CHM135H1 Chapter Notes - Chapter 15: Conjugate Acid, Chemical Species, Hydronium

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CHM135 CHEMISTRY TEXTBOOK CH. 15 (Excluding 15.5)
NOTES/SUMMARY
Chapter 15 Aqueous Equilibria: Acids and Bases
Chapter 15.1: Acid Base Concepts: The Bronsted Lowry Theory
- So far, Arrhenius theory of acids and bases has been used according to Arrhenius, acids are
substances that dissociate in water to produce hydrogen ions, and bases dissociate in water to
yield hydroxide ions
- Generalized Arrhenius Acid: HA(aq) H+(aq) + A-(aq)
Generalized Arrhenius Base: MOH(aq) M+(aq) + OH-(aq)
- Arrhenius theory accounts for properties of many common acids and bases; has important
limitations Arrhenius theory restricted to aqueous solutions, doesn’t account for basicity of
substances
- Bronsted Lowry theory is more general theory of acids and bases where an acid is any
substance (molecule or ion that can transfer a proton (H+ ion) to another substance, and a base is
any substance that can accept a proton
- Acids are proton donors, bases are proton acceptors, and acid base reactions are proton
transfer reactions
- Bronsted Lowry Acid: a substance that can transfer H+
Bronsted Lowry Base: a substance that can accept H+
- (Textbook, Page 604)
- The products of a Bronsted Lowry acid base reaction, BH+ and A-, are themselves,
respectively, acids and bases; he species BH+ produced when base B accepts a proton from HA
can itself donate a proton back to A- ; similarly, when HA loses a proton, A- can accept a proton
back from BH+
- Chemical species whose formulas differ only by one proton are said to be conjugate acid base
pairs A- s the conjugate base of the acid HA, and HA is the conjugate acid of the base A-; B is
conjugate base of acid BH+, and BH+ is conjugate acid of base B
- (Textbook, Page 605)
- E.g. when a Bronsted Lowry acid HA dissolves in water, it reacts reversibly with water in an
acid dissociation equilibrium acid transfers proton to solvent, which acts as a base, and
products are hydronium (conjugate acid of water), and A- (conjugate base of HA)
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o In reverse reaction, hydronium ion acts as proton donor and A- acts as proton acceptor
- For a molecule or ion to accept a proton, it must have at least on unshared pair of electrons that it
can use for bonding to the proton
- (Textbook, Page 605)
- Proton is fundamental to both Arrhenius and Bronsted Lowry definitions of an acid in
Arrhenius definition, acid (HA) dissociates to give aqueous hydrogen ion, or hydrated proton
(H+(aq))
- As bare proton, positively charged H+ ion is too reactive to exist in aqueous solutions; bonds to
oxygen atom of a solvent water molecule to give trigonal pyramidal hydronium ion
- Hydronium is simplest hydrate of the proton; it can associate through hydrogen bonding with
additional water molecules to give higher hydrates with general formula:
   
- It’s likely that acidic aqueous solutions contain a distribution of  ions having
different values of n
- A proton hydrated by an unspecified number of water molecules
Chapter 15.2: Acid Strength and Base Strength
- In acid dissociation equilibrium, two bases, H2O and A- are competing for protons:
HA(aq) + H2O(l) H3O+(aq) + A-(aq)
- If H2O is stronger proton acceptor than A- , H2O molecules will get protons and solution will
contain mainly products; if A- is stronger base than H2O, the A- ions will get the protons and
solution will contain mainly reactants
- When beginning with equal concentrations of reactants and products, the proton is always
transferred to the stronger base direction of reaction to reach equilibrium is proton transfer
from stronger acid to the stronger base to give weaker acid and weaker base:
Stronger acid + Stronger base Weaker acid + Weaker base
- Different acids differ in their ability to donate protons; strong acid is almost completely
dissociated in water and is thus, a strong electrolyte acid dissociation equilibrium of a strong
electrolyte; acid dissociation equilibrium of a strong acid lies nearly 100% to right , and
solution contains almost entirely of products
- Strong acids have very weak conjugate bases
- Weak acid is only partially dissociated in water and is thus a weak electrolyte; only a small
fraction of the weak acid molecules transfer a proton to water, and solution contains mainly
undissociated HA molecules along with small amounts products
- Very weak acids have strong conjugate bases
- (Textbook, Page 609)
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- (Textbook, Page 610)
Chapter 15.3: Factors That Affect Acid Strength
- Strength of the H A bond is given by the bond dissociation energy (D), which is amount of
energy required to dissociate HA into an H atom and an A atom polarity of the H A bond
increases with an increase in electronegativity of A and is related to ease of electron transfer from
an H atom to an A atom to give an H+ cation and an A- anion
- The weaker and more polar the H A bond, the stronger the acid
- Electrostatic potential maps show range of polarity in hydrohalic acids HF, HCl, HBr and HI
(Textbook, Page 611)
- Variation in polarity in this series is much less important than variation in bond strength
- In general, for binary acids of elements in the same group of the periodic table, the H A bond
strength is the most important determinant of acidity
- H A bond strength generally decreases with increasing size of element A down a group, so
acidity increases
(Textbook, Page 611)
- For binary acids of elements in the same row of the periodic table, changes in the H A bond
strength are smaller and the polarity of the H A bond is the most important determinant of acid
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