Download aq - Harding Charter Preparatory High School

Chapter 14
Acids and Bases
Properties of Acids and Bases
• Acids
– Give foods a tart or sour taste
• What acidic foods might you eat?
– Aqueous solutions of acids are electrolytes
• Conduct Electricity
– Some are strong electrolytes (strong acids)
– Some are weak electrolytes (weak acids)
• Cause indicator dyes to change colors
• Many metals react with acids producing hydrogen gas
• React with compounds containing hydroxide ions to form
water and a salt
Properties of Acids and Bases
• Bases
– Have bitter taste, and slippery feel
– Aqueous solutions of bases are also
electrolytes
• Conduct Electricity
– Some are strong electrolytes (strong bases)
– Some are weak electrolytes (weak bases)
• Cause indicator dyes to change colors
• Water and salt are formed when a base that
contains hydroxide ions react with an acid
Arrhenious Acids and Bases
• Acids are hydrogen-containing compounds
that ionize to yield hydrogen ions (H+) in
aqueous solution
• Bases are compounds that ionize to yield
hydroxide ions (OH-) in aqueous solution
Arrhenious Acids
• Can be monoprotic, diprotic, or triprotic
– Monoportic: HNO3 → H+ + NO3• Ionization yields one hydrogen ion
– Diprotic: H2SO4 → 2H+ + SO42• Complete ionization yields 2 hydrogen ions
– Triprotic: H3PO4 → 3H+ + PO43• Complete ionization yields 3 hydrogen ions
• Not all the hydrogens in an acid may be released as hydrogen ions
• Not all hydrogen containing compounds are acids
– Only hydrogens joined to very electronegative elements, and thus have
very polar bonds, are ionizable in water
H O
H C C
O-
H
Ethanoic Acid
H+
Nonionizable
Hydrogen
Ionizable
Hydrogen
Arrhenious Bases
• NaOH → Na+(aq) + OH-(aq)
• KOH → K+(aq) + OH-(aq)
– Bases formed with group one metals are very soluble
and caustic
– Can be made by reacting group one metals with
water
• Na + H2O → Na+(aq) + OH-(aq) H2 (g)
• Bases of group 2 metals are very weak resulting
low solubility
– Examples are Ca(OH)2 and Mg(OH)2
Bronsted-Lowry Acids and Bases
• Arrhenious definition of acids and bases is not very
comprehensive and does not explain why certain
substances have basic or acidic properties
– Ammonia (NH3) is a base, but there is no hydroxide (OH-) in the
compound to ionize
• The Bronsted-Lowry theory defines an acid as a
hydrogen-ion donor, and a base as a hydrogen-ion
acceptor
– Why ammonia is a base
NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq)
Hydrogen ion
aceptor, BronstedLowry Base
Hydrogen ion
donar, BronstedLowry Acid
Makes the
solution basic
Conjugate Acids and Bases
conjugate acid-base pair
NH4+(aq) + OH-(aq)
NH3(aq) + H2O(l)
Base
Conjugate
Acid
Acid
Conjugate
Base
conjugate acid-base pair
•
•
•
A conjugant acid is the particle formed when a base gains a hydrogen ion
A conjugant base is the particle that remains when an acid has donated a
hydrogen ion
A conjugate acid-base pair consists of two substances related by the loss or
gain of a single hydrogen ion
conjugate acid-base pair
HCl(aq) + H2O(l)
Acid
Base
H3O+(aq) + Cl-(aq)
Conjugate
Acid
conjugate acid-base pair
Conjugate
Base
Conjugate Acids and Bases
conjugate acid-base pair
HCl(aq) + H2O(l)
Acid
Base
H3O+(aq) + Cl-(aq)
Conjugate
Acid
Conjugate
Base
conjugate acid-base pair
• A water molecule that gains a hydrogen ion becomes a positively
charged hydronium ion H3O+
– In above equation, what is the hydrogen ion donor (acid) and which is
the hydrogen ion acceptor (base)
• Notice, water can both accept and donate a hydrogen ion and thus
act as an acid and a base
– A substance that can act as both an acid and a base is said to be
amphoteric
• Amino Acids as an example
Lewis Acids and Bases
• Acids accept a pair of electrons during a
reaction while a base donates a pair of
electrons
– Lewis acid – a substance that can accept a
pair of electrons to form a covalent bond
– Lewis base – a substance that can donate a
pair of electrons to form a covalent bond
NH3 + BF3 → NH3BF3
Identify the Lewis Acid and the Lewis Base in the above equation
Acid-Base Definitions Review
Type
