Chapter 3

The covalent bond of a water molecule Oxygen being more electronegative than hydrogen attracts the electron of Hydrogen close to it. This results in a slightly positive charge on the Hydrogen side and a slightly negative charge on the oxygen side. They form a covalent bond This results in a polar covalent molecule H 2 O

The polarity of water results in Hydrogen bonding

Weak hydrogen bonds break and reform in trillionths of a second giving water its fluidity in the liquid state The polarity of one water molecule makes it attracted to other molecules resulting in the formation of weak Hydrogen bonds between water molecules

Four emergent properties as a result of hydrogen bonding in Water 1.Cohesion 2.Moderation of temperature 3.Insulation of bodies of water by floating ice 4.Universal solvent

1. Cohesive Forces 1.Cohesion - water molecules are attracted to other water molecules water’s “stickiness” -surface tension = cohesion of water molecules at the surface of a body of water

Basilisk: The “Jesus Christ” Lizard

Adhesion of water molecules to other surfaces or substances Adhesion -Water is attracted to other materials Ex. Capillary action in a straw -water “climbs” up a straw -water molