Water and solutions. Prof. Ramune Morkuniene, Biochemistry Dept., LUHS

Transcription

1 Water and solutions Prof. Ramune Morkuniene, Biochemistry Dept., LUHS

2 Characteristics of water molecule Hydrophylic, hydrophobic and amphipatic compounds Types of real solutions Electrolytes and non- electrolytes Concentration of solution Osmolarity of real solutions Tonicity of solutions Dissociation of weak electrolytes. Principle of Le-Chatelier

3 Water Water plays a central role in the chemistry of all life = most abundant molecule in the cell (60-90% of total mass) - as a biological solvent - as a participant in many biochemical reactions - used for transport - blood, cerebrospinal fluid, lymph, urine - as an essential buffer to regulate temperature and ph

4 The Water Molecule is Polar Angle of o between two covalent bonds Polar covalent O-H bonds due to uneven distribution of charge (oxygen (δ - ), hydrogen (δ + )) Angled arrangement of polar bonds creates a permanent dipole for a water molecule

5 Hydrogen Bond in Water Water molecules attract each other due to their polarity As type of electrostatic interaction

6 The hydrogen bonds are weak 29 kj (7 kcal) are needed to break a hydrogen bond in water (the breaking of a covalent oxygen-hydrogen bond in the water molecule requires 450 kj (110 kcal)). Breaking the hydrogen bonds requires no more than heating the water to 100 C. Relative strengths of bonds in biological molecules.

7 Hydrogen Bond in Water: A water molecule can form up to four hydrogen bonds Hexagonal lattice structure of ice: The structure of liquid water is dynamic average life of hydrogen bond in water ~ s. The geometric regularity of these hydrogen bonds contributes to the strength of the ice crystal.

8 Hydrogen bonds can form between electronegative atoms and a hydrogen attached to another electronegative atom (O, N, F)

9 Ionic and Polar Substances Dissolve in Water (1) Polar and ionic substances readily dissolve in H 2 O - Hydrophilic (water-loving) Polar water molecules align themselves around ions or other polar molecules A molecule or ion surrounded by solvent molecules is solvated When the solvent is water the molecules or ions are hydrated

10 Ionic and Polar Substances Dissolve in Water (2) Ionic substances: Dipole-ion interaction Polar substances: Dipole-dipole interaction

11 Nonpolar Substances Are Insoluble in Water Hydrophobic (water-fearing) molecules are nonpolar Nonpolar compounds like fats (lipids, petrol, benzene) Little or no hydrogen bounding Hydrophobic substance

12 Amphipathic molecules Some materials have both polar and nonpolar ends - amphipatic. Amphipatic molecules tend to organize into micelle structures: the polar heads point towards the aqueous environment and nonpolar tails are inside. Detergents (surfactants) are examples

13 Detergents in water Polar (hydrophylic) end Soap nonpolar parts make multiple weak interaction with greasy contamination and pull them apart from skin. Nonpolar (hydrophobic) end

14 Solubilities of molecules in water Solubility in water depends upon the ratio of polar to nonpolar groups in a molecule The larger the portion of nonpolar groups the less soluble the molecule is in water The larger the portion of polar groups (e.g. hydroxyl groups (-OH)) the more soluble the molecule is in water

15 Solubilities of short-chain alcohols in water

16 Problem Based on the structure, do you expect vitamin C to be hydrophobic or hydrophylic?

17 Solutions Are homogenous mixtures of two or more substances Consist of a solvent and one or more solutes. Solute: the substance present in lesser amount. Solvent: the substance present in greater amount. The solute and solvent in a solution can be a solid, liquid and/or a gas. In relation to the solvent, solutions can be of three types: 1. Liquid solutions 2. Solid solution 3. Gaseous solutions

18 Liquid solution Similar size of solute and solvent particles Solute molecules must be able to interact with solvent molecules through noncovalent interaction

19 Like dissolves like A solution forms when there is an attraction between the particles of the solute and solvent. A polar solvent such as water dissolves polar solutes such as sugar and ionic solutes such as NaCl. A nonpolar solvent such as hexane (C 6 H 14 ) dissolves nonpolar solutes such as oil

20 Properties of real solutions A mixture of two or more components is homogenous The dissollved solute is molecular or ionic in size The solute remains uniformly distributed throughout the solution and will not settle out with time It is usually transparent (colored or colorless)

21 Formation of a solution (1): Ionic solution Solution, in which the solute is in dissociated state - Ionic solution Polar H 2 O molecules attract Na + & Cl -. In solution, Na + & Cl - are surrounded by H 2 O molecules (Hydration).

22 Formation of a solution (2): Molecular solution Solution, in which the solute is in non-dissociated state (as molecules) - Molecular solution

23 Solubility rules Insoluble salts attraction between its positive and negative ions are too strong for the polar molecules to break.

