Chapter 6 Chemistry of Water; Chemistry in Water

Transcription

1 Chapter 6 Chemistry of Water; Chemistry in Water Water is one of the most remarkable and important of all chemical species. We, and all living things, are mostly water about 80% of our brain; 65% of our muscle. Even bone has water. Not surprisingly, the chemistry of water and the chemistry that can occur in water is a critical component of Chemistry. Most of the reactions that we will discuss in the course of first year take place in aqueous solution. Chapter 6 1

2 The Density of Water: Water has a putative density of 1.0 gram/cubic centimeter or 1.0 gram/ml. In fact, it is a little more subtle than that because the density of water depends upon temperature. The discontinuity in density at 0ºC is unusual water is the only liquid that does this and is responsible for ice floating. Ice is about 10% less dense than water at the same temperature. Chapter 6 2

3 Specific Heat Capacity: The amount of heat or energy required to change the temperature of one gram of a substance by 1ºC or 1 K is the specific heat capacity and varies from substance to substance. In comparison to other liquids, water has a very high specific heat capacity indeed, compared to most other substances. This has interesting consequences. Chapter 6 3

4 Enthalpy Change of Vapourization: Liquids (and to some extent solids) establish an equilibrium with their surroundings, with the molecules on or near the surface changing state. In an open environment, the material will continue to evaporate until it is entirely gone. Water evaporates very slowly or requires much more energy than other liquids for vapourization. Chapter 6 4

5 Equilibrium Vapour Pressure: In a closed vessel (or environment), a liquid will establish an equilibrium with the vapour above it. For water: H 2 O (l) H 2 O (g) This equilibrium depends on the energy of vapourization and condensation which are one and the same. They are simply different directions. The consequence is that energy released during condensation vapourizes an almost equivalent amount of water. Chapter 6 5

6 As a consequence of the relatively high heat of vapourization, which is a consequence of the hydrogen bonding structure in liquid water, water has a relatively low equilibrium vapour pressure the pressure that the liquid will exert in a container. Water is much less volatile than substances like diethyl ether and acetone. Put another way, if containers of water and acetone are left open, the acetone will disappear more quickly. Gasoline, for example, is very volatile which is why it smells so strongly. Chapter 6 6

7 Boiling point: If we heat a liquid, eventually the kinetic energy of the molecules will overcome the forces holding the molecules together in the liquid. This can be seen by an increase in vapour pressure with temperature. At some temperature, the vapour pressure from the liquid will equal that of the atmosphere and the liquid will start to boil. This is the normal boiling point and is defined for 1.0 atm of pressure. Note that this implies that boiling points will change if atmospheric pressure changes. This is why water boils at a much lower temperature at altitude and food takes longer to cook. Water (b.pt.) 100ºC Methane (b.pt.) ºC Chapter 6 7

8 Surface Tension: Liquids have both adhesive and cohesive forces at work. Cohesive forces pull the molecules together pull the molecules towards each other. The result is that most liquids will maximize their volume to surface ratio. They will adopt a spherical geometry if unconstrained as this provides the most volume for the least amount of surface. This is true of water droplets, although gravity and air resistance do play a part. Chapter 6 8

9 So where do these properties of water come from? What is our model? Our model of water depends on the intermolecular forces within a bulk sample. Intermolecular forces are the forces of attraction between molecular (or chemical) species. These differ from covalent interactions ( intramolecular forces ) which account for the structures of different chemical species. Intermolecular forces lead to the physical properties of chemical species. The forces between molecules arise from polarity either permanent or momentary. Chapter 6 9

10 A dipole moment leads to a molecule aligning itself with an electric field, where the positive pole of the molecule is attracted to the negative pole of the field and vice versa. Molecules develop dipoles because of a difference in electronegativity between the atoms in a covalent bond. Note that in a perfect covalent bond (homo-atomic species such as H 2 or Cl 2 ), the atoms each have equal claim to the pair of electrons joining the atoms. The electrons are shared 50:50. However, in the case of hetero-atomic species, such as HCl, the pair of electrons are not shared equally. They spend more time with one atom than the other. The consequence is that one end of the bond has more electrons and is negative. The other end is positive. Chapter 6 10

