1
ACIDS, BASES AND SALTS
Water (H2O) is made of two ions
H+ (aq) – hydrogen ion
OH- (aq) – hydroxide ion
Acids – substances that increase H+ (aq) conc.
- definition of Arrhenius acid
Strong acids
- strong electrolytes, i. e., completely dissociate
HCl (aq) = H+ (aq) + Cl- (aq)
- memorize list
HCl, HBr, HI, HNO3, H2SO4, HClO3, HClO4
Weak acids
- weak electrolytes
HF (aq)  H+ (aq) + F- (aq)
HF (aq)  HF (aq) + H+ (aq) + F- (aq)
a lot
a little a little
+
H3PO4 (aq)  H (aq) + H2PO4- (aq)
Bases – substances that increase OH- concentration
- definition of Arrhenius base
Strong bases
- strong electrolytes
- all alkali metal hydroxides
LiOH, NaOH, KOH, RbOH, CsOH
- some alkaline earth metal hydroxides
Ca(OH)2, Sr(OH)2, Ba(OH)2
Weak bases
- weak electrolytes
- usually increase OH- (aq) conc. “indirectly” by decreasing H+ (aq) conc.
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
C2H3O2- (aq) + H2O (l)  HC2H3O2 (aq) + OH- (aq)
Neutralization reaction
Acid + Base  Water + Salt
Salt
- ionic compound
- cations and anions remaining after water is made
HNO3 (aq) + KOH (aq)  H2O (l) + KNO3 (aq)
base
acid
water
salt
BRØNSTED – LOWRY ACID/BASE THEORY
2
Acid/base reactions are simply transfer of proton (hydronium ion) from acid to base.
Brønsted – Lowry Acid
- proton donor
- note HCl, H2SO4, HC2H3O2 all donate protons
Brønsted – Lowry Base
- proton acceptor
- note OH- and NH3 accept protons
In Brønsted – Lowry theory, water is an acid or base depending on the circumstances.
HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq)
- water accepts proton; therefore, it is a base.
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
- water donates proton; therefore, it is a acid.
Conjugate Acids and Bases
Consider a typical acid/base reaction.
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
- H2O has donated proton, it is an acid.
- NH3 accepted proton, it is a base.
Consider the reverse reaction.
NH4+ (aq) + OH- (aq)  NH3 (aq) + H2O (l)
- NH4+ has donated proton, it is an acid.
- OH- accepted proton, it is a base.
- An acid is changed into a base and a base is changed into an acid.
- These pairs of acids and bases are called conjugate acid-base pairs.
- A conjugate acid-base pair differ from each other only by a proton.
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
Another Example:
H2SO4 (aq) + H2O (l)  HSO4- (aq) + H3O+ (aq)
Another Example:
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HC2H3O2 + NaOH  H2O + NaC2H3O2
acid
base
conjugate conjugate
acid
base
Relative strengths of acids and bases
The strength of a conjugate base is related to the strength of its acid.
The strength of a conjugate acid is related to the strength of its base.
Consider perchloric acid:
HClO4 (aq) – acid
ClO4- (aq) – conjugate base
HClO4 is a very strong acid, meaning that it fully dissociates.
HClO4 (aq)  H+ (aq) + ClO4- (aq)
- This means that the ClO4- ion has no desire to accept a proton.
- Therefore the perchlorate ion is a very weak (negligible) base.
Consider aqueous sodium hydroxide:
OH- – base
H2O – acid
NaOH (aq)  Na+ (aq) + OH- (aq)
- OH- ion reacts fully (almost) with H+ ions.
- If any additional H+ is added, OH- will grab it to form water.
- OH- is a very strong base.
- The conjugate acid, H2O is a very weak acid.
As the strength of acid increases, the strength of its conjugate base decreases.
As the strength of base increases, the strength of its conjugate acid decreases.
Example: HC2H3O2 is a weaker acid than HF. Which ion will more readily accept protons
in aqueous solution, C2H3O2- or F-?
HC2H3O2 is weaker acid; therefore, C2H3O2- is stronger base.
Thus C2H3O2- will accept a proton more readily.
Example: Pyridine, C5H5N, is a weaker base than ammonia, NH3. Which ion is the stronger
acid, NH4+, C5H5NH+?
C5H5N is weaker base; therefore, C5H5NH+ is stronger acid.
Therefore, C5H5NH+ will donate its proton easier that NH4+.
Pyridine is a base used in methods that determine the moisture content of foods.
THE AUTOIONIZATION OF WATER
- Water is not a nonelectrolyte, actually it is a very weak electrolyte.
4
- One out of every 10 million water molecules dissociates into an H+ ion and a OH- ion.
H2O (l)  H+ (aq) + OH- (aq)
Ion product of water
Write equilibrium expression
 H  OH 


