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Biochemistry 403 > Chapter 2: Water > Flashcards

Chapter 2: Water Flashcards

1
Q

_____ ______ Gives Water Its Unusual Properties

A

Hydrogen Bonding

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2
Q

Lecture

because of the unequal electron sharing between each H and O in water, each hydrogen atom bears a ____ _____ charge, and the oxygen atom bears a ______ ______ charge equal in magnitude to the _____ of the _____ ______ _____. As a result, there is an _____ ______ between the oxygen atom of one water molecule and the hydrogen of another called a _____ ______

A
  • partial positive charge (δ+)
  • partial negative charge
  • sum
  • two partial positives (2δ−)
  • electrostatic attraction
  • hydrogen bond
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3
Q

Lecture

The nearly tetrahedral arrangement of the orbitals about the oxygen atom allows each water molecule to form hydrogen bonds with as many as ____ neighboring water molecules

A

four

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4
Q

Lecture

When ice melts or water evaporates, _____ is taken up by the ____

A
  • heat
  • system
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5
Q

Lecture

Gibs free energy focuses on the _____

A

system

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6
Q

Lecture

Gibs free energy formula

A

ΔG = ΔH - TΔS

ΔG: gibbs free energy
ΔH: enthalpy
T: temp in Kelvin
ΔS: entropy

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7
Q

Lecture

Gibs free energy summary

A
  • ΔG is proportional to the negative of ΔSuniv
  • ΔG < 0 = spontaneous process
  • ΔG > 0 = non-spontaneous process
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8
Q

Lecture

Gibs free energy formula when
ΔH < 0 and ΔS > 0

A

ΔH (enthalpy) < 0: heat is emitted ↑ entropy
ΔS (entropy) > 0: entropy ↑

  • both are increasing entropy so universe ↑ entropy → spontaneous
  • both enthalpy and entropy favor the reaction
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9
Q

Lecture

Gibs free energy formula when
ΔH < 0 and ΔS < 0

A

enthalpy < 0: heat is emitted ↑ entropy
entropy < 0: entropy ↓
temp: low ↑ entropy

@ low temps, heat emitted = ↑ entropy
so universe ↑ entropy → spontaneous

@ hi temps, heat dispersed to warmer sorroundings makes change in ↑ of entropy small, insignificant

  • enthalpy favors the reaction
  • entropy opposes
  • “enthalpy driven” reaction
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10
Q

Lecture

Reactions will tend to go from _____ to _____ if they give _____ _____ OR if they involve a _____ of order (increase in _____). The actual direction depends on the balance of favorable and unfavorable contributions to both the enthalpy and entropy.

A
  • left
  • right
  • off heat
  • loss
  • randomness
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11
Q

Lecture

Gibs free energy formula when
ΔH > 0 and ΔS > 0

A

enthalpy > 0: heat is absorbed ↓ entropy
entropy > 0: entropy ↑
temp: high ↑ entropy

@ hi temps, heat emitted = ↑ entropy
so universe ↑ entropy → spontaneous

@ absorption of heat (by ΔH) from surrounding has less effect on entropy as temp increases

  • unusual, entropy driven
  • absorbs heat but is favored by the large entropy increase resulting from the formation of gaseous products
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12
Q

Lecture

Hydrogen bonds form between an ______ atom (the hydrogen _____) and a hydrogen atom covalently bonded to another ______ atom (hydrogen _____). common hydrogen acceptors are

A
  • electronegative
  • acceptor
  • electronegative
  • donor
  • oxygen, nitrogen, flourine
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13
Q

Hydrogen atoms covalently bonded to carbon atoms ______ participate in hydrogen bonding, because carbon is ______ more electronegative than hydrogen and thus the C—H bond is only very _____ _____

A
  • don’t
  • slightly
  • weakly polar
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14
Q

Lecture

Hydrogen bonds are strongest when the bonded molecules are oriented to maximize electrostatic interaction. That is when the H & the two atoms that share it are in a _____ _____

A

straight line

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15
Q

Lecture

Amphipathic compounds

A

compounds that contain regions that are polar (or charged) and
regions that are nonpolar