Acid
Base
Arrhenius
H+ producer
OH- producer
Bronsted-Lowry
H+ donor
H+ acceptor
Lewis
electron-pair acceptor
electron-pair donor
Hydrogen Ions and Acidity
• Occasionally collusions between water molecules
cause them to react forming hydroxide ions and
hydronium ions
– The reaction in which water molecules produce ions is called
the self ionization of water
H2O(l)
H2O (l) + H2O (l)
H+(aq) + OH-(aq)
H3O+(aq) + OH-(aq)
• In aqueous solution, hydrogen ions H+ are always joined
to a water molecule as hydronium ions
• In pure (neutral) water, the self-ionization of water results
in 1 x 10-7 M of H+ ions and 1 x 10-7 M of OH- ions
– Any aqueous solution in which H+ and OH- ions are equal is
described as a neutral solution
Ion Product Constant for Water
•
For aqueous solutions, the product of the hydrogen-ion concentration and
the hydroxide-ion concentration equals 1.0x10-14
Kw = [H+][OH-] = 1.0x10-14 M
•
•
The product of the concentrations of the hydrogen ions and hydroxide ions
in water is called the ion-product constant
Aqueous acids and bases sift the ratio of hydrogen ions to hydroxide ions in
solution causing it to become either acidic or basic
– In a basic solution aka alkaline solution, the hydroxide ion (OH-) is greater than
1x10-7 M and the hydrogen ion (H+) is less 1x10-7 M
– In a acidic solution, the hydrogen ion (H+) is greater than 1x10-7 M and the
hydroxide ion (OH-) is less 1x10-7 M
•
Regardless of the acidity or alkalinity of the solution, the product of the
Molarity (M) concentration of H+ and OH- always equals 1x10-14 at 25ºC
– If the hydrogen ion H+ concentration in a soft drink is 1 x 10-5 M, what is the
concentration of the hydroxide ion OH-?
– Is the solution basic, neutral, or acidic?
The pH Concept
• Expressing hydrogen-ion concentrations in molarity is cumbersome
– Soren Sorensen suggested that the hydrogen ion concentration be
expressed as the negative log of the hydrogen-ion concentration giving
us much smaller numbers to work with
pH = -log[H+]
• The pH of a solution is the negative logarithm of the hydrogen ion
concentration
• A neutral solution H+ = 1x10-7 has a pH = -log[1x10-7]= 7
• A solution in which the [H+] is greater than 1x10-7 M and has a pH less than
7.0 is acidic
• The pH of pure water or a neutral solution has a pH of 7
• A solution with a pH greater than 7 is basic and has a [H+] concentration of
less than 1x10-7 M
– You can also calculate pOH which is the negative logarithm of
hydroxide ion concentration
pOH = -log[OH-]
Work some pH problems (pH of 1 x10-5 M H+?)
Relationship between pH and pOH
pH + pOH = 14
pH = 14 – pOH
pOH = 14 - pH
pH and significant figures
• Hydrogen ion concentrations should
always be reported to two significant
figures
• pH and pOH calculations should always
be reported to two decimal places
– Rules are due to the sensitivity of pH meters
Acid-Base Indicators
Dyes
• An indicator (HIn) is an acid or base that undergoes
dissociation in a known pH range
– An indicator is a valuable tool for measuring pH because its
acid form and base form have different colors in solution
HIn (aq)
Acid Form
H+
OH-
H+ (aq) + In- (aq)
Base Form
– The acid form of the indicator dominates the disassociation
equilibrium at low pH
– The basic form of the indicator dominates the disassociation
equilibrium at high pH
– Color change of an indicator occurs in a narrow pH range ≈ 2
pH units
– Thus it takes many indicators to span the entire pH spectrum
• Indicator dyes have limitations
Acid-Base Indicators
pH Meter
• Makes rapid, accurate pH
measurements
• Can record pH continuously over time
when performing a reactions
• Measures pH to two decimal places
• Color and cloudiness of solution does
not interfere with reading
• Are many different types specialized
for different jobs were pH
measurements are required
http://www.vittbi.com/photogallery/biotech/PH-Meter.jpg
http://personals.galaxyinternet.net/tunga/Meter.jpg
Strengths of Acids and Bases
Strong and Weak Acids and Bases
• Acids are classified as strong or weak
depending on the degree to which they ionize in
water
– A strong acid have an equilibrium that lies far to the
right (almost all the original HA ionizes)
– Weak acids have an equilibrium that lies far the left
(ionize only slightly in aqueous solution)
– What are some example of strong acids and weak
acids
HCl(q) + H2O(l)
CH3COOH(aq) + H2O(l)
H3O+(aq) + Cl-(aq) 100% ionized
H3O+(aq) + CH3COO-(aq) <1% ionized
pH of 0.10 M Solutions of Common Acids and Bases
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Compound
HCl (hydrochloric acid)
H2SO4 (sulfuric acid)
NaHSO4 (sodium hydrogen sulfate)
H2SO3 (sulfurous acid)
H3PO4 (phosphoric acid)
HF (hydrofluoric acid)
CH3CO2H (acetic acid)
H2CO3 (carbonic acid)
H2S (hydrogen sulfide)
NaH2PO4 (sodium dihydrogen phosphate)
NH4Cl (ammonium chloride)
HCN (hydrocyanic acid)
Na2SO4 (sodium sulfate)
NaCl (sodium chloride)
NaCH3CO2 (sodium acetate)
NaHCO3 (sodium bicarbonate)
Na2HPO4 (sodium hydrogen phosphate)
Na2SO3 (sodium sulfite)
NaCN (sodium cyanide)
NH3 (aqueous ammonia)
Na2CO3 (sodium carbonate)
Na3PO4 (sodium phosphate)
NaOH (sodium hydroxide, lye)
pH
1.1
1.2
1.4
1.5
1.5
2.1
2.9
3.8 (saturated solution)
4.1
4.4
4.6
5.1
6.1
6.4
8.4
8.4
9.3
9.8
11.0
11.1
11.6
12.0
13.0
http://www.cartage.org.lb/en/themes/sciences/chemistry/Inorganicchemistry/AcidsBases/Common/Common.htm
Acid Disassociation Constant
• The equilibrium constant for weak acids (HA) can be written as:
Acid
Conjugate base
H3O+(aq) + A-(aq)
HA(aq) + H2O(l)
Keq =
[H3O+] [A-]
[HA] [H2O]
• For dilute solutions, the concentration of water is a constant, and
can be combined with Keq to give the acid dissociation constant
(Ka)
Keq X H2O = Ka=
[H3O+] [A-]
[HA]
• Ka reflects the fraction of an acid in the ionized form and thus is
sometimes referred to as the ionization constant
• Weak acids have small Ka values, while stronger acids have
larger Ka values; why?
Base Disassociation Constant
•
The equilibrium constant for weak Bases (B) can be written as:
base
Conjugate acid
BH+(aq) + HO-(aq)
B(aq) + H2O(l)
Keq =
[BH+] [HO-]
[B] [H2O]
• For dilute solutions, the concentration of water is a constant, and
can be combined with Keq to give the base dissociation constant
(Kb)
Keq X H2O = Kb=
[BH+] [OH-]
[B]
• Kb is the ratio of the concentration of the conjugate acid times the
concentration of the hydroxide ion to the concentration of the
base
• The magnitue of Kb indicates the ability of a weak base to
compete with the very strong base OH- for hydrogen ions
• The smaller the Kb the weaker the base
Relationship Between Ka and Kb
𝑯+ [𝑨− ] 𝑯𝑨 [𝑶𝑯− ]
+
−
𝑲𝒂 × 𝑲𝒃 =
×
=
𝑯
𝑶𝑯
= 𝑲𝒘 = 𝟏. 𝟎 × 𝟏𝟎−𝟏𝟒
−
[𝑯𝑨]
[𝑨 ]
• Allow you to calculate Kb from Ka or Ka from Kb
Concentration and Strength
• Remember, the word strong and weak acids and bases
refers to the number particles of the acid or base that
completely dissociate into their respective ions in
solution
• Concentration and dilute refer to how many moles of an
acid or base is diluted in a constant volume of solution
• Even though an acid may be “weak”, if it is highly
concentrated, it will result in much lower pH of the
solution it is dissolved in than a dilute solution of the
same weak acid
Various Ways to Describe Acid Strength
Property
𝐾𝑎 Value
Position of the dissociation
(ionization) equilibrium
Strength of conjugate
base compared with that
of water
Strong Acid
Weak Acid
𝐾𝑎 is large
𝐾𝑎 is small
𝐻+ ≈ [𝐻𝐴]0
𝐻 + ≪ [𝐻𝐴]0
𝐴− much weaker base
than 𝐻2 𝑂
𝐴− much stronger base
than 𝐻2 𝑂
Practice Problems
𝑲𝒂 Values
HF
7.2 × 10−4
𝐻𝑁𝑂2
4.0 × 10−4
HCN
6.2 × 10−10
• Relative base strength
– Using the table above and your knowledge of
strong acids, arrange the following species
according to their strength as bases:
– 𝐻2 𝑂, 𝐹 − , 𝐶𝑙 − , 𝑁𝑂2− , 𝐶𝑁 −
Practice Problems
• Calculating [𝐻+ ] or 𝑂𝐻−
– Calculate [𝐻 + ] or 𝑂𝐻 − as required for each
of the following solutions at 25°C, and state
whether the solution is neutral, acidic, or basic
a) 1.0 × 10−5 𝑀 𝑂𝐻 −
b) 1.0 × 10−7 𝑀 𝑂𝐻 −
c) 10.0 M 𝑀 𝐻 +
Practice Problems
• Calculating pH and pOH
– Calculate the pH and pOH for each of the
following solutions at 25°C.
a) 1.0 × 10−3 𝑀 𝑂𝐻 −
b) 1.0 𝑀 𝐻 +
Practice Problems
• Calculating pH
– The pH of a sample of human blood was
measured to be 7.41 at 25°C. Calculate pOH,
[H+], and [OH-] for the sample.
Practice Problems
• pH of Strong Acids
– Calculate the pH of 0.10 M HNO3.
– Calculate the pH of 1.0 × 10−10 M HCl.
Calculating the pH of Weak Acid Solutions
1. List the major species in the solution
2. Choose the species that can produce H+, and write balanced
equations for the reactions producing H+
3. Using the values of the equilibrium constants for the reaction you
have written, decide which equilibrium will dominate in producing H+
4. Write the equilibrium expression for the dominant equilibrium
5. List the initial concentrations of the species participating in the
dominant equilibrium
6. Define the change needed to achieve equilibrium; that is define x
7. Write the equilibrium concentration in terms of x
8. Substitute the equilibrium concentrations into the equilibrium
expression
9. Solve for x the “easy” way, that is, by assuming that [𝐻𝐴]0 − 𝑥 ≈
[𝐻𝐴]0
10.Use the 5% rule to verify whether the approximation is valid
–
𝑥
[𝐻𝐴]
× 100 ≤ 5%
11.Calculate [H+] and pH
Practice Problem
• Calculating the pH of a weak acid
1. Calculate the pH of 0.100 M solution of
hypochlorous acid. (𝐾𝑎 = 3.5 × 10−8 𝑀)
Practice Problems
• The pH of week acid mixtures
1. Calculate the pH of a solution that contains 1.00 M
HCN (𝐾𝑎 = 6.2 × 10−10 ) and 5.00 M HNO2 (𝐾𝑎 =
4.0 × 10−4 ). Also calculate the concentration of the
cyanide ion in this solution at equilibrium.
Calculating Dissociation Constants
• Disassociation constants are calculated from
experimental data
• To find the Ka of weak acid or the Kb of a weak
base, substitute the measured concentrations of
all the substances present at equilibrium into the
expression for Ka or Kb
• A 0.1000M solution of ethanoic acid is only
partially ionized and has a pH of 2.87. What is
the acid dissociation constant (Ka) of ethanoic
acid?
Percent Dissociation
• The percent dissociation of a weak acid is
𝑚𝑜𝑙
𝑎𝑚𝑜𝑢𝑛𝑡 𝑑𝑖𝑠𝑠𝑜𝑐𝑖𝑎𝑡𝑒𝑑 (
)
𝐿
𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝐷𝑖𝑠𝑠𝑜𝑐𝑖𝑎𝑡𝑖𝑜𝑛 =
× 100%
𝑚𝑜𝑙
𝑖𝑛𝑖𝑡𝑖𝑎𝑙 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 (
)
𝐿
1. Calculate the percent dissociation of acetic acid (𝐾𝑎 = 1.8 × 10−5 ) in each of the
following solutions.
a)
b)
–
1.00 M 𝐻𝐶2 𝐻3 𝑂2
0.100 M 𝐻𝐶2 𝐻3 𝑂2
For solutions of any weak acid HA, [𝐻 + ] decreases as [𝐻𝐴]0 decreases, but the
percent dissociation increases as [𝐻𝐴]0 decreases
Practice Problems
• Calculating Ka from percent dissociation
1. Lactic acid (HC3H5O3
2. ) is a waste product that accumulates in muscle
tissue during exertion, leading to pain and feeling of
fatigue. In a 0.100 M aqueous solution, lactic acid is
3.7% dissociated. Calculate the value of Ka for this
acid.