Kc

 H O
2
- Since H2O is solvent, we will neglect it in equilibrium expression.
No matter what conditions exist in an aqueous solution, the following expression is true. (at
25 C)
 H   OH    1.0 1014
Ion product is also known as Kw.
K w   H   OH    1.0  1014
Aside: Other solvents autoionize as well.
NH3 + NH3  NH4+ + NH2CH3OH + CH3OH  CH3OH2+ + CH3OSince  H   OH    1.0 1014 is always true
- if we know [H+], we can calculate [OH-]
- if we know [OH-], we can calculate [H+]
Example: If a solution has [H+] = 1.8 x 10-5 M, what is [OH-]?
Example: If a solution has [OH-] = 9.4 x 10-12 M, what is [H+]?
Example: What is the [H+] and [OH-] for pure water?
H2O(l)  H+ (aq) + OH- (aq)
Kw   H
5
 OH 

[H+]
0
x
x
Initial
Change
Equil.
x 2  1.0  1014

[OH-]
0
x
x
x 2  x  1.0  107
 H    1.0  107 M


OH    1.0  107 M
Note: When [H+] = [OH-], the amount of acid and base are in equal amounts. The solution
is called neutral.
THE HYDRONIUM ION
H+ is simply a proton.
- A proton will hydrogen bond with water
H
H
H
H
O
O
+
H+
H
H3O+ is a hydronium ion.
Current theories of water treat proton as surrounded by 4 or 6 water molecules.
Therefore, H+ may be truly H9O4+ or H13O6+.
H
H
O
H
H
H+
O
O
H
H
O
H
H
H3O+ is a concept, not an actual structure.
Reconsider autoionization of water as
H2O (l) + H2O (l)  H3O+ (aq) + OH- (aq)
pH SCALE
Definition: pH = - log [H+] = -log[H3O+]
pH scale is used strictly for convenience
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- convenient to avoid scientific notation.
[H+] for pure water equals 10-7; therefore, pH = - log(10-7) = - (-7) = 7
All neutral solutions have pH = 7.
If solution is acidic, pH < 7.
If solution is basic, pH > 7.
pH scale of common substances
very acidic
slightly acidic
neutral
1
2
3
4
5
6
7
gastric acid
lemon juice
coke
slightly basic
coffee
very basic
8
9
10
11
12
13
milk of magnesia
household ammonia
household bleach
blood
**In a number that is a logarithm such pH, only the digits to the right of decimal point
are significant.**
[H+] = 5.38  10-2  pH = 1.269
[H+] = 5.38  10-5  pH = 4.269
[H+] = 5.38  10-8  pH = 7.269
[H+] = 5.38  10-12  pH = 11.269
Other definitions:
pOH = - log [OH-]
pKw = - log Kw = 14.00
Autoionization of water in terms of pH
- log([H+][OH-]) = -log Kw