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16
Q

Lecture

When an amphipathic compound is mixed with water, the ____ (_____) region interacts favorably with the water and tends to dissolve, but the _____ (______) region tends to avoid contact with the water

A
  • polar
  • hydrophilic
  • nonpolar
  • hydrophobic
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17
Q

Lecture

hydrophobic effect

A

in amphipathic compounds the nonpolar regions of the molecules cluster together to present the smallest hydrophobic area to the aqueous solvent, and the polar regions are arranged to maximize their interaction with the solvent

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18
Q

Lecture

micelles

A
  • lipid molecules that arrange themselves in a spherical form in aqueous solutions
  • form in response to the amphipathic nature of fatty acids, meaning that they contain both hydrophilic regions (polar head groups) as well as hydrophobic regions (the long hydrophobic chain)
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19
Q

Lecture

The _____ ______ on interactions among lipids, and between lipids and proteins, is the most important determinant of structure in biological membranes

A

hydrophobic effect

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20
Q

Lecture

Water Forms Hydrogen Bonds with _____ Solutes

A

polar

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21
Q

Lecture

Entropy Increases as _____ Substances Dissolve

A

Crystalline (ie. NaCl)

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22
Q

_____ Gases Are Poorly Soluble in Water

A

Nonpolar

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23
Q

van der Waals Interactions Are Weak _____ Attractions

A

Interatomic

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24
Q

van der Waals interactions aka London forces

A
  • When two uncharged atoms are very close together
  • their surrounding electron clouds influence each other
  • variations in the positions of electrons create a transient electric dipole in one, and a opposite in the other
  • The two dipoles weakly attract each other, bringing the two nuclei closer
  • As the two nuclei draw closer together, their electron clouds begin to repel each other
  • where net attraction is maximal, the nuclei are said to be in van der Waals contact
  • Each atom has a their own van der Waals radius, a measure of how close that atom will allow another to approach
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25
Q

Weak Interactions Are Crucial to _____ _____and Function

A
  • Macromolecular Structure
  • function
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26
Q

We can calculate the stability of a noncovalent interaction, (i.e. hydrogen) from the _____ ______ which is

A
  • binding energy
  • the reduction in the energy of the system when binding occurs
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27
Q

the energy released when an enzyme binds noncovalently to its substrate is the main source of the enzyme’s _____ _____

A

catalytic power

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28
Q

One consequence of the large size of enzymes and receptors (relative to their substrates or ligands) is

A

that their extensive surfaces provide many opportunities for weak interactions

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29
Q

Lecture

For many proteins, tightly bound _____ molecules are essential to their function. These molecules, have distinctly different properties from those of the “bulk” _____ of the solvent and can be used in things like

A
  • water
  • water
  • “proton hopping” or form an essential part of a protein’s ligand-binding site
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30
Q

colligative properties

A
  • properties of solutions that depend on the ratio of the number of solute particles to the number of solvent molecules in a solution, and not on the nature of the chemical species present
    *
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31
Q

Colligative properties include

A
  • Relative lowering of vapour pressure
  • Elevation of boiling point
  • Depression of freezing point
  • Osmotic pressure
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32
Q

Osmosis

A

spontaneous net movement of solvent molecules through a selectively permeable membrane into a region of higher solute concentration, in the direction that tends to equalize the solute concentrations on the two sides

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33
Q

Osmotic pressure

A
  • pressure required to stop water from diffusing through a barrier by osmosis
  • represented by Π
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34
Q

Lecture

van’t Hoff equation

A
  • Π = icRT
  • Π = RT(i1c1 + i2c2 + i3c3 + ··· + incn)
  • R = gas constant: 8.315 J/mol • K
  • T = absolute temperature
  • i = van’t Hoff factor
  • c = solute’s molar concentration
  • ic = together is the osmolarity of the solution
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35
Q

Lecture

van’t Hoff factor

A
  • for an ionic compound is the number of ions that form when the compound dissociates.
  • for a non-ionizing molecule is 1.
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36
Q

Solutions of osmolarity equal to that of a cell’s cytosol are said to be _____ relative to that cell. Surrounded by an isotonic solution, a cell …