11 There are a number of systems for measuring electronegativity the relative ability of each element to draw electrons to itself in a bond. They all result in some form of the periodic table with assigned electronegativity values. Chapter 6 11

12 The difference in electronegativity, along with some structural considerations, tells us the bond polarity and the direction/size of the dipole moment within a molecule: Chapter 6 12

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14 There are many tables of dipole moments calculated for a variety of molecules and molecular geometries. This is not something that we have to do in this course (but it is not that hard.). Chapter 6 14

15 Another way of representing a dipole is with an electrostatic potential map which is a space filling diagram, where red represents the region of the molecule with the greatest negative charge and blue represents the positive regions. For water, this looks like: Chapter 6 15

16 Dipole-dipole interactions result in attraction between adjacent molecules. In the case of water, the dipoledipole interaction is particularly strong, due to the presence of hydrogen bonding. (This is an extreme form of dipole interaction where one of the components is a Hydrogen atom and the other is an Oxygen, Nitrogen, or Fluorine atom.). Chapter 6 16

17 The force of attraction resulting from hydrogen bonding in water means that the molecules are more tightly held than would be expected when compared to similar compounds: Chapter 6 17

18 Even if a molecule doesn t have a dipole, there are still forces of attraction between molecules. These are collectively referred to as dispersion forces, such as instantaneous dipole-dipole interactions. Venn Diagram Dispersion forces are important in explaining the physical properties of non-polar substances. Generally speaking, the larger an atom is (the more electrons it has), the more easily its electrons can be shifted and the more polarizable it is. Chapter 6 18

19 Dispersion forces depend on the instantaneous development of a dipole which has a rippling effect throughout the remaining molecules of the liquid: Chapter 6 19

20 Our model of water H 2 O is a chemical species with two O-H covalent bonds that is capable of forming a hydrogen bonding network. How does this explain the physical properties that we have observed for water? The physical properties of molecular compounds depend on the strengths of the intermolecular forces. The ways in which water is remarkable, compared to other substances with molecules of similar size and/or mass, can be attributed to strong hydrogen bonding between water molecules. Chapter 6 20

21 Density: The open structure of ice means fewer molecules (less mass) in any given volume. Chapter 6 21

22 Heat Capacity: The inherent energy tied up in the hydrogen bonding network means that a lot more energy is required to get the molecules moving freely and to break down the intermolecular interactions. Vapourization: The stronger the intermolecular interaction, the more energy needed to separate the molecules. Same reasoning for Equilibrium Vapour Pressure. Chapter 6 22

23 Surface Tension: Molecules in the middle of the liquid are surrounded on all sides by equal attractive forces but at the surface, the net attractive force is into the bulk of solution. Chapter 6 23

24 A solution arises when a solute is dissolved in a solvent. A solvent is the chemical species that dissolves the solute to give a solution. A solute dissolves in a solvent to give a solution. The definitions are a bit circular or self-reliant, but a solvent can be any number of chemical compounds or species such as gasoline, iron, nitrogen gas, or water. And it is aqueous solutions ones in which water is the solvent - that we are most interested in because they permeate chemistry, biochemistry, and biology. Chapter 6 24

25 One of the general principles of solubility can be summed up with the expression: like dissolves like Non-polar substances tend to favour dissolving in non-polar solvents. Polar substances tend to favour dissolving in polar solvents. And charged compounds tend to favour dissolving in water. Chapter 6 25