pH + pOH = pKw = 14.00
Example: The pH of milk of magnesia is 10.0, what is the pOH?
pOH = pKw – pH = 14.00 – 10.0 = 4.0
Example: An sodium cyanide solution used to extract gold from low-grade ore has a
pH of 10.67, what is the hydronium ion concentration of the solution?
4 Au (s) + 8 CN- (aq) + O2 (aq) + H2O (l)  4 Au(CN)2- + 4 OH- (aq)
Example: A basic solution of trisodium phosphate (TSP), Na3PO4, from the hardware
store has a pOH of 2.14, what is its hydroxide ion concentration?
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You know that the proper name for TSP is sodium phosphate!
STRONG ACIDS AND BASES
Remember lists
Strong Acids: HCl, HBr, HI, HNO3, H2SO4, HClO3, HClO4
Strong Bases: Alkali metal hydroxides, Ba(OH)2, Sr(OH)2, Ca(OH)2, Mg(OH)2
limited solubility
Calculating pH of strong acid and strong base solutions
Strong acids and bases completely dissociate; therefore, [H+] (or [OH-]) equals the
concentration of the solute.
Example: What is the pH of a 0.0142 M solution of HBr?
Since HBr is a strong acid, [H+] = 0.0142 M
pH = - log (0.0142) = - (-1.848) = 1.848
Example: What the pH of 1.5 M Sr(OH)2 solution?
Since Sr(OH)2 is a strong base,
OH  
1.5molSr(OH)2
L

2 molOH
1molSr(OH)2
 3.0 M
1.0 1014 1.0 1014
 H   

 3.3 1015 M

3.0
OH 
pH = - log (3.3  10-15) = 14.48
OR
pOH = - log (3.0) = -0.48
pH = 14.00 – pOH = 14.00 – (-0.48) = 14.48
WEAK ACIDS AND BASES
For weak acids and bases, the concentration of acid (base) is not the concentration of
[H+] ([OH-]).
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Weak Acids
Dissociated ions are in equilibrium with undissociated molecule.
HC2H3O2 (aq)  H+ (aq) + C2H3O2- (aq)
Equilibrium expression is
 H C H O 



Ka
2
3
2
 HC H O 
2
3
2
Ka – acid-dissociation constant
- only difference from Kc is the label.
Each weak acid has its own Ka which does not change (at 25 C).
Example: What is the pH of 2.0 M ascorbic acid (HC6H7O6), also known as vitamin C?
Ka (HC6H7O6) = 8.0  10-5.
Vitamin C is necessary for the production of collagen, a fibrous protein that is found in skin,
tendon, blood vessels, organs, bone. Citrus fruit contains a substantial amount of vitamin C.
Example: What is the pH of 0.15 M iodic acid (HIO3)? Ka (HIO3) = 0.17.
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Assume that x is small.
x2
 017
.
015
.
x2  015
.  017
.   0.026
x = 0.16
Hmmm! Looks like x is not small. We will need the quadratic equation.
x2
Ka 
 017
.
 015
.  x
x2  017
.  015
.  x  0.026  017
. x
x2  017
. x  0.026  0
 b  b 2  4ac 0.17 
x

2a

 0.17 
2
 4 1  0.026
2 1
0.17  0.1329 0.17  0.36

 0.097
2
2
x = [H+] = 0.097  pH = 1.01
Iodic acid is used in the production of sodium iodate and potassium iodate. Both iodate
compounds are added to table salt to “iodize” the salt. Iodine is necessary for healthy thyroid
function.
Example: Calculate the pKa of an unknown acid solution with a concentration of
0.0988 M and a pH of 4.43.
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Polyprotic Acids
- Many acids have more than one proton to donate.
Consider H2SO4
H2SO4 is a strong acid.
 H2SO4 (aq)  H+ (aq) + HSO4- (aq)
Bisulfate ion also has a proton to donate, though it is less willing to do so.
HSO4- (aq)  H+ (aq) + SO42- (aq)
H SO   12.  10
HSO 
2

Ka 
4

2
 Ka 2
4
2nd proton disassociated
Consider H3PO4
Phosphoric acid has 3 protons to donate
- before second proton leaves, almost all of the first protons have left.
H  H PO   7.5  10
H PO 
H HPO   6.2  10

