1. Water Introduction Unusual Properties of Water Hydrogen Bond
HB in water
HB in other molecules
Interaction between water and charged molecules
Interaction between water and polar molecules
Interaction between water and non-polar molecules
Structure and function of biomolecules depended on weak interaction
Solute’s effect on colligative properties of aqueous solution
colligative properties
hypertonic solution, isotonic solution, hypotonic solution
Ionization of water

3. O: 1s22s22p4 4 sp3 orbitals Tetrahedron Geometry
FIGURE 2-1a Structure of the water molecule. (a) The dipolar nature of the H2O molecule is shown in a ball-and-stick model; the dashed lines represent the nonbonding orbitals. There is a nearly tetrahedral arrangement of the outer-shell electron pairs around the oxygen atom; the two hydrogen atoms have localized partial positive charges (δ+) and the oxygen atom has a partial negative charge (δ–).

4. Properties of Alkanes
Alkanes show regular increases in both boiling point and melting point as molecular weight increases
Due to the presence of weak dispersion forces between molecules
Dispersion forces increase as molecular size increases
A sufficient amount of energy is needed to overcome the dispersion forces and melt a solid or boil a liquid

6. FIGURE 2-1b Structure of the water molecule
FIGURE 2-1b Structure of the water molecule. (b) Two H2O molecules joined by a hydrogen bond (designated here, and throughout this book, by three blue lines) between the oxygen atom of the upper molecule and a hydrogen atom of the lower one. Hydrogen bonds are longer and weaker than covalent O—H bonds.

7. FIGURE 2-2 Hydrogen bonding in ice
FIGURE 2-2 Hydrogen bonding in ice. In ice, each water molecule forms four hydrogen bonds, the maximum possible for a water molecule, creating a regular crystal lattice. By contrast, in liquid water at room temperature and atmospheric pressure, each water molecule hydrogen-bonds with an average of 3.4 other water molecules. This crystal lattice structure makes ice less dense than liquid water, and thus ice floats on liquid water.

8. Atom Dependent Directional Dynamic Number Limited
FIGURE 2-5 Directionality of the hydrogen bond. The attraction between the partial electric charges (see Figure 2-1) is greatest when the three atoms involved in the bond (in this case O, H, and O) lie in a straight line. When the hydrogen-bonded moieties are structurally constrained (when they are parts of a single protein molecule, for example), this ideal geometry may not be possible and the resulting hydrogen bond is weaker.

9. Periodic Table of Elements: Electronegativity

10. FIGURE 2-3 Common hydrogen bonds in biological systems
FIGURE 2-3 Common hydrogen bonds in biological systems. The hydrogen acceptor is usually oxygen or nitrogen; the hydrogen donor is another electronegative atom.

11. FIGURE 2-4 Some biologically important hydrogen bonds.

14. FIGURE 2-6 Water as solvent
FIGURE 2-6 Water as solvent. Water dissolves many crystalline salts by hydrating their component ions. The NaCl crystal lattice is disrupted as water molecules cluster about the C– and Na+ ions. The ionic charges are partially neutralized, and the electrostatic attractions necessary for lattice formation are weakened.

15. FIGURE 2-7a Amphipathic compounds in aqueous solution
FIGURE 2-7a Amphipathic compounds in aqueous solution. (a) Long- chain fatty acids have very hydrophobic alkyl chains, each of which is surrounded by a layer of highly ordered water molecules.

16. FIGURE 2-7b (part 1) Amphipathic compounds in aqueous solution
FIGURE 2-7b (part 1) Amphipathic compounds in aqueous solution. (b) By clustering together in micelles, the fatty acid molecules expose the smallest possible hydrophobic surface area to the water, and fewer water molecules are required in the shell of ordered water. The energy gained by freeing immobilized water molecules stabilizes the micelle.

17. FIGURE 2-7b (part 2) Amphipathic compounds in aqueous solution
FIGURE 2-7b (part 2) Amphipathic compounds in aqueous solution. (b) By clustering together in micelles, the fatty acid molecules expose the smallest possible hydrophobic surface area to the water, and fewer water molecules are required in the shell of ordered water. The energy gained by freeing immobilized water molecules stabilizes the micelle.

18. FIGURE 2-7b (part 3) Amphipathic compounds in aqueous solution
FIGURE 2-7b (part 3) Amphipathic compounds in aqueous solution. (b) By clustering together in micelles, the fatty acid molecules expose the smallest possible hydrophobic surface area to the water, and fewer water molecules are required in the shell of ordered water. The energy gained by freeing immobilized water molecules stabilizes the micelle.

