Water, Weak Acids and Bases Flashcards
Strong Bonds/Primary Bonds
Strong bonds are between ELEMENTS
Ionic
Covalent
Metallic
Ionic Bonds
TRANSFER of electrons
Strongest of the 3 primary bond types
Metal - Nonmetal
Covalent Bonds
SHARING of electrons
Second strongest primary bond
Nonmetal - Nonmetal
Metallic
SHARING of CONSTANTLY MOVING electrons
Weakest primary bond
Metal - Metal
Weak Bonds/Secondary Bonds (Strong–>Weak)
- Weak bonds are between MOLECULES*
1. Ion-Dipole
2. H Bonds
3. Dipole-Dipole
4. Ion-Induced Dipole
5. Dipole-Induced Dipole
6. Dispersion Forces
Hydrogen Bonds
When a Hydrogen is COVALENTLY BONDED to a F,N, or O it has the ABILITY to H-bond. Can only H-bond to F, N, or O (electronegative elements)
Weak Interactions among Biomolecules in Aq Solvent
Hydrogen Bonds- between neutral groups and peptide bonds
Ionic Interactions- Attraction and Repulsion
Hydrophobic Interactions
Van Der Waals- caused by delocalization of e- around 2 molecules
Water as a Liquid and Solid
Liquid: ONE molecule of water binds to ONLY 3.4 molecules of water.
-why? its liquid so its moving
Solid: ONE molecule of water binds to 4 molecules of water.
-why? Solid lattice structure
Ice is less dense than water because it expands
Hydrogen Bond Directionality
H-Bonds are stronger when they are straight (puts the positive charge of H directly between two negative charges)
H-Bonds are weaker when there is a bend in the alignment of X—H——X
Amphipathic Molecule
Has both Polar and Nonpolar regions
Water as a Solvent for Polar Molecules
Think Salt in Water
Ionic bonds (and other polar bonds) dissolve in water spontaneously because the entropy (S) is large due to the disordering of a highly structured lattice breaking down.
^G= ^H - T^S
^G is NEGATIVE
Water as a Solvent for Amphipathic/ Amphiphilic Compounds (Think Oil)
Non-Spontaneous Interaction because the non polar region of the molecule causes water to form highly organized “cages” around the polar tail. Organization=negative entropy (S)
^G= ^H - T^S
^G is POSITIVE
Water as a Solvent Overview
No matter what molecule you put in water, the molecule will be interrupting water-water hydrogen bonding (and the enthalpy is always +)
- Polar solvents compensate via INCREASE in entropy = negative deltaG = spontaneous
- Non-polar solvents cause a DECREASE in entropy because the water is being forced to organize around the non polar region = positive deltaG = non spontaneous
Lipids in Water
Clusters of lipid molecules are more energetically favorable than dispersed lipids because the organization of water is decreased when it doesn’t have to surround each individual lipid.
Micelles are even more favorable because only the polar regions interact with the water causing even less water organization.
Strong interactions NOT due to polar interactions with water, but because its the MOST THERMODYNAMICALLY STABLE
Lipid Bilayer
Two rows of lipid molecules with polar heads facing out and hydrophobic/nonpolar tails facing inward. Water inside and outside the bilayer interacting with the polar heads.
Micelles
Single row of lipid molecules in a sphere. Polar heads out lipid tails inside the ball. Used to deliver things into the body (drugs)
Liposome
Two rows of lipid molecules in a spherical shape where there is water both inside and outside the lipids. Similar to a bilayer but it forms a vesicle rather than a membrane.
Enzyme-Substrate Complexes in Water
Thermodynamically favorable because the binding of the Substrate to the Enzyme disrupts the order of the water molecules on the surfaces on the substrate and the enzyme. Entropy increases upon binding.
Solutes in Solution
Adding a Solute changes the colligative properties:
- Boiling Point
- Vapor Pressure
- Melting Point
- Osmotic Pressure
Isotonic Solution
No net water movement because solute levels are the same inside and outside the cell
Hypertonic Solution
More solute outside of the cell than inside
Causes the water to leave the cell in an attempt to equalize solute levels.
Water leaving the cell causes it to shrink/shrivel
Hypotonic Solution
Solute levels are higher inside the cell than outside.
Causes water to flow into the cell = swelling
pH and pOH formulas
pH= log (1/[H+]) or pH= -log(H+) --------> pH= log ( 1/ [ 1x 10^-7]) = 7
pOH is the same but with [OH-]
pH degree of change
pH scale is logarithmic. Therefore two solutions that differ by 1 pH have a difference of 10 times the amount of [H+]
Small changes of pH in the body can be fatal (blood pH of 7.2 is normal, 7.4 can be fatal)
Strong vs. Weak Acids and Bases
Strong Acids/Bases COMPLETELY dissociate in water
Weak Acids/Bases PARTIALLY dissociate in water
Titration
Add H+ to bases
Add OH- to acids
When pH=pKa there is an equal amount of acid and basin solution. Called the buffer region.
Strong Acid: pKa and Ka values
Large Ka
Small pKa
Most of the Acid has been converted into H+ and A-
Weak Acid: pKa and Ka
Small Ka
Large pKa
Less acid has dissociated = weaker acid
Henderson-Hasselbalch Equation
pH = pKa + log A-/HA
So when [A-] = [HA] you have log 1/1 = 0
so pH=pKa when you have equal concentrations of acid and base
(literally just explains titration curve)
Buffers
Minimize pH changes
Bicarbonate Buffer System
Carbonic acid and bicarbonate –> CO2
Most important buffer in the body because it keeps our blood pH constant
Acidosis
Too much acid in the arterial blood
pH lower than 7.35
Alkalosis
Too much base in the arterial blood
pH higher than 7.45
Respiratory Acidosis
Impairment of the disposal of CO2 by the lungs so pH levels are too low
Respiratory Alkalosis
Results from hyperventilation, excessive exercise etc
Breathe too fast and get rid of too much CO2, pH increases
Metabolic Acidosis
Over production of organic acids or inability of kidneys to excrete excess acid. pH is too low
Metabolic Alkalosis
Abnormal loss of acid from the body.