H PO 
H  PO   4.2  10

HPO 


K a1 
2
3
4
4
2

Ka 2
4
8

2
4
3

Ka 3
3
4
13
2
4
When we add H3PO4 to water, the disassociation is
H3PO4 (aq)  H+ (aq) + H2PO4- (aq)
How can we get the second proton to leave?
- Remove H+; i.e., add base.
H2PO4- (aq)  H+ (aq) + HPO42- (aq)
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Weak Bases
Conc. of OH- is not concentration of weak base.
Consider ammonia
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
Equilibrium expression is
 NH  OH 

 NH 

Kb

4
3
Kb – base-dissociation constant
- only difference from Kc is the label.
Consider general base
B (aq) + H2O (l)  BH+ (aq) + OH- (aq)
Equilibrium expression is
 BH  OH 


Kb

 B
Example: Calculate the OH- concentration of 0.012 M solution of pyridine, C5H5N.
Kb(C5H5N) is 1.7  10-9.
H
N
C
H
C
+
C
H
H2O

C
C
H
Assume that x is small.
H
H+
N
H
H
H
C
C
C
C
C
H
+
H
OH-
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Example: Calculate the pH of 3.4 M solution of soda ash, otherwise known as sodium
carbonate, Na2CO3. Kb(CO32-) = 1.8 10-4
When Na2CO3 is dissolved in water, it is fully disassociated. (All sodium salts are soluble.)
Na2CO3 (aq) = 2 Na+ (aq) + CO32- (aq)
Consider that CO32- is the conjugate base of HCO3-.
Assume that x is small.
x2
 18
.  104 x2  34
. 18
.  104   61
.  104
3.4
x = [OH-] = 2.510-2 M  pOH = 1.60  pH = 12.40
Check assumption:
0.025
100%  0.74%  0.74%  5%
3.4
Also 3.4 – 0.025 = 3.375  3.4
Assumption is good.
Soda ash is combined with calcium carbonate and sand (SiO 2) to make soda lime glass, the type of glass
most commonly used in bottles and windows.
Percent Ionization (Degree of Ionization)
While the dissociation constant gives a measure of the strength of a weak acid or base, it
doesn’t give any indication of how much electrolyte actually dissociates.
By definition, the percent ionization is the percentage of weak electrolyte molecules that
have dissociated in solution
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For a weak acid, the percent ionization is
 H  
electrolytedissociated
% ionization 
100% 
100% or
electrolyteinitial
 HA0
 A  
100% .
 HA0
Note the ambiguity in the definition.
- For pure weak acid, either definition is valid.
- If extra A- is added, use [H+] definition.
- If extra H+ is added, use [A-] definition.
- If in doubt, smaller number is correct.
For a weak base, the percent ionization is
OH  
electrolytedissociated
% ionization 
100% 
100%
electrolyteinitial
 B0
or
 HB 
100% .
 B0
- For pure weak base, either definition is valid.
- If extra HB+ is added, use [OH-] definition.
- If extra OH- is added, use [BH+] definition.
- If in doubt, smaller number is correct.
Example: Calculate the percent ionization of hydrofluoric acid in a 1.0 M, 0.10
M and 0.010 M solutions. Ka(HF) = 6.8  10-4
To calculate percent ionization, we need [H+], therefore, we need an ICE table.
 H    F 
K a       6.8 104
 HF
For a variable initial concentration c
Initial
Change
Equil.
[H+]
0
x
x
[F-]
0
x
x
[HF]
cM
–x
c–x
x2
Ka 
 6.8 104
c  x 
Assume that x is small.
x2
 6.8 104
c
x 2  c  6.8 104   x  c  6.8 104   [H  ]
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x
% ionization  100%
c


For c = 1.0 M, x  1.0 6.8 104  0.026
% ionization 
0.026
100%  2.6%
1.0


For c = 0.10 M, x  0.10 6.8 104  0.0082
0.0082
100%  8.2%
0.10
Oops! Now x is not small!
Ka 
x2
c  x 
 x2  Ka c  x   x 2  Ka x  Kac  0
2
 b  b 2  4ac  K a  K a  4 1   K a c 
x

