Biochemistry: Biochemistry Fall 2022- Water Flashcards


Set Details Share
created 12 days ago by Cahickey77
9 views
book cover
Biochemistry
Chapter 3
Water
updated 10 days ago by Cahickey77
Subjects:
biochemistry
show moreless
Page to share:
Embed this setcancel
COPY
code changes based on your size selection
Size:
X
Show:

1

Molecular structure of water

  • Two atoms of hydrogen
  • One atom of Oxygen
  • Tetrahedral geometry (can form hydrogen bonds with four other water molecules)
  • Bent geometry
  • Hydrogens are bound to oxygen by single covalent bonds

2

Molecular Structure of water (Charges)

  • Oxygen is partially negative
  • the two hydrogen atoms are partially positive
  • the bond is polar
  • Because the charge is separated there forms dipoles

3

Hydrogen Bond with respect to two waters

A hydrogen bond results when the two electronegative oxygen atoms of two water molecules compete for the same electron deficient hydrogen atom. The hydrogen bond is represented by short parallel lines designating the weak covalent character and directionality of the bond

4

Electrostatic Interactions

occur between any two opposite partial charges (polar molecules) or full charges (ions or charged molecules)

5

Covalent bonds

Involve electron sharing with orbital overlap or mixing. Covalent character confers directionality to the bond or interaction m unlike the uniformly spherical force field around an ion

6

Most important Non-covalent interactions

ionic interactions

Hydrogen bonds

van der Waals interactions

7

Liquid Water

  • Structure is irregular
  • On average there are 2-3 H-bonds per water molecule that are breaking and reforming
  • Weak hydrogen bonds

8

Ice

  • 4 hydrogen bonds per water molecule
  • hydrogen bonds are directional, straight, and stable
  • less dense than the liquid form

9

Specific heat

the amount of heat that must be absorbed or lost for one gram of a substance to change its temperature by 1 degree Celsius

10

Density

Mass per volume. The more dense a substance the more mass it has per volume

11

When is water the most dense?

At 4 degrees Celsius

Once water reaches this and goes to 0 degrees C it expands (ie, develops a lower density)

12

Types of interactions with water

Hydrophilic interactions

Hydrophobic interactions

13

Hydrophilic Interactions

The polar nature of water makes water an excellent solvent for

  • Polar Solutes (H-bonds) (which are like water)
  • Ionic Materials (solvation Spheres) (a hydration shell is what water forms around the dissolved ionic compounds in water)

14

What do hydration shells do for the water molecules?

Ions/ionic compounds, in the presence of water, form hydration shells that create local order in the water molecules

15

Hydration Shell

A shell of water that surrounds ions when in water

For example, when NaCl is dissolved water molecules will surround Na and another group will surround Cl

16

Water's dielectric constant

Is high and decreases the force holding ionic interactions between molecules

17

Coulombs law of forces

F=(q1q2)/(4(r^2)D)

F-force holding the two charges (force of ionic interaction

D-Dielectric constant (a property of solvent where the charges exist in

r-distance between two charges

q-charge

18

Ions in water

Always hydrated and stay separate from one another because of their hydration shell

19

Hydrophobic Interactions

  • Affinity of non-polar compounds for another
  • Initially the hydrogen bonding of water reorganizes to accommodate the non-polar solute forming a local cage like structure (clathrate)
  • In the process, there is an increase in the "order" of water around the non-polar solute resulting in a decrease in entropy (deltaS) or -delta S during clathrate formation
  • Reactions are spontaneous
  • Oil in water is an example

20

Hydrophobic Interactions continued

  • Minimize solvation cages, hence, result in an increase of entropy (delta S>0, becomes positive) because fewer ordered water molecules are formed around the solute
  • promote an increase in enthalpy (delaH=positive: endothermic rxn)

21

Amphiphile

A chemical compound possessing both hydrophilic (water loving, polar) and lipophilic (fat-loving) properties such a compound is called amphiphilic or amphipathic

  • Heads toward water and tails away from water (toward molecule)

22

Van der Waals

can be dipole-dipole (strongest weak interaction), dipole-induced dipole, induced dipole-induced dipole (Londan dispersion which are the weakest)

23

Type of bonds

Covalent - sharing of electrons

Noncovalent

  • Ionic interaction (electrostatic interaction) (salt bridges)
  • Hydrophobic interactions
  • Hydrogen bond
  • van der Waals forces

24

Colligative Properties

Physical properties not affected by the specific structure of dissolved solutes but by the number of solute particles per unit volume of solvent

25

Colligative properties include

  • Freezing point depression
  • Boiling point elevation
  • Lowering of vapor pressure
  • Osmotic pressure