A
  • isotonic
  • neither gains nor loses water
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37
Q

Surrounded by an isotonic solution, a cell neither gains nor loses water

A

Surrounded by an isotonic solution, a cell neither gains nor loses water

38
Q

in a _____ solution, one with higher osmolarity than that of the cytosol, the cell _____ as water moves _____

A
  • hypertonic
  • shrinks
  • out
  • out
39
Q

In a _____ solution, one with a lower osmolarity than the cytosol, the cell _____ as water _____

A
  • hypotonic
  • swells
  • enters
40
Q

osmalarity

A

refers to the number of particles of solute per liter of solution

41
Q

In their natural environments, cells generally contain _____ concentrations of biomolecules and ions than their surroundings, so osmotic pressure tends to drive water _____ cells. Several mechanisms have evolved to prevent _____ of the cell.

A
  • higher
  • into
  • lysis
42
Q

Because the effect of solutes on osmolarity depends on the number of _____ _____, not their _____, macromolecules (proteins, nucleic acids, polysaccharides) have far ______ effect on the osmolarity of a solution than would an equal mass of their _____ components

A
  • dissolved particles
  • mass
  • less
  • monomeric
43
Q

Water molecules have a slight tendency to undergo reversible ionization to yield

A

a hydrogen ion (a proton) and a hydroxide ion

44
Q

H+, free protons do not exist in solution; hydrogen ions formed in water are immediately _____ to form _____ ______.

A
  • hydrated
  • hydronium ions (H3O+)
45
Q

_____ _____ between water molecules makes the hydration of dissociating protons virtually instantaneous

A

Hydrogen bonding

46
Q

Lecture

Proton Hopping

A
  • Short “hops” of protons between a series of hydrogen-bonded water molecules
  • results in extremely rapid net movement of a proton over a long distance
47
Q

Lecture

equilibrium constant

Keq or Kc
ᶜ = concentration

A
  • depicts [] of reactants/products in a reaction at equilibrium
  • measure of how far a reaction proceeds
  • gases and aqueous states only
  • changes only w/change in temperature
  • Given a reaction aA+bB ⇌ cC+dD the equilibrium constant Kc or Keq, is defined in terms of the concentrations of reactants (A and B) and products (C and D) at equilibrium
  • molarity is the unit of concentration
48
Q

Lecture

The equilibrium constant is _____ and ______ for any given chemical reaction at a specified temperature. It defines the composition of the final equilibrium mixture, regardless of the starting amounts of reactants and products

A
  • fixed
  • characteristic
49
Q

Lecture

Intermolecular forces

A
  • dipersion: Nonpolar solute and solvent
  • Dipole - Dipole: polar solute and solvent
  • Hydrogen bond: Solute and Solvent that has H Attached to F,O,N
  • Ion - Dipole: Polar solvent – Ionic solute
50
Q

Lecture

what does a large value of K mean

A
  • K >> 1
  • reaction lies far to the right at equilibrium
  • a high [] of products and a low [] of reactants
  • reaction goes to completion then reverses back to reach equilibrium
51
Q

Lecture

what does a low value of K mean

A
  • K << 1
  • reaction lies far to the left at equilibrium
  • reactant [] high & product [] low
52
Q

Lecture

equilibrium constant (Keq) for a reaction formula

A
  • coefficients = exponents in equilibrium constant
  • numerators: amt of products in equilibrium
  • denominator: amt of reactants in equilibrium
53
Q

give Keq for
Cu(s) + 2Ag(aq) ⇄ Cu(aq) + 2Ag(s)

A

[Cu²⁺] / [Ag⁺]²

54
Q

give Keq for
CaCO₃(s) ⇄ CaO(s) + 2CO₂

A

[CO₂]²

55
Q

give Keq for
2H₂O(l) ⇄ 2H₂ + O₂

A

[H₂]²[O₂]

56
Q

electrolyte

A

a substance, such as an acid, that forms ions in solution

57
Q

Lecture

a strong electrolyte, acid or base _____ dissociates into ions in solution

A

completely

58
Q

Lecture

a weak electrolyte, acid or base _____ dissociates into ions in solution

A

partially

59
Q

Lecture

water is _____, it can act as either an acid or base.