19. FIGURE 2-8 Release of ordered water favors formation of an enzyme- substrate complex. While separate, both enzyme and substrate force neighboring water molecules into an ordered shell. Binding of substrate to enzyme releases some of the ordered water, and the resulting increase in entropy provides a thermodynamic push toward formation of the enzyme- substrate complex (see p. 192).

22. COLLIGATIVE PROPERTIES
Melting Point
Boiling Point
Vapor Pressure
Osmotic Pressure

23. FIGURE 2-11 Osmosis and the measurement of osmotic pressure
FIGURE 2-11 Osmosis and the measurement of osmotic pressure. (a) The initial state. The tube contains an aqueous solution, the beaker contains pure water, and the semipermeable membrane allows the passage of water but not solute. Water flows from the beaker into the tube to equalize its concentration across the membrane. (b) The final state. Water has moved into the solution of the nonpermeant compound, diluting it and raising the column of water within the tube. At equilibrium, the force of gravity operating on the solution in the tube exactly balances the tendency of water to move into the tube, where its concentration is lower. (c) Osmotic pressure (Π) is measured as the force that must be applied to return the solution in the tube to the level of that in the beaker. This force is proportional to the height, h, of the column in (b).

24. Hypertonic solutions are those in which more solute (and hence lower water potential) is present. Isotonic solutions have equal (iso-) concentrations of substances. Water potentials are thus equal, although there will still be equal amounts of water movement in and out of the cell, the net flow is zero. Hypotonic solutions are those with less solute (again read as higher water potential).

27. FIGURE 2-12 Effect of extracellular osmolarity on water movement across a plasma membrane. When a cell in osmotic balance with its surrounding medium—that is, a cell in (a) an isotonic medium—is transferred into (b) a hypertonic solution or (c) a hypotonic solution, water moves across the plasma membrane in the direction that tends to equalize osmolarity outside and inside the cell.

28. IONIZATION OF WATER

29. Concentration Definition
Molarity
Nomality
Molality
Percentage Concentration
Unit Conversion

30. CONCENTRATION
Concentration: amount of substance present in a given volume of solution.
Concentrated Solution: solution with a high concentration.
Dilute Solution: solution with low concentration.
Mole: amount of substance.
Molar Concentration: the number of moles of the substance contained in 1L of solution, i.e., moles of solute/L of a solution.
Percentage Concentration.
Very Low Concentration.
Normality

31. c = molar concentration (mol/L) n = number of moles (mol)
To find the Molarity of a solution, one has to use the following formula:
c = molar concentration (mol/L)
n = number of moles (mol)
V = volume (L)
For example: What is the [NaCl] in a solution containing 6.25 g of NaCl in L of solution?
*1 mole of NaCl = 58.5 g
*since mass and volume are given, all we need to do now is to first convert the mass into moles and then use the moles and volume to find Molarity of the solution.

32. If a 2.0 L of solution contains 8.0 mol of NaCl,
Molarity can be determined if the amount of solution and solvent is given. For example:
If a 2.0 L of solution contains 8.0 mol of NaCl,
When expressed in words, the unit symbol “M” is written as “molar”.
The easier way of writing “molar concentration of…” is using a set of brackets: […].
For example: If a 2.0 L of solution contains 5.0 mol of NaCl, the molar concentration can be written as:
molar concentration of NaCl= 2.5 M
[NaCl]= 2.5 M
The Molarity of the sodium chloride is 2.50 molar.

33. Percentage Concentration
1) % concentration can be in V/V, W/W, or W/V
Like most %s, V/V and W/W need to have the same units on top and bottom.
W/V is sort of in the same units; V is mostly water and water’s density is 1 g/mL or 1 kg/L
3 g H2O2/100 mL solution ≈ 3 g H2O2/100 g solution

34. Very Low Concentration
2) Expressing concentrations in parts per million (ppm) requires the unit on top to be 1,000,000 times smaller than the unit on the bottom
E.g. 1 mg/kg or μg/g
Multiples of 1000 are expressed in this order
μ_, m_, _, k_ (“_” is the base unit)
For parts per billion (ppb), the top unit would have to be 1,000,000,000 times smaller
3) Molar concentration is the most commonly used in chemistry. Ensure that units are mol/L.

