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Transcript of Water
Ionization of water
To understand the chemistry of living systems (biochemistry), it is important to understand the physical and chemical properties of water.
Why is water so important?
1. Water determines the shape and consequently, the function of all biological macromolecules.
2. Water is the solvent of life! The human body is 70% water. Most biochemical reactions occur in water.
3. Water participates in many biochemical reactions.
4. The oxidation of water to produce O2, is a fundamental reaction of photosynthesis.
The geometry of water is nearly tetrahedral because of the two lone pairs of electrons on the oxygen atom.
What is a polar bond:
• Electrons are unequally shared,more negative charge found closer to one atom.
• Due to difference in electronegativity of atoms involved in bond.
What makes water polar? – it’s physical properties
Electronegativity: a measure of the force of an atom’s attraction for electrons it shares in a chemical bond with another atom
Oxygen and Nitrogen, more electronegative than carbon and hydrogen
Fluorine is most electronegative (4)
Molecules such as CO2 have polar bonds but, given their geometry, are nonpolar molecules; that is, they have a zero dipole moments
Polar Bonds & Molecules
(van der Waals forces)
Key factors in
Various types of
Hydration Shells Surrounding Ions in Water
Ion-dipole and Dipole-dipole Interactions
• Ion-dipole and dipole-dipole interactions help ionic and polar compounds dissolve in water
The formation of a H-bond requires a donor (D = O-H or N-H) and an acceptor (A = O or N).
Hydrogen bond: the attractive interaction between dipoles when:
positive end of one dipole is a hydrogen atom bonded to an atom of high electronegativity, most commonly O or N, and
the negative end of the other dipole is an atom with a lone pair of electrons, most commonly O or N
The electrostatic attraction of the water dipoles is responsible for the ability of water to form hydrogen bonds (H-bonds).
H-bonds - weak electrostatic interactions where a hydrogen atom covalently bonded to an electronegative atom is shared with another electronegative atom (O or N, sometimes S).
Hydrogen bond is non-covalent
• Each water molecule can be involved in 4 hydrogen bonds: 2 as donor, and 2 as acceptor
• Due to the tetrahedral arrangement of the water molecule (Refer to Figure 2.1).
Interesting and Unique Properties of Water
The energy of a H-bond is about 20 KJ/mol (small compared to energy of O-H covalent bond (460 KJ/mol)).
If H-bond is so weak, how does it that explain its properties above?
Despite the relative weakness of a H-bond, the enormous amount of these bonds formed in water, give water its remarkable properties.
Does methane have H-bonding?
Even though hydrogen bonds are weaker than covalent bonds, they have a significant effect on the physical properties of hydrogen-bonded compounds
• Hydrogen bonding is important in stabilization of 3-D structures of biological molecules such as: DNA, RNA, proteins.
Other Biologically Important Hydrogen bonds
Ionic compounds (e.g.,KCl) and low-molecular- weight polar covalent compounds (e.g., C2H5OH and CH3COCH3) tend to dissolve in water
The underlying principle is electrostatic attraction of unlike charges; the positive dipole of water for the negative dipole of another molecule, etc. (Coulomb’s law and dielectric constants)
What are the types of interactions that can occur between molecules?
Interactions that water can have with other molecules:
ion-dipole interaction: e.g., KCl dissolved in H2O
dipole-dipole interactions: e.g., ethanol or acetone dissolved in H2O
dipole induced-dipole interactions: weak and generally do not lead to solubility in water
Solvent Properties of H2O
Has a high dielectric constant
For a solute to be soluble in a solvent requires that the solvent interact with the solute more strongly than solute particles can interact with each other.
Water is an excellent solvent for polar or ionic (hydrophilic) substances
“like dissolves like”
Non-polar (hydrophobic) substances are insoluble in water
“oil and water don’t mix”
Solvent Properties of H2O
A: Water tends to hydrate the hydrophilic portion and exclude the hydrophobic portion, leading to the aggregation of the amphiphiles.
What happens when water interacts with amphiphilic molecules ?
• both polar and nonpolar character
Interaction between nonpolar molecules is very weak
called van der Waals interactions
They are stabilized by hydrophobic forces (hydrophobic effect) - describes the tendency of water to minimize its contact with nonpolar molecules (oil and water
formation depends on the attraction between temporary induced dipoles
When water interacts with amphiphilic molecules
Micelles and bilayers form…
A/B versus CA/CB
pH vs. pKa meaning
HA vs. A-
B- vs. BH
Acid: a molecule that behaves as a proton donor
Strong base: a molecule that behaves as a proton acceptor
Acids, Bases and pH
One can derive a numerical value for the strength of an acid (amount of hydrogen ion released when a given amount of acid is dissolved in water).
