Chapter 5: Water for Life

1. Chapter 5: Water for Life

2. “Water has never lost its mystery. After at least two and a half millennia of philosophical and scientific inquiry, the most vital of the world’s substances remains surrounded by deep uncertainties. Without too much poetic license, we can reduce these questions to a single bare essential: What exactly is water?” Philip Ball, in Life’s Matrix: A Biography of Water, University of California Press, Berkeley, CA, 2001, p. 115 Do you know where your drinking water comes from? Do you know if your drinking water is safe to drink? How would you know?

3. Different Representations of Water Lewis structures Space-filling 5.1

4. Water is a very unique molecule. • It has a very high boiling point for such a small molecule • It is an excellent solvent for many types of compounds • The solid form is less dense than the liquid (a very rare property) • Has a very high heat capacity

5. The properties of water are due to: • It's molecular geometry • It's small size • And the type of bonds it contains

6. Bond – an attractive force that holds two atoms together. Atoms bond to obtain a more stable electronic configuration. They do this by gaining, losing, or sharing electrons with other atoms

7. Covalent (molecular) bonds Between non-metals and non-metals Atoms tied together by sharing electrons Forms molecules – fixed numbers of atoms in a particular geometry Ionic bonds Between metals and non-metals Attraction between oppositely charged atoms/molecules (ions) Formula shows the ratio of ions

10. Metallic bonds Between two or more metals Outer electrons are not linked to a particular atom, as in covalent and ionic bonding These electrons are shared between all atoms in a 'sea of electrons' This is why metals conduct electricity and heat so well

11. For Covalent and Ionic Bonding: • The electronegativity determines what type of bond will form between two atoms • Electronegativity (EN)– attraction for shared electrons • Fluorine is the most EN element • Francium the least EN element • Larger EN – stronger attraction for bonding electrons

13. There are two types of covalent bonds: Non-polar (covalent) Bonding electrons are equally shared When a non-metal bonds to itself, or to another non-metal with a similar EN Polar (covalent) Bonding electrons are shared, but not equally One atom has a larger EN than the other Covalent Bonds

16. Polar (covalent) Usually between non-metals two or more spaces apart on the periodic table The atom with greater EN (closer to F) pulls harder on the shared electrons This pull creates a polarity, or dipole, across the bond: The atom with the higher EN has a slight negative charge The other has a slight positive charge Covalent Bonds

17. A difference in the electronegativities of the atoms in a bond creates a polar bond. O H H A polar covalent bond is a covalent bond in which the electrons are not equally shared, but rather displaced toward the more electronegative atom. Partial charges result from bond polarization. 5.1

18. Polarity of hydrogen covalent bonds is difficult to tell from the PT: H-C – non-polar (some books list as very slightly polar) H-O – very polar H-N – very polar H-F – very polar

19. If the electronegativity difference is large enough between two atoms, they will not share the 'bonding' electrons: The atom with the greater EN takes the 'bonding' electrons from the other atom This atom then has a negative charge The atom that lost the electron(s) has a positive charge Ionic Bonds

20. Ions • Cations – lost one or more electrons • positively charged • Anions – gained one or more electrons • negatively charged • When forming ions, atoms usually want to get to the same number of electrons as the nearest Noble Gas

23. Naming Ions: ●Cation (metal) – name is the same as the element, + 'ion' ●Fixed charge cations – metals that only form one cation (such as Group 1 and 2 metals): Li+1 → lithium ion, Ca+2 → calcium ion ●Variable charged cations – metals that may form different cations (most transition metals)Use Roman numerals to show the charge: Fe+2 → iron (II) ion Fe+3 → iron (III) ion

27. ●Anion (non-metal) – use the root of the element name, change the ending to 'ide', + 'ion': S → S-2 sulfur → sulfide ion N → N-3 nitrogen → nitride ion O → O-2 oxygen → oxide ion

29. Naming Binary Ionic Compounds: ●List the cation first, then the anion ●Do not include 'ion' in the name ●Names must be distinctive, in order to distinguish between similar compounds, such as with variable-charged metals NaCl – sodium chloride CaF2 – calcium fluoride FeI2 – iron (II) iodide FeI3 – iron (III) iodide

30. Writing formulas for binary ionic compounds: ●The formula shows a ratio of one ion to the other. ●The ionic charges must cancel out so that the overall charge is neutral ●Always list the metal first, then the non-metal ●Select subscripts to balance charges ●Reduce subscripts if needed to obtain the lowest whole number ratio between ions