Practice Problem
• The pH of strong bases
1. Calculate the pH of 5.0 × 10−2 𝑀 NaOH
solution.
Practice Problem
• pH of a weak base
1. Calculate the pH for a 15.0 M solution of
NH3 (𝐾𝑏 = 1.8 × 10−5 )
2. Calculate the pH of a 1.0 M solution of
methylamine 𝐶𝐻3 𝑁𝐻2 𝐾𝑏 = 4.38 × 10−4
Polyprotic Acids
• Polyprotic Acids are Acids that can furnish more than
one proton
• Example is the triprodic acid phosphoric acid
𝐻3 𝑃𝑂4 𝑎𝑞 ↔ 𝐻 + 𝑎𝑞 + 𝐻2 𝑃𝑂4− 𝐾𝑎1 = 7.5 × 10−3
𝐻2 𝑃𝑂4− 𝑎𝑞 ↔ 𝐻 + 𝑎𝑞 + 𝐻𝑃𝑂42− 𝐾𝑎2 = 6.2 × 10−8
𝐻𝑃𝑂42− 𝑎𝑞 ↔ 𝐻 + 𝑎𝑞 + 𝐻2 𝑃𝑂43− 𝐾𝑎3 = 4.8 × 10−13
– For a typical weak polyprotic acid 𝐾𝑎1 ≫ 𝐾𝑎2 ≫ 𝐾𝑎3
• Means that only the first dissociation step makes an important contribution to
[H+]
Practice Problem
• Calculate the pH of a 5.0 M H3PO4 solution and
the equilibrium concentrations of the species
𝐻3 𝑃𝑂4 , 𝐻2 𝑃𝑂4− , 𝐻𝑃𝑂42− , 𝑎𝑛𝑑 𝑃𝑂43−
Sulfuric Acid
• Sulfuric Acid is unique in the common acids
since it is a strong acid in its first dissociation
step and a weak acid in its second step:
𝐻2 𝑆𝑂4 𝑎𝑞 → 𝐻 + + 𝐻𝑆𝑂4− 𝑎𝑞 𝐾𝑎 𝑖𝑠 𝑣𝑒𝑟𝑦 𝑙𝑎𝑟𝑔𝑒
𝐻𝑆𝑂4− 𝑎𝑞 → 𝐻 + + 𝑆𝑂42− 𝑎𝑞 𝐾𝑎2 = 1.2 × 10−2
1. Calculate the pH of a 1.0 M sulfuric acid solution
2. Calculate the pH of a 0.0100 M sulfuric acid solution
Acid-Base Properties of Salts
• A salt is an ionic compound
– In dilute solutions, it is assumed that salts break up into their respective
ions
• Sometimes, the ions can behave as acids or bases
Acid-Base Properties of Various Types of Salts
Type of Salt
Examples
Comment
pH of
Solution
Cation is from strong base; anion is
from strong acid
NaCl, KNO3,
NaCl, NaNO3
Neither acts as an acid or a
base
Netural
Cation is from strong base; anion is
from weak acid
NaC2H3O2,
KCN, NaF
Anion acts as a base; cation has
no effect on pH
Basic
Cation is conjugate acid of weak base;
anion is from strong acid
NH4Cl,
NH4NO3
Cation acts as acid; anion has
no effect on pH
Acidic
Cation is conjugate acid of weak base;
anion is conjugate base of weak acid
NH4C2H3O2,
NH4CN
Cation acts as an acid; anion
acts as a base
Acidic if 𝐾𝑎 > 𝐾𝑏 ,
basic if 𝐾𝑎 < 𝐾𝑏 ,
neutral if 𝐾𝑎 = 𝐾𝑏
Cation is highly charged metal ion;
anion is from strong acid
Al(NO3)3,
FeCl3
Hydrated cation acts as an acid;
anion has no effect on pH
Acidic
Practice Problem
• Salts as a weak base
– Calculate the pH of a 0.30 M NaF solution.