26

Osmosis

  • A spontaneous process whereby solvent molecules (e.g. water) passes through a semi permeable membrane from a solution of lower solute concentration to a solution of higher solute concentration
  • Pores in the membrane allow solvent molecules to pass through in both directions, but are narrow for larger solute molecules or ions to pass
  • The side with more particles would draw the water over

27

Osmotic Pressure

The pressure that would have to be applied to a pure solvent to prevent it from passing into a given solution by osmosis and is based on the number of particles

Pi=iMRT

28

Pi=iMRT

i=vanHoff

M=Molarity (mol/L)

R= constant (0.082 L*atm/Kmol)

T= Temperature in Kelvin

29

Hypertonic

Higher solute concentration in solution than in cells

30

Isotonic

Equal solute concentrations in solution and in cells

31

Hypotonic

Higher solute concentration in cells than outside of cell

32

Dissociation of Water Molecules

Sometimes a hydrogen atom shared by two waters shifts from one molecule to the other

The hydrogen atom leaves its electron behind and is transferred as a single proton - a hydrogen ion (H+)

The water molecule that lost a proton is now a hydroxide ion (OH-)

The water molecule with the extra proton is a hydronium ion (H3O+)

33

Bronsted Concept of Acids and Bases

An acid is a substance capable of donating protons

A base is a substance capable of accepting protons

There are conjugate acid base pairs

Strong acids and bases dissociate completely in water

34

Weak acids

Partially dissociate in aqueous solution, the equilibrium lies well to the left

Has the following acid dissociation constant:

Ka=[H+][A-]/[HA]

35

Weak Acids and pKa

  • The strength of an acid can be determined by its dissociation constant, Ka
  • When Ka<1, HA is not significantly dissociated and [HA]>[H+] and [A-]. In this case, HA is the weak acid
  • The larger the Ka the stronger the acid, conversely, the smaller the Ka the weaker the acid
  • the value of Ka can be represented as pKa

pKa=-log10(Ka)........10^(-pKa)=Ka

  • The larger the pKa, the weaker the acid
  • At a constant temperature, pKa is a constant for a conjugate acid and its conjugate base pair (ie. Ka of an acid changes with temperature)

36

Buffers

Buffer solutions are made up of a mixture of two substances, a conjugate acid and it conjugate base

Buffers can be divided according to their chemical nature into two types:

  • Acidic buffers
  • Basic Buffers

37

Acidic Buffers

Contain a weak acid and its salt which is a strong base

CH3COOH ----> CH3COO-+ H+

38

Basic Buffers

Contain a weak base and its salt which is a strong acid

NH3----> NH4++OH-

39

Buffer Solutions

  • Buffer solutions resist pH changes when H+ or OH- is added (Le Chatelier's principle
  • Various buffers have different pH ranges where they can prove useful
  • If you need to select a buffer around a certain pH range, perform the selection based on the pKa of the buffer
  • A good buffering range occurs at one pH unit above and below the pKa of a buffer

40

Normal pH for Human blood

7.4 with a range of 7.35-7.45

41

Acidosis

a condition when the blood drops below 7.35

This condition occurs in diabetes and kidney disease and can result in coma and death

42

Alkalosis

When the pH rises above 7.45

If not corrected convulsions and respiratory arrest occurs

43

Some other things Buffers do

  • Buffer prevents a rapid change in pH
  • If you add acid to a buffer, the buffer absorbs the h+, preventing the pH from falling fast
  • various buffers have different pH ranges where they can prove useful

44

Henderson-Hasselbach's Equation

pH=pKa+log([A-]/[HA])

When [A]=[HA] then pH=pKa

Note: Buffers are most effective when [A-]/[HA]=1 or in the pH range of pKa+- 1

45

Buffer Capacity

The capacity of a buffer to maintain a specific pH depends on two factors:

  1. Total buffer concentration (TBC) (TBC=[HA]+[A-] because the more H+ or OH- can be absorbed without changing the pH
  2. Ration of [A-]/[HA] for buffers consisting of a weak acid and its conjugate base

46

Buffer Concentration

Most buffers include weak acids and their conjugate bases

For these buffers, their total concentration is: SUM of ([HA] weak acid + [A-] conjugate base)

47

Physiological Buffers

Three most important buffers in the body are:

  1. The Phosphate buffer (higher concentration in cells versus blood)
  2. The bicarbonate buffer (higher concentration in blood versus cells)
  3. The protein buffers, especially hemoglobin and albumin in blood

48

Bicarbonate Buffer

Most significant buffer compound in human blood

Buffering capacity of blood relies on two equilibria