31. To determine the charge on a variable charge cation, treat the formula as an algebraic expression:To determine the iron charge in Fe2O3 ●let Fe = x and O = y (x and y are ionic charges) ●the charges of the ions must add up to the overall charge, which is 0 in this case, so 2x + 3y = 0●we know that y = -2 (oxide ion) 2x + 3 (-2) = 0 x = +3●so Fe2O3 is named iron (III) oxide

33. H2 has a nonpolar covalent bond. A water molecule is polar – due to polar covalent bonds and the shape of the molecule. NaCl NaCl has an ionic bond – look at the EN difference. Na = 1.0 Cl = 2.9 EN = 1.9 5.1

34. Polyatomic Ions These are covalently bonded atoms with an overall charge (an ionic molecule):NO3-1 – nitrate ion ClO3-1 – chlorate ion C2H3O2-1 – acetate ion OH-1 – hydroxide ion SO4-2 – sulfate ion CO3-2 – carbonate ion PO4-3 – phosphate ion H3O+1 – hydronium ion NH4+1 – ammonium ion (NH3 – ammonia)

35. Polar bonds can result in polar molecules. For molecules like CO2, the polar bonds cancel each other out. For other molecules, like water, the polar bonds cause slight positive and negative ends on each molecule. Intermolecular Forces

36. The dipoles on one molecule are attracted to the dipoles on other molecules. This is an example of intermolecular attractive force. Water molecules are extremely polar, and so have strong intermolecular attraction. This is why water has such a high boiling point. N2 is a heavier molecule, but with little intermolecular attraction, it's boiling point is 300°C lower than that of water Intermolecular Forces

37. Polarized bonds allow hydrogen bonding to occur. A hydrogen bond is an electrostatic attraction between an atom bearing a partial positive charge in one molecule and an atom bearing a partial negative charge in a neighboring molecule. The H atom must be bonded to an O, N, or F atom. Hydrogen bonds typically are only about one-fifteenth as strong as the covalent bonds that connect atoms together within molecules. H–bonds are intermolecular bonds. Covalent bonds are intramolecular bonds. 5.2

38. The dipole/dipole interaction is like a weak ionic bond. For this reason, water is also able to dissolve many ionic compounds. Several water molecules can surround ions and bring them into solution. Intermolecular Forces

39. Substances that will dissociate in solution are called electrolytes. Ions are simply charged particles – atoms or groups of atoms. They may be positively charged – cations. Or negatively charged –anions. Dissolution of NaCl in Water The polar water molecules stabilize the ions as they break apart (dissociate). H2O NaCl(s)Na+(aq) + Cl–(aq) 5.8

40. Polar bonds can result in polar molecules. For molecules like CO2, the polar bonds cancel each other out. For other molecules, like water, the polar bonds cause slight positive and negative ends on each molecule. Intermolecular Forces

41. Water on Earth • Only 3% of all water is fresh (potable) • Of this: • 68% is in glaciers • 30% is underground • 1% in the atmosphere • only 0.3% in lakes, rivers, streams

42. Water Footprint • The average person needs 1E6 (1 million) liters per year. • This is equivalent to 250000 gallons, half of an Olympic sized pool. • Some of this water is used directly, and some indirectly.

43. Directly used water includes: • Drinking water • Bathing water • Water used for washing dishes • Water used for washing clothes • Toilet water

44. Indirectly used water includes: • Water for crops • Water for livestock • Water needed for services, • Electrical power • Waste treatment • Water needed for industry • Production of consumer goods • Construction

45. Water Footprint Water is necessary to produce food: 5.3

46. Water Footprint Water is necessary for products: 5.3

47. Fresh Water • Surface water – Lakes, rivers, streams • Easily accessible • Not abundant enough to meet our needs • May need to be filtered/treated to drink • Ground water – Underground in aquifers (trapped in geological formations) • Harder to access • More abundant • Often drunk without treatment

48. The average American usesalmost 100 gallons of water a day. Nearly ¾ of the water enteringour homes goes down the drain. Much of our clean water comes from underground aquifers. The Ogallala Aquifer is shown in dark blue. While normally free of pollutants, groundwater can be contaminated by a number of sources: Abandoned mines Runoff from fertilized fields Poorly constructed landfills and septic systems Household chemicals poured down the drain or on the ground 5.4

49. Salt Water • Very abundant (97% of all water on Earth) • Easily accessible, at least near coasts • Not potable due to high salt content • Difficult to purify for human consumption

50. Access to safe drinking water varies widely across the world. 5.4