The Ka value for HF is 7.2 × 10−4 .
Practice Problem
• Salts as weak acids
– Calculate the pH of a 0.10 M NH4Cl solution.
The Kb value for NH3 is 1.8 × 10−5 .
– Calculate the pH of a 0.010 M AlCl3 solution.
−5
The Ka value for 𝐴𝑙(𝐻2 𝑂)3+
is
1.4
×
10
6
2+
+ (𝑎𝑞)
𝐴𝑙(𝐻2 𝑂)3+
(𝑎𝑞)
↔
𝐴𝑙(𝑂𝐻)(𝐻
𝑂)
𝑎𝑞
+
𝐻
2
6
5
Practice Problem
• The Acid-Base Properties of Salts
– Predict whether an aqueous solution of each of the
following salts will be acidic, basic, or neutral
a) 𝑁𝐻4 𝐶2 𝐻3 𝑂2
b) 𝑁𝐻4 𝐶𝑁
c) 𝐴𝑙2 (𝑆𝑂4 )3
•
•
•
•
𝑁𝐻4+ , 𝐾𝑎 = 5.6 × 10−10
𝐶2 𝐻3 𝑂2− , 𝐾𝑎 = 5.6 × 10−10
𝐶𝑁 − , 𝐾𝑏 = 1.6 × 10−5
𝑆𝑂42− , 𝐾𝑏 = 8.3 × 10−13
The affection of structure on Acid-Base Properties
• The two main factors that determine whether a molecule containing
a X—H bond will behave as a Bronsted-Lowry acid are the bond
strength and polarity
– C—H bonds are do not produce acidic solutions since a C—H bond is
strong an nonpolar
– The H—Cl bond, however, is stronger than the C—H bond, but it also
much more polar and readily dissociates in water
Bond Strengths and Acid Strengths for Hydrogen Halides
H—X Bond
Bond Strength
(kJ/mol)
Acid Strength in
water
H—F
565
Weak
H—Cl
427
Strong
H—Br
363
Strong
H—I
295
Strong
Most Polar
H—F
Bond Polarity
>
H—Cl
>
Least Polar
HBr
>
H--I
Oxyacids
Oxyacid
𝐻𝐶𝑙𝑂4
𝐻𝐶𝑙𝑂3
Ka Value
Structure
Large
O
H O Cl O
O
~1
H O Cl
𝐻𝐶𝑙𝑂2
1.2 × 10−2
H O Cl O
𝐻𝐶𝑙𝑂
3.5 × 10−8
H O Cl
• Oxyacids
characteristically contain
the grouping (H—O—X)
– Acid strengths increases
with increasing oxygen
• Due to very
electronegative oxygen
atoms are able to draw
electrons away form the
chlorine atom and O—H
bond making it weaker
and more polar
Oxyacids
Comparison of electronegativity of X and Ka Values for a series of oxyacids
Acid
X
Electronegativity of X
Ka for Acid
HOCl
Cl
3.0
4 × 10−8
HOBr
Br
2.8
2 × 10−9
HOI
I
2.5
2 × 10−11
HOCH3
CH3
2.3 (for Carbon in CH3)
~10 × 10−15
• The more electronegative X is, the stronger the acid is
Acids of Highly Charged Metal Ions
• The acidity of water molecules attached to the
metal ion is increased by the attraction of
electrons to the positive metal ion:
H
Al3+ O
H
Acid Base Properties of Oxides
• A compound containing the H—O—X group will produce
an acidic solution in water the O—X bond is strong and
covalent
𝑆𝑂3 (𝑔) + 𝐻2 𝑂(𝑙) → 𝐻2 𝑆𝑂4 𝑎𝑞
𝐶𝑂2 𝑔 + 𝐻2 𝑂(𝑙) → 𝐻2 𝐶𝑂3 (𝑎𝑞)
2 𝑁𝑂2 𝑔 + 𝐻2 𝑂(𝑙) → 2𝐻𝑁𝑂3 (𝑎𝑞)
• If the O—X bond is ionic, the compound will produce a
basic solution in water
𝐶𝑎𝑂 𝑠 + 𝐻2 𝑂(𝑙) → 𝐶𝑎(𝑂𝐻)2 (𝑎𝑞)
𝐾2 𝑂(𝑠) + 𝐻2 𝑂(𝑙) → 2𝐾𝑂𝐻(𝑎𝑞)