Describe by Ka:
Lets quantitatively examine the dissociation of water:
• Molar concentration of water (55M)
• Kw is called the ion product constant for water.
• Must define a quantity to express hydrogen ion concentrations…pH
Ionization of H2O and pH
Fig. 2-12, p. 47
Equation to connect Ka to pH of solution containing both acid and base.
We can calculate the ratio of weak acid, HA, to its conjugate base, A-, in the following way
From this equation, we see that
when the concentrations of weak acid and its conjugate base are equal, the pH of the solution equals the pKa of the weak acid
WHY? What does this mean? (Evident in the buffer curves, follow arrow)
pKa is the pH at which the group will be ionized (lose an H+ from its ionizable group)
when pH < pKa, the weak acid predominates
Means 'group' is protonated – the acid in buffer donated H+ to the 'group'
when pH > pKa, the conjugate base predominates
Means 'group' is deprotonated – the base in buffer accepted H+ from the 'group'
Henderson-Hasselbalch Equation (Cont’d)
Aspirin is an acid with a Pka of 3.5; its structure includes a carboxyl group. To be absorbed into the bloodstream, it must pass through the membrane lining the stomach and the small intestine. Electrically neutral molecules can pass through a membrane more easily than can charged molecules.
Would you expect more aspirin to be absorbed in the stomach, where the pH of gastric juice is about 1, or in the small intestine, where the pH is about 6? Explain your answer.
Titration: an experiment in which measured amounts of acid (or base) are added to measured amounts of base (or acid)
Equivalence point: the point in an acid-base titration at which enough acid has been added to exactly neutralize the base (or vice versa)
a monoprotic acid releases one H+ per mole
a diprotic acid releases two H+ per mole
a triprotic acid releases three H+ per mole
Calculate pH of solution after calculating the relative amts of acetic acid (pKa = 4.76) and acetate ion present at the following points when 1 mol acetic acid it titrated with NaOH:
A. 0.1 mol of NaOH added
B. 0.5 mol of NaOH added
C. 0.9 mol of NaOH added
buffer: a solution whose pH resists change upon addition of either more acid or more base
consists of a weak acid and its conjugate base
Examples of acid-base buffers are solutions containing
CH3COOH and CH3COONa
H2CO3 and NaHCO3
NaH2PO4 and Na2HPO4
Buffers prepared by mixing a weak acid or weak base with a salt of that acid or base
H2PO4-/HPO42- is the principal buffer in cells
H2CO3/HCO3- is an important (but not the only) buffer in blood
hyperventilation can result in increased blood pH
hypoventilation can result in decreased blood pH
Naturally Occurring Buffers
The following criteria are typical
suitable pKa (+/- 1 pH unit from pH of rxn; ½ even better)
no interference with the reaction or detection of the assay
suitable ionic strength
its non-biological nature
Selecting a Buffer
Maintaining a relatively constant pH is immensely important in biological systems. This is achieved through buffers
Fig. to right illustrates how buffers work
When [HA] appox = [A-] (at the pKa), the pH of the solution is relatively insensitive to the addition of strong base or strong acid
pH = pKa + log [A-]/[HA]
A buffer is effective at a pH = pKa ± 1 pH unit.
Fig. 2-15a, p. 53
pH = pKa +/- 1 pH unit
Selecting a buffer
Buffer capacity is related to the concentrations of the weak acid and its conjugate base
the greater the concentration of the weak acid and its conjugate base, the greater the buffer capacity
0.1 < [A-]/[HA] < 10
Calculate the pH of a buffer solution prepared by mixing 75 mL of 1.0 M lactic acid (pKa = 3.86) and 25 mL of 1.0 M sodium lactate.
Define buffering capacity. How do the following buffers differ in buffering capacity? In pH?
Buffer A: 0.01 M Na2HPO4 and 0.01 M NaHPO4
Buffer B: 0.10 M Na2HPO4 and 0.10 M NaHPO4
Buffer C: 1.0 M Na2HPO4 and 1.0 M NaHPO4
Ex) NH3, CO, NaCl
Calculating pH: strong acid vs. weak acid -->
how do you solve?
Calculating pH directly or using an ICE table???
Write out rxn:
Characterisitics of buffers
You want a buffer with a pH of 4.90. Would sodium acetate/acetic acid (Ka = 1.7 x 10^-5) be a suitable buffer?
Practice Problems from Text:
1, 2, 6, 11, 18-20, 24