26 A charged atom dissolved in water is an example of an ion-dipole interaction, where the dissolution is called aquation or hydration. This is what the (aq) means that the ions are interacting with water molecules to give ion-dipole complexes in which the ion is surrounded by water molecules. Compounds which dissolve to conduct electricity are called electrolytes and ionic species are invariably electrolytes. Substances that dissociate completely are called strong electrolytes (i.e. NaCl). Other compounds, that don t dissolve as readily or dissociate completely are weak electrolytes. Chapter 6 26

27 Not all ionic compounds are soluble. Some give ionic precipitates, with a general guide: Chapter 6 27

28 Organic compounds can also be polar (i.e. CH 3 CH 2 OH) and miscible but they can also be non polar (i.e. CH 3 CH 2 CH 2 CH 2 CH 2 CH 2 CH 2 CH 2 CH 2 CH 3 ) and immiscible. The question of solubility depends on the relative strengths of the interactions solute/solute, solvent/solvent, and solvent/solute. Chapter 6 28

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30 One of the more unusual properties of water is that it undergoes self-ionization (aka auto-ionization ): H 2 O (l) + H 2 O (l) H 3 O + (aq) + OH - (aq) The species H 3 O + is called the hydronium ion and is a common form for writing H + but, in fact, the ionic species for H + in solution is H 9 O 4 + (aq) : Chapter 6 30

31 Since water is ubiquitous, much chemistry is done in aqueous solution. The reactions that occur in water can be loosely categorized in the following types of reactions: Some of these we will address in more detail in later chapters. And some reactions are a combination of types (i.e a redox reaction might also result in a precipitation). Chapter 6 31

32 Precipitation reactions result when the products exceed their solubility, resulting in one or all of the products forming solids. For example: NaCl (aq) + AgNO 3(aq) AgCl (s) + NaNO 3(aq) Note that the reaction can be better described as: Cl - (aq) + Ag + (aq) AgCl (s) Both Na + (aq) and NO 3 - (aq) are spectator ions as they do not participate in the process they just watch. Chapter 6 32

33 A oxidation-reduction reaction or redox reaction result from the transfer of electrons from one chemical species to another. This means that while one species is losing electrons, the other is gaining and that the two reactions reduction and oxidation must occur together. Oxidation is the loss of electrons. (LEO or OIL) Reduction is the gain of electrons. (GER or RIG) Chapter 6 33

34 Redox reactions can be analyzed by separating the reaction into two halves, called half-cell equations. Balancing a redox reaction is simply a matter of identifying the two half-cells, and then finding a common denominator for the electrons: O 2(g) + 4H + (aq) + 4e - 2H 2 O (aq) 2H 2(g) 4H + (aq) + 4e - O 2(g) + 2H 2(g) 2H 2 O (aq) Chapter 6 34

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36 Acid-Base reactions combine an acid (a proton donor) with a base (a proton acceptor) to give a salt (a neutral ionic species) and water. NaOH (aq) + CH 3 COOH (aq) Na(CH 3 COO) (aq) + H 2 O In one sense, all reactions can either be viewed as acidbase (2 electrons) or redox (1 electron) reactions. Chapter 6 36

37 Chapter 6 37

38 Complexation is an example of an acid-base reaction in which the two species are a Lewis Acid and a Lewis Base. The Lewis definition says that an acid is a species capable of accepting a pair of electrons while a base is a species capable of donating a pair of electrons. These are critically important reactions in biochemistry the complexation of oxygen by the iron atoms in hemoglobin is what allows us to transport oxygen throughout our bodies. Chapter 6 38

39 Simple solvation or aquation of metal ions is an example of complexation. Chapter 6 39

40 The question of whether or not something dissolves is qualitative. Sometimes we would like to be quantitative to know how much is in solution. To do this, we define a variety of concentration terms but the most common and the one that we will use heavily is molarity. molarity = moles of solute / volume of solvent c = mol / L Note that because chemical species don t always completely dissociate, the concentration of the individual species may vary from the overall concentration of the compound (i.e. acetic acid). Chapter 6 40