35. Nomality
The number of “equivalent weights” of solute per liter of solution
Another way to express concentration
Used when it’s important to know how many “reactive groups” are present in a solution rather than how many molecules

36. Reactive groups and equivalent weight
Used in reference to acids and bases
Acids dissociate in solution to release H+ ions
Bases dissociate to release OH- ions
For an acid:
1 equivalent weight is equal to the number of grams of that acid that reacts to yield 1 mole of H+ ions, that is, 1 mole of reactive groups
For a base:
1 equivalent weight is equal to the number of grams of that base that supplies 1 mole of OH-, that is , 1 mole of reactive groups

37. NaOH  Na+ OH-
1 mole of NaOH dissociates to produce 1 mole of Na+ and 1 mole of OH-
Because 1 equivalent weight of NaOH is the number of grams that will produce 1 mole of OH- ions, the equivalent weight of NaOH is the same as its molecular weight.

38. 1 mole of sulfuric acid dissociates to form 2 moles of H+ ions.
H2SO4SO H+
1 mole of sulfuric acid dissociates to form 2 moles of H+ ions.
It only requires 0.5 mole of H2SO4 to produce 1 mole of H+ ions.
Therefore, the FW of H2SO4 = 98.1, but its equivalent weight is 98.1 x 0.5 = 49.1
1 M = 2 N

39. Concentration after Dilution
For Example: 20 mL 0.5 M NaCl is diluted to 200 mL, to find the final NaCl concentration, we do the following calculation:

40. Unit Conversion Weight: 1 g = 1000 mg = 106 mg
Length: 1 m = 10 dm = 100 cm = 1000 mm
Area: 1 m2 = 100 dm2
Volume: 1 dm3 = 1 L = 1000 cm3 = 1000 c.c. =1000 mL=106 mL

43. Concentration Conversion

44. BrØnsted-Lowry Acids & Bases
A BrØnsted-Lowry acid is a proton donor.
A BrØnsted-Lowry base is a proton acceptor.
H+ = proton

45. Reactions of BrØnsted-Lowry Acids and Bases
Loss of a proton from an acid forms its conjugate base.
Gain of a proton by a base forms its conjugate acid.
A double reaction arrow is used between starting materials and products to indicate that the reaction can proceed in the forward and reverse directions. These are equilibrium arrows.
Examples:

46. Lewis Acids and Bases A Lewis acid is an electron pair acceptor.
A Lewis base is an electron pair donor.
Lewis bases & BrØnsted-Lowry bases - both have an available electron pair—a lone pair or an electron pair in a π bond.
A BrØnsted -Lowry base always donates this electron pair to a proton, but a Lewis base donates this electron pair to anything that is electron deficient.
A Lewis acid must be able to accept an electron pair, but there are many ways for this to occur.
All BrØnsted-Lowry acids are also Lewis acids, but the reverse is not necessarily true.
Any species that is electron deficient and capable of accepting an electron pair is also a Lewis acid.

47. Lewis Acids A Lewis acid is also called an electrophile.
When a Lewis base reacts with an electrophile other than a proton, the Lewis base is also called a nucleophile. In this example, BF3 is the electrophile and H2O is the nucleophile.
Lewis acid-base reactions illustrate a general pattern in organic chemistry. Electron-rich species react with electron-poor species.
In the simplest Lewis acid-base reaction one bond is formed and no bonds are broken..
electrophile
nucleophile

48. Lewis Acids & Bases Other good examples involve metal ions.
Such bonds as the H2O ---> Co bond are often called COORDINATE COVALENT BONDS because both electrons are supplied by one of the atoms of the bond.

49. Acid Strength and pKa
Acid strength is the tendency of an acid to donate a proton.
The more readily a compound donates a proton, the stronger an acid it is.
Acidity is measured by an equilibrium constant
When a BrØnsted-Lowry acid H—A is dissolved in water, an acid-base reaction occurs, and an equilibrium constant can be written for the reaction.

50. Acid Strength and pKa
Because the concentration of the solvent H2O is essentially constant, the equation can be rearranged and a new equilibrium constant, called the acidity constant, Ka, can be defined.
It is generally more convenient when describing acid strength to use “pKa” values than Ka values.

51. Factors that Determine Acid Strength
Anything that stabilizes a conjugate base A:¯ makes the starting acid H—A more acidic.
Four factors affect the acidity of H—A. These are:
Element effects
Inductive effects
Resonance effects
Hybridization effects
No matter which factor is discussed, the same procedure is always followed. To compare the acidity of any two acids:
Always draw the conjugate bases.
Determine which conjugate base is more stable.
The more stable the conjugate base, the more acidic the acid.