A

amphoteric

60
Q

autoionization

A

when a substance acts as an acid and a base w/itself

61
Q

autoionization reaction of water

A

H₂O(l) + H₂O(l) ⇄ H₃O⁺ + OH⁻

62
Q

Lecture

Kw

A
  • ion product constant for water
  • H₂O(l) + H₂O(l) ⇄ H₃O⁺ + OH⁻
  • Kw = [H₃O⁺] [OH⁻] or [H⁺][OH⁻]
  • Kw = 1.0 × 10⁻¹⁴
  • if [H⁺] ↓ then [OH⁻] must ↑ to = 1.0 × 10⁻¹⁴
63
Q

Lecture

acidic solution

A
  • contains an acid that creates more H₃O⁺
  • [H₃O⁺] > [OH⁻]
64
Q

Lecture

basic solution

A
  • contains a base that creates more OH⁻
  • [H₃O⁺] < [OH⁻]
65
Q

Lecture

in all aqueous solutions, _____ and _____ are present and their [] multiplied = _____.

A
  • H₃O⁺
  • OH⁻
  • 1.0 × 10⁻¹⁴
66
Q

Lecture

What is the concentration of H+ in a solution of 0.1 M NaOH?

A
  • Kw = [H⁺][OH⁻]
  • [H⁺] = Kw / [OH⁻]
  • [H⁺] = 1.0 × 10-14 M2 / 0.1 M
  • [H⁺] = 10-14 M2 / 10-1 M
  • [H⁺] = 10-13 M
67
Q

Lecture

What is the concentration of OH− in a solution with an H+ concentration of 1.3 × 10−4 M?

A

Kw = [H⁺][OH⁻]

[H⁺] = Kw / [OH⁻]

[H⁺] = 1.0 × 10-14 M2 / 1.3 × 10−4 M

[H⁺] = 10-14 M2 / 10-1 M

[H⁺] = 7.7 × 10-11 M

68
Q

Lecture

pH

A
  • the negative log of hydronium ion []:
    • pH = -log[H₃O⁺]
  • has 2 decimal places or as many sig figs from []
  • # to right of decimal are sig in a logarithm
69
Q

Lecture

find the pH of a solution w/
[H₃O⁺] = 1 × 1.0⁻³ M

A
  • pH = -log[H₃O⁺]
  • pH = -log[1 × 1.0⁻³]
  • pH = -(-3.00)
  • pH = 3.00
70
Q

Lecture

pH scale is a logarithmic scale. a change of 1 pH corresponds to a _____ change in [H₃O⁺]

A

tenfold

71
Q

calculate pH of [OH⁻] = 7.1 × 10⁻³ M

A
  • Kw = 1.0 × 10⁻¹⁴
  • Kw = [H₃O⁺] [OH⁻] or [H⁺][OH⁻] = 1.0 × 10-14
  • use 1st formula
  • [H₃O⁺] (7.1 × 10⁻³) = 1.0 × 10⁻¹⁴
  • [H₃O⁺] = 1.4 × 10⁻¹²
  • pH = -log(1.4 × 10⁻¹²)
  • pH = 11.85
72
Q

calculate [H₃O⁺] for a solution with a pH of 4.80

A
73
Q

Lecture

Acids may be defined as proton _____ and bases as proton _____. When a proton donor loses a proton, it becomes the corresponding proton _____. A proton donor and its corresponding proton acceptor make up a _____ _____ _____, related by the _____ reaction

A
  • donors
  • acceptors
  • acceptor
  • conjugate acid-base pair
  • reversible
74
Q

Lecture

pKₐ

A
  • another common p scale
  • pKₐ = - log Kₐ
  • the smaller the pKₐ, the stronger the acid
  • Keq = ([H+] [A<span>-</span>]) / [HA] = Ka
75
Q

Lecture

_____ is used to determine the amount of an acid in a given solution

A

titration

76
Q

Buffers

A
  • aqueous systems that tend to resist changes in pH when small amounts of acid (H+) or base (OH–) are added
  • consists of a weak acid (the proton donor) and its conjugate base (the proton acceptor)
77
Q

