Chapter 2 - Small Molecules
Atoms: the constituents of matter
Nucleus: protons and neutrons
Mass of the proton = 1 dalton (1 D)
Electrons in orbitals about the nucleus
An element is made up of only one kind of atom
The number of protons identifies the element
Isotopes differ in the number of neutrons
*Electron behavior determines chemical bonding*
Chemical reactions are changes in the atomic composition of substances
All chemical reactions involve rearrangements of electrons as atoms change partners
When atoms share electrons, they are bonded together
A molecule is defined as two or more atoms linked by chemical bonds
Chemical bonds: linking atoms together (Table 2.1)
Covalent bonds consist of shared pairs of electrons
Electronegativity: unequal sharing of electrons
Some atoms are electronegative (O and N) - they draw the electrons closer
Polar covalent bond: partial electrical charges (d -; d +) on the atoms forming the bond. Water is a good example (Figure 2.8)
*Hydrogen Bonds*: Very important in biological systems
Hydrogen bonds may form between any covalently bonded hydrogen and an electronegative atom (usually O or N) nearby
Ions form bonds by electrical attraction (anions and cations)
Nonpolar substances have no attraction for polar substances
Hydrocarbons are nonpolar
Substances with electronegative atoms (such as O or N) are often polar
Nonpolar molecules show van der Waals attractions
*Hydrophobic Interactions*: Very important in biological systems
Nonpolar molecules come into contact with one another (coalesce) in aqueous solutions, not because of any intrinsic affinity for one another, but because clustering of nonpolar molecules allows more water molecules to achieve a more stable, H-bonded arrangement.
1. In water, the water molecules interact extensively with one another through hydrogen bonds (3.4 H bonds per water molecule at 0° C; 2.3 H bonds per water molecule at 10° C). Stable state
2. The introduction of nonpolar molecules into water puts nonpolar molecules in the way of water molecules seeking to form H bonds with their neighbors. Each nonpolar molecule forces water to form a"cage" around it (Figure 2.14). Instability (water molecules can't achieve their most stable state)
3. Nonpolar molecules form clusters within a common"cage" (hydrophobic interactions) so that the total number of water molecules taken up in forming"cages" is minimized (more water molecules are free to interact via mutual H-bonding). Recovery of stable state
Eggs by the Dozen, Molecules by the Mole
Molecular formulas, structural formulas, molecular weight, definition of a mole, Avogadro's number, molarity (concentrations of solutes in solution)
Chemical Reactions: Atoms Change Partners
reactants, products, changes in chemical potential energy, calories and joules
WATER: Structure and Properties
Water (H2O): Emergent properties - water is very different from hydrogen (H2) and oxygen (O2)
Water is a polar molecule (Figure 2.18)
Water readily forms H bonds
These strong interactions between water molecules (cohesiveness) account for its high heat capacity
(1 cal/g/° C), heat of vaporization (540 cal/g), heat of fusion (-80 cal/g), dielectric constant (78.5 D), surface tension, and capillarity.
Properties of Solutions: Acids, Bases, and Buffers
Dissociation of Water Molecules
KD = [H+][OH-]/[H2O] = 1.8 x 10-16 M
where KD is the dissociation constant (equilibrium constant) for the dissociation of a proton from a water molecule.
Since KD is very small, there is very little propensity for an H+ to dissociate from an H2O molecule.
Or, put the other way, OH- has a very high affinity for H+
Since [H2O] = 55.5 M,
[H+][OH-] = KD(55.5 M) = 10-14 M2
From the stoichiometry of the reaction for dissociation of water, it is obvious that
[H+] = [OH-] = 10-7 M in pure water
pH - an index of the relative concentration of H+ ions in solution
pH º -log10[H+]
In pure water, pH = 7
Note also that, because [H+][OH-] must always equal 10-14 M2 in aqueous solutions,
pH + pOH = 14
pH values of some familiar substances: Figure 2.21
The pH scale ranges from 0 to 14
Keep in mind that the pH scale is a negative logarithmic scale
The higher the pH, the lower the [H+] (alkaline, basic solutions have a high pH)
The lower the pH, the higher the [H+] (acidic solutions have a lower pH)
Example:
orange juice: pH = 4.3
grapefruit juice: pH = 3.2
The concentration of H+ in 12 times higher in grapefruit juice than in orange juice.
Biological systems are very sensitive to [H+] and [OH-]
Therefore, pH is very important
The normal pH of blood plasma = 7.4 ([H+] = 0.00000004 M)
Life is jeopardized if the pH of blood plasma falls to 6.8
([H+] = 0.00000016 M, a 4-fold increase).
Acids and Bases: Acids donate H+; Bases accept H+
Two classes: Strong Electrolytes and Weak Electrolytes
Strong Electrolytes: Substances that, when dissolved in water, dissociate completely. (By readily forming ions, such substances dramatically raise the electrical conductivity of water.)
Sodium Chloride: NaCl ® Na+ + Cl- salt
Hydrochloric Acid: HCl ® H+ + Cl- strong acid
Sodium Hydroxide: NaOH ® Na+ + OH- strong base
For strong electrolytes, KD » ¥
Weak Electrolytes: Substances that dissociate in water only to a small extent (KD » 10-3 M to 10-11 M) (and only slightly increase the conductivity of water).
Carbonic Acid: H2CO3 _ H+ + HCO3- KD = 1.7 x 10-4 M
Carbonic acid dissociates only 1% or so in water. That is, HCO3- (bicarbonate ion) has a high affinity for H+ (especially when compared to Cl-)
Buffers: Substances that minimize changes pH (in [H+] or [OH-]) when H+ (or OH-) are added to, or generated in, a solution. (Figure 2.22)
Buffers are typically prepared by combining a solution of a weak acid and a solution of its salt, such as H2CO3 and NaHCO3, to form a mixture of the weak acid and its conjugate base.
The carbonic acid/bicarbonate system (a mixture of H2CO3 and HCO3-) is an important biological buffer
H2CO3 _ H+ + HCO3-
Response to an increase in H+ (HCO3- acts as a proton acceptor):
H+ + HCO3- ® H2CO3
Response to an increase in OH- (H2CO3 acts as a proton donor):
OH- + H2CO3 ® H2O + HCO3-
Changes in pH, the concentration of a weak acid (HA) and its conjugate base (A-) as OH- is added to a solution of HA.