52. 1. Element Effects—Trends in the Periodic Table.
Across a row of the periodic table, the acidity of H—A increases as the electronegativity of A increases.
Positive or negative charge is stabilized when it is spread over a larger volume.

53. 2. Inductive Effects
An inductive effect is the pull of electron density through σ bonds caused by electronegativity differences in atoms.

54. 3. Resonance Effects
Resonance is a third factor that influences acidity.
In the example below, when we compare the acidities of ethanol and acetic acid, we note that the latter is more acidic than the former.
When the conjugate bases of the two species are compared, it is evident that the conjugate base of acetic acid enjoys resonance stabilization, whereas that of ethanol does not.

55. 3. Resonance Effects
Resonance delocalization makes CH3COO¯ more stable than CH3CH2O¯, so CH3COOH is a stronger acid than CH3CH2OH.
The acidity of H—A increases when the conjugate base A:¯ is resonance stabilized.

56. 4. Hybridization Effects
The final factor affecting the acidity of H—A is the hybridization of A.
Let us consider the relative acidities of three different compounds containing C—H bonds.
The higher the percent of s-character of the hybrid orbital, the closer the lone pair is held to the nucleus, and the more stable the conjugate base.

57. 4. Hybridization Effects

58. Summary of Factors that Determine Acidity

63. FIGURE 2-13 Proton hopping
FIGURE 2-13 Proton hopping. Short “hops” of protons between a series of hydrogen-bonded water molecules result in an extremely rapid net movement of a proton over a long distance. As a hydronium ion (upper left) gives up a proton, a water molecule some distance away (lower right) acquires one, becoming a hydronium ion. Proton hopping is much faster than true diffusion and explains the remarkably high ionic mobility of H+ ions compared with other monovalent cations such as Na+ and K+.

64. pH of pure Water

69. FIGURE 2-14 The pH of some aqueous fluids.

70. How to Find pH of A Solution?
A Strong Acid or Strong Base Solution
A Weak Acid or A Weak Base Solution
A Solution with Both Weak Acid and Weak Base (Buffer)

71. pH of Strong Acid/Base
Strong Acids: HCl, HBr, HI, HClO4, H2SO4, HNO3, etc. (Leveling Effect: which one is the strongest acid?)
Strong Bases: NaOH, KOH, RbOH, CsOH
For HCl, [H+] = [HCl], so pH = -log[H+] = - log[HCl]
For NaOH, [OH-] = [NaOH], pOH = -log[OH-] = -log[NaOH], pH = 14 - pOH

72. Conjugate Acid-Base Pairs
Chemical entities that differ only by H+ in the equation are considered conjugate acid-base pairs:
e.g. CH3COOH(aq) and CH3COO–(aq)
H3O+(aq) and H2O(aq)
H H+
e.g. CH3COOH (aq) + H2O (aq) H3O+ (aq) + CH3COO– (aq)
B-L acid B-L base B-L acid B-L base
conj. acid-base pair

73. Strengths of Conjugate Acid – Base pairs
In chemical equilibrium, we have a competition between forward and reverse reactions. For acid and base reactions, this means we have a competition between two sets of acids and bases
i.e acid 1 + base acid 2 + base 2
If acid 1 and base 1 are stronger than acid 2 and base 2, then the forward reaction will dominate
If acid 2 and base 2 are stronger than acid 1 and base 1, then the reverse reaction will dominate
The stronger the acid, the weaker it’s conjugate base and the weaker the acid, the stronger it’s conjugate base.

74. pH of Weak Acid

75. If Ca / Ka > 100, you can simplify Ka expression to just Ka = (x2 / Ca)

76. Ka = X2/C, X = 0.00209 M, almost no error.
pH = -log( ) = 4.18.

77. Buffers Buffer Solutions resist a change in pH
Buffers contain relatively large concentrations of either
An acid, HA and its conjugate base A-
A base, B, and its conjugate acid (BH+)

78. Conjugate Acid-Base Pairs
Chemical entities that differ only by H+ in the equation are considered conjugate acid-base pairs:
e.g. CH3COOH(aq) and CH3COO–(aq)
e.g. CH3COOH (aq) + H2O (aq) H3O+ (aq) + CH3COO– (aq)

79. For a general equation:

80. Titration Terminology
Titrant - solution of known concentration used in titration
Analyte - substance being analyzed
Standard Solution - A solution whose concentration is known accurately.

81. Primary Standard Solution -Prepared by dissolving carefully weighed solid sample in enough water to give an accurately known volume of solution.
Equivalent Point - point at which stoichiometrically equivalent amounts of reactants have reacted.
End Point - the point at which the indicator changes color

82. FIGURE 2-15 Conjugate acid-base pairs consist of a proton donor and a proton acceptor. Some compounds, such as acetic acid and ammonium ion, are monoprotic; they can give up only one proton. Others are diprotic (carbonic acid and glycine) or triprotic (phosphoric acid). The dissociation reactions for each pair are shown where they occur along a pH gradient. The equilibrium or dissociation constant (Ka) and its negative logarithm, the pKa, are shown for each reaction. *For an explanation of apparent discrepancies in pKa values for carbonic acid (H2CO3), see p. 63.

83. FIGURE 2-16 The titration curve of acetic acid
FIGURE 2-16 The titration curve of acetic acid. After addition of each increment of NaOH to the acetic acid solution, the pH of the mixture is measured. This value is plotted against the amount of NaOH added, expressed as a fraction of the total NaOH required to convert all the acetic acid (CH3COOH) to its deprotonated form, acetate (CH3COO–). The points so obtained yield the titration curve. Shown in the boxes are the predominant ionic forms at the points designated. At the midpoint of the titration, the concentrations of the proton donor and proton acceptor are equal, and the pH is numerically equal to the pKa. The shaded zone is the useful region of buffering power, generally between 10% and 90% titration of the weak acid.

84. FIGURE 2-17 Comparison of the titration curves of three weak acids
FIGURE 2-17 Comparison of the titration curves of three weak acids. Shown here are the titration curves for CH3COOH, H2PO4–, and NH4+. The predominant ionic forms at designated points in the titration are given in boxes. The regions of buffering capacity are indicated at the right. Conjugate acid-base pairs are effective buffers between approximately 10% and 90% neutralization of the proton-donor species.

85. FIGURE 2-18 The acetic acid–acetate pair as a buffer system
FIGURE 2-18 The acetic acid–acetate pair as a buffer system. The system is capable of absorbing either H+ or OH– through the reversibility of the dissociation of acetic acid. The proton donor, acetic acid (HAc), contains a reserve of bound H+, which can be released to neutralize an addition of OH– to the system, forming H2O. This happens because the product [H+][OH– ] transiently exceeds Kw (1 x 10–14 M2). The equilibrium quickly adjusts to restore the product to 1 x 10–14 M2 (at 25°C), thus transiently reducing the concentration of H+. But now the quotient [H+][Ac–]/[HAc] is less than Ka, so HAc dissociates further to restore equilibrium. Similarly, the conjugate base, Ac–, can react with H+ ions added to the system; again, the two ionization reactions simultaneously come to equilibrium. Thus a conjugate acid-base pair, such as acetic acid and acetate ion, tends to resist a change in pH when small amounts of acid or base are added. Buffering action is simply the consequence of two reversible reactions taking place simultaneously and reaching their points of equilibrium as governed by their equilibrium constants, Kw and Ka.

86. pKa of molecules for buffers in biochemistry studies

87. FIGURE 2-20 The bicarbonate buffer system
FIGURE 2-20 The bicarbonate buffer system. CO2 in the air space of the lungs is in equilibrium with the bicarbonate buffer in the blood plasma passing through the lung capillaries. Because the concentration of dissolved CO2 can be adjusted rapidly through changes in the rate of breathing, the bicarbonate buffer system of the blood is in near equilibrium with a large potential reservoir of CO2.

88. FIGURE 2-21 The pH optima of some enzymes
FIGURE 2-21 The pH optima of some enzymes. Pepsin is a digestive enzyme secreted into gastric juice, which has a pH of ~1.5, allowing pepsin to act optimally. Trypsin, a digestive enzyme that acts in the small intestine, has a pH optimum that matches the neutral pH in the lumen of the small intestine. Alkaline phosphatase of bone tissue is a hydrolytic enzyme thought to aid in bone mineralization.

89. Find out the pH of 0.01 M HCl and 1.0 M NaOH solution
Calculate the pH of 0.2 M AcOH with a Ka of x 10-5M
Calculate the pH of sodium acetate and acetic acid buffer at the following concentration: [NaAc] = 0.17 M, [AcOH] = 0.35 M
Prepare the 4 L of PBS buffer by given the condition that [NaCl] = 0.15 M, [PO43-] = M, pKa = 6.91.