24 Electrolytes and nonelectrolytes Substances that release ions are called electrolytes (whose aqueous solutions are conductors of electricity) Electrolytes are capable to release ions in solution Strong electrolytes completely ionized Weak electrolytes partially ionized Nonelectrolytes produce molecules

25 Equivalents An equivalent - the amount of a electrolyte that carries one mole of positive or negative charge

26 Calculating Equivalents Example: How many equivalents are in 6.0 g of Ca 2+? Solution: There are 2 equivalents in 1 mole of calcium ion. Molecular mass of calcium g. 6.0g Ca 2+ 1mole Ca 40.1g Ca Eq Ca 1mole Ca = 0.30 Eq Ca 2+

27 Electrolytes in body fluids are expressed in milliequivalents (meq/l) 1 Eq = 1000 meq Typical concentrations of electrolytes in blood plasma Electrolyte concentrations in intravenous replacement solutions

28 Factors affecting solution formation 1. Temperature: 1. For most solids dissolved in a liquid, an increase in temperature results in increased solubility 2. For a gas in water the solubility decreases with increasing the temperature 2. Pressure has a marked effect on solubility of gases in liquids. Henry s law states that the solubility of gas in liquid is directly related to the pressure of that gas above the liquid. Pressure leads to Solubility of gases A hyperbaric (high pressure) chamber high pressure moves more O2 into the blood Divers Slow decompression

29 Concentration Concentrat ion of a solution = amount of solute amount of solution 1. Percent concentration: mass % (m/m) mass of solute (g) 100% mass of solution (g) = =weight/weight (w/w) volume of solute (ml) volume % (v/v) = 100% volume of solution (ml) mass of solute (g) mass/volume % (m/v) = 100% volume of solution (ml)

30 Calculating percent concentrations Example 1: What is the weight percent of a solution prepared by dissolving 30.0 g NaOH in g of H 2 O? Solution: We find the mass of the solution by adding the mass of the solute and solvent: 30.0 g NaOH g H 2 O = g of solution 30.0 g NaOH weight % = 100% = g 20.0%(w/w)

31 2. Molar concentration: Molarity(M ) = moles of liters of solute (mol) solution(l) 1 mole - a standart unit that contains atoms Avogadro s number 1 mole = molar mass of compound (in grams)

32 Calculating molar concentration Example : What is the molarity (M) of 60.0 g NaOH in L of solution? Solution: We convert grams of NaOH to moles of NaOH using the molar mass NaOH (40.0): 1mole NaOH 60.0 g NaOH = 1.50 moles NaOH 40.0 g NaOH Grams of NaOH Molar mass Moles of solute 1.50 moles NaOH M = = L Liters of solution 6.00 M (or mol/l) solution

33 Diluting a solution In a dilution, an amount of solute does not change A dilution increases the volume of the solution As a result, the concentration of the solution decreases The dilution calculation can be expressed as an equation: c 1 V 1 = c 2 V 2 c concentration V - volume Example: How many milliliters of 6 M NaOH solution are needed to prepare 1 L of a 0.15 M NaOH solution? Solution: c 2V c 1 = 6 M; V 1 = x; c 2 = 0.15 M; V 2 = 1 L; V = = L = L c 6 1 = 1 25 ml

34 Effect of solutes on water properties Dissolved solutes change structure and properties of H 2 O. The properties of solution which depend on the number of solute molecules, not on the chemical nature of solute are called colligative properties: lowered vapour pressure; increased boiling point; decreased melting point; osmotic pressure

35 Osmosis In osmosis, the solvent (H 2 O) moves through a semipermeable membrane from the solution that has a lower concentration of solute into the solution where the solute concentration is higher. Osmotic pressure is the pressure that prevents the flow of additional water into the more concentrated solution (force required to counterbalance the force of osmotic flow (H 2 O flow) through a semipermeable membrane)

36 Osmotic pressure Osmotic pressure is directly proportional to the concentration of a solution. This is Van t Hoff s rule: Π = c RT c molarity of a solution, R the gas constant, = atm/mol/k T the absolute temperature in Kelvin degrees.

37 Osmolarity Osmolarity is a characteristic of osmotic pressure, which shows the concentration of a solution expressed in the number of osmoles. Osmolarity = Molarity x Number of particles of solute Osmole - a unit indicating how many particles of solute are in a solution. Example NaCl (s) Na + (aq) + Cl - (aq) Osmolarity of 1M NaCl =1M x 2 =2 osmoles (Osm) Na 2 SO 4 2Na + + SO 4 2- = 3 Osm

38 Tonicity of solutions: isotonic, hypotonic and hypertonic Solutions with the same osmotic pressure are called isotonic. Isotonic solution - solute concentration the same as in cell fluid - cell shape maintained. Hypertonic solution - solute concentration higher than in cell fluid - water leaves cell, causing it to collapse. Hypotonic solution - solute concentration lower than in cell fluid - water enters cell, causing cell to expand and burst.

39 Red blood cells Normal plasma osmolarity 0.3 Osm 0.9% NaCl

40 Dissociation Dissociation is a general process in which ionic compounds (complexes, molecules, or salts) split into smaller molecules or ions. The dissociation constant K d is the ratio of dissociated to undissociated compound [K + ][A - ] KA K + + A - K c d = [KA]

41 Dissociation of electrolytes - ionisation Ionization - electrolytes ability to form ions in water solutions a concentration of ionized molecules Ionization degree = α = concentration of solution c Strong electrolytes are those, whose α > 30%, e.g. almost all mineral salts, HNO 3, H 2 SO 4, HCl, KOH, NaOH and etc. Their dissociation is irreversible. NaCl Na + + Cl - Mild electrolytes are those, whose α = 3 30%, e.g. H 3 PO 4, H 2 SO 3. Weak electrolytes are those, whose α = 0 3%, e.g. organic acids, NH 4 OH, H 2 CO 3 and etc.

42 Dissociation of weak electrolytes It is reversible: KA K + + A - Dissociation (ionization) constant: [K + ][A - ] K c = [KA] K c =const.; Example: dissociation of acetic acid [CH CH 3 COOH CH 3 COO - + H + 3 COO - ][H + ] K c = [CH 3 COOH] Ions of the same type influence equilibrium - decrease dissociation of weak electrolyte.

43 Le Chatelier s principle The system at equilibrium resists to outside changes. At very beginning, the equilibrium is perturbed, but then it is regained by shifting the equilibrium in the direction, which is opposite to the change done.