Henderson-Hasselbalch equation

A
  • pH = pKa - log ( [A-] / [HA] )
  • A-: conjugate base
  • HA: acid
  • allows us to
    • calculate pKa, given pH and the molar ratio of proton donor and acceptor
    • calculate pH, given pKa and the molar ratio of proton donor and acceptor
    • calculate the molar ratio of proton donor and acceptor, given pH and pKa
78
Q

Two especially important biological buffers are the _____ and _____ systems

A
  • phosphate
  • bicarbonate
79
Q

Enzymes typically show maximal catalytic activity at a characteristic pH, called the _____ _____. On either side of this optimum pH, catalytic activity often _____ _____

A
  • pH optimum
  • declines sharply
80
Q

The formation of ATP from ADP and inorganic phosphate is an example of a _____ reaction in which the elements of water are _____. The reverse of this reaction, _____ accompanied by the addition of the elements of _____ is a _____ reaction. _____ reactions are also responsible for the enzymatic depolymerization of proteins, carbohydrates, and nucleic acids, they are catalyzed by enzymes called _____, are almost invariably _____

A
  • condensation
  • eliminated
  • cleavage
  • water
  • hydrolysis reaction
  • hydrolases
  • exergonic
81
Q

Water is both the _____ in which metabolic reactions occur and a _____ in many biochemical processes, including hydrolysis, condensation, and oxidation-reduction reactions

A
  • solvent
  • reactant
82
Q

Lecture

how do you get the Ka value if you know the pKa

A
  • take the antilog of pKa
  • raise both sides of the equality to exponents of 10
  • pKa = -log Ka
    • ​pKa = 3.75
    • Ka = 10(-pKa)
    • Ka = 10-3.75
83
Q

Lecture

  • In pure water at 25°C, the concentration of water is _____ M:
    • grams of H2O in 1 L divided by its gram molecular weight: (1,000 g/L)/(18.015 g/mol)
A
  • 55.5
84
Q

from Ka to Henderson-Hasselbalch equation

A
85
Q

Lecture

Boltzmann’s constant

A
  • allows you to rewrite the ideal gas law formula to use # of molecules versus the # of moles
  • ideal gas law formula: PV = nRT
    • n: moles
    • R: gas constant
    • T: temp in Kelvin 8.31 J/mol K
  • deal gas law formula: PV = NRT
    • N: # of molecules
    • Kb: Boltzmann’s constant
      • = (gas constant / avogadro’s #)
      • 1.380649 × 10-23
    • T: temp in Kelvin
86
Q

Lecture

Coulomb’s law

A

a law stating that like charges repel and opposite charges attract, with a force proportional to the product of the charges and inversely proportional to the square of the distance between them

87
Q

Lecture

Coulomb’s law formula

A
  • q1 and q2 are the charges on particles 1 and 2
  • k is a constant of proportionality (Boltzmann again)
  • ε is the dielectric constant.
88
Q

Lecture

Energy of interaction

A
  • the contribution to the total energy that is caused by an interaction between the objects being considered
  • q1 and q2 are the charges on particles 1 and 2
  • k is a constant of proportionality (Boltzmann again)
  • ε is the dielectric constant.
89
Q

Lecture

Van der Waals interactions

  • Attraction: depends on _____ weak, but operates at relatively _____ distances
  • Repulsion: depends on _____ strong, but operates at relatively _____ distances
  • “6-12” potential (Lennard – Jones potential) connects them
A
  • 1/r6
  • long
  • 1/r12
  • short
90
Q

Lecture

H-bonds connect each molecule to an average of _____ nearest neighbors in liquid water and to _____ in ice.

A
  • 3.6
  • 4
91
Q

Lecture

volume “collapse” yields

A

smaller volume less motional freedom and thus lower entropy.

92
Q

Lecture

Thermodynamics

  1. q > 0
  2. q < 0
  3. w > 0
  4. w < 0
A
  1. Heat is absorbed; System’s energy increases
  2. Heat is released; System’s energy decreases
  3. System does work; System loses energy
  4. Surroundings do work on system; System gains energy

Biochemistry 403 (12 decks)

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