2a
2 1

6.8 104 
 6.8 10 
4 2
2
4
6.8 10  2.7 10
2
 0.0079

 4 1   6.8 10 4  0.10 
c = [HF]0
1.0
0.10
0.010
4
[H+]
0.026
0.0079
0.0023

6.8 10 4  1.7 10 2
2
% ionized
2.6
7.9
23
pH
1.58
2.10
2.64
Note that as the concentration of weak acid decreases, more of it dissociates!
Artists use hydrofluoric acid to etch glass.
Relationship between Ka and Kb
Consider Kb of ammonia. Kb(NH3) = 1.8  10-5
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
 NH 4   OH  
Kb 
 NH 3 
15
Now consider Ka of conjugate acid, NH4+
NH4+ (aq) + H2O (l)  NH3 (aq) + H3O+ (aq)
H 3O   NH 3 

Ka 
 NH 4  
Now multiply Ka and Kb.
 NH  H O    NH  OH    H O  OH 
 NH 
 NH 


Ka  Kb 
3
3


4

4
Thus for a conjugate acid-base pair:
3
3
Ka  Kb = Kw
Example: If Kb(NH3) = 1.8  10-5, what is Ka for NH4+?
Example: If Ka(HF) = 6.8  10-4, what is Kb for F-?
Ka  Kb = Kw also implies pKa + pKb = 14.00
HYDROLYSIS
Many salts when dissolved in water will change its pH.
- Cations may attract the hydroxide ion.
- Anions may attract the hydrogen ion.
The process where ions cause water to split apart is called hydrolysis.

16
The acidic or basic properties of a salt solution depend on the relative abilities of the
cation and the anion to cause hydrolysis.
Cations are acidic (usually)
1. Cations from strong bases have negligible acidity.
- Li+, Na+, K+, Rb+, Cs+, Ca2+, Sr2+, Ba2+
2. Higher cationic charge causes greater hydrolysis.
- Higher charges attract OH- more effectively.
- pH of TlCl soln. higher than PbCl2 soln.
Pb(H2O)62+  Pb(H2O)5(OH)+ + H+
3. Smaller cations are more acidic than larger cations.
- Higher charge density attracts OH- more effectively.
- pH of Ga(NO3)3 sol. higher than Al(NO3)3 sol.
4. Transition metal ions are more acidic than alkali or alkaline earth metal ions.
- Incomplete d subshell attracts OH- and H2O.
- Co(ClO3)2 more acidic than Mg(ClO3)2.
Co(H2O)62+  Co(H2O)5(OH-)+ + H+
5. Conjugate acids of nitrogen bases are acidic.
- NH4+, C5H5NH+, (CH3)3NH+, etc…
NH4I (aq) + H2O (l)  NH3 (aq) + H3O+ (aq) + I- (aq)
Anions are basic
1. Anions from strong acids are not basic.
- Cl-, Br-, I-, NO3- HSO4-, ClO3-, ClO42. Other soluble monatomic anions are basic.
- K2S, Cs2O, BeH2 solutions are basic.
S2- + H2O  HS- + OHO2- + H2O  OH- + OHH- + H2O  H2 + OH3. Conjugate bases of weak acids are basic.
- Al(C2H3O2)3, SrF2, Li2C2O4 are basic salts.
F- + H2O  HF + OH- Note: This is the dissociation of a weak base.
Examples:
Predict whether the following salt solutions are acidic, basic or neutral.
a) KClO2
The K+ comes from a strong base; therefore, it should not affect the pH (it doesn’t attract
OH- to itself).
ClO2- comes from a weak acid; therefore, it should raise the pH (it attracts H+ a little).
17
b) Ni(NO3)2
The Ni2+ will attract OH-; therefore, it has the tendency to make the solution acidic.
The NO3- comes from a strong acid; therefore, it won’t affect the pH.
c) SrBr2
Sr2+ – neutral Br- – neutral
Salts derived from weak acid and weak base may be acidic or basic
- Both cation and anion hydrolyze water.
- pH depends on which ion has stronger acid/base properties.
- if base is stronger; i.e., Kb > Ka then pH > 7.
- if acid is stronger; i.e., Ka > Kb then pH < 7.
Example: Is a NH4NO2 solution acidic, basic or neutral?

K a NH 4



K b  NO2   
Kw
1.00 1014

 5.6 1010
K b  NH 3  1.8 105
Kw
Ka  HNO2 

100
.  1014
 2.2  1011
4.5  10 4
Ka > Kb, thus pH < 7. NH4NO2 solution is acidic.
Sodium nitrite is used in some meats as a preservative and to prevent the color of meat changing
from red to brown.
CHEMICAL STRUCTURE OF ACIDS AND BASES
Basic Principles of Acidity
1. Acidity is related to bond strength.
- H2Se is more acidic than H2S because the H – Se bond is weaker than the H – S
bond (Se is larger than S; therefore the bond overlap for the H – Se bond is weaker
than for the H – S bond).
2. Acidic hydrogens ionize from a polar bond.
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3. Polarity of the acidic hydrogen bond can be affected to a lesser extent by other
atoms.
4. Strong attraction of the water (solvent) molecules is the driving force.
Binary Acids
- Bonds between water molecules and ions are stronger than bonds between acid
molecules.
H
O
H
O
H
H
Cl
H
H
H
+
H
Cl
H
These bonds are
stronger than these
bonds.
O
O
H
H
Cl
Consider HF
- Bonds between HF molecules are very strong.
-
HF (aq) + HF (aq)  [F – H – F]- (aq) + H+ (aq)
H+
H
These bonds are stronger
than these bonds.
O
H
O
H
H
F-
- Even though HF is very polar, it is a weak acid since it forms strong intermolecular
bonds with itself.
HF etches glass (and bone!). It cannot be stored in a glass bottle.
Oxyacids
- Oxyacids are essentially molecular compounds.
- Bonds holding H+ must be strongly polar to allow H+ to disassociate.
- Must be considerable difference in electronegativity to create polar bond.
O
O Cl O H
O
polar bond
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- Secondary effects are very important.
- Electron withdrawing of other oxygen atoms is the difference between HClO4
being a strong acid rather than a weak acid.
Compare to HClO2.
O Cl O H
bond is less polar since fewer oxygen atoms
are drawing away electron density
Example: Which is the stronger acid, acetic acid or trichloroacetic acid?
H O
H C
Cl O
C O H
Cl
C
C O H
H
acetic acid
Cl
trichloroacetic acid
Chlorine atoms are much more electronegative that hydrogen atom
Chlorine atoms are much more electronegative that hydrogen atoms; therefore, the
chlorine atoms are making the O – H bond more polar thus increasing the acidity.
Trichloroacetic acid is used in some tattoo removal processes.
LEWIS ACIDS AND BASES
Definitions
Lewis acid – electron pair acceptor
Lewis base – electron pair donor
Lewis acid/base theory is more general than Brønsted – Lowry acid/base theory.
- Brønsted – Lowry acids are Lewis acids, though not vice versa.
- Brønsted – Lowry bases are Lewis bases, though not vice versa.
Consider H3O+ (aq) + OH- (aq)  H2O (l)
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H
O
H
H
O
H
H
O
H
O
H
H
+
H is electron pair acceptor
- Lewis acid
OH- is electron pair donor
- Lewis base
Consider NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
H
H
N
H
H
H
O
H
H
N
H
H
O
H
- NH3 is electron pair donor
- H2O is electron pair acceptor
Not all Lewis acid/base reactions are Brønsted-Lowry acid/base reactions.
Lewis acid – accepts e- pair from Cl-
Lewis base – donates e- pair via Cl- ion
AlCl3 + CH3Cl  AlCl4- + CH3+  other products
This is the first step in the “Friedel-Crafts alkylation” that adds a methyl group to an organic molecule.
More next year! 