Equation:

Example #2:

Molecular Equation:

H3PO4(aq) + KOH(aq) (

Total Ionic Equation:

Net Ionic Equation:

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9. Acid Salts and Basic Salts

Depending on stoichiometrical relationships, different types of salts can be formed. Now we look at

acid/base reactions in terms of limiting and excess reactants.

Examples:

Note: The production of Normal Salt. This is a salt that contains no ionizable H or OH.

H3PO4(aq) + NaOH(aq) (

Note: If less than the appropriate stoichiometrical amounts are added, then the resulting salts

are known as acidic salts…

H3PO4(aq) + NaOH(aq) ( NaH2PO4(aq) + H2O(l)

H3PO4(aq) + 2 NaOH(aq) ( Na2HPO4(aq) + 2 H2O(l)

Note: The acidic salts above can also react with bases…

NaH2PO4(aq) + 2 NaOH(aq) ( Na3PO4(aq) + 2 H2O(l)

Na2HPO4(aq) + NaOH(aq) ( Na3PO4(aq) + H2O(l)

Note: Polyhydroxl bases react with acids to produce Normal Salts.

Al(OH)3(s) + HCl(aq) (

Note: If less than the appropriate stoichiometrical amounts are added, then the resulting salts

are known as basic salts…

Al(OH)3(s) + HCl(aq) ( Al(OH)2Cl(s) + H2O(l)

Al(OH)3(s) + 2 HCl(aq) ( Al(OH)Cl2(s) + 2 H2O(l)

Think about how would you make Al(OH)Br2, K2HPO4, etc.

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-----------------------

Coordinate

Covalent Bond

Bond formed

Coordinate

Covalent Bond

Bond formed

Cl

Coordinate

Covalent Bond

+

’!

Cl

Cl

B

N

Bond formed

+

’!

H

H

H

Cl

Cl

Cl

N

B

Each Cl donates an electron to create two coordinate covalent bonds

Coordinate

Covalent Bonds

’!

2Cl-1

+→

Cl

Cl

B

N

Bond formed

+

→

H

H

H

Cl

Cl

Cl

N

B

Each Cl donates an electron to create two coordinate covalent bonds

Coordinate

Covalent Bonds

→

2Cl-1

+

Cl

Cl

Cl

Cl

Cl

Cl

Sn

Sn

Cl

Cl

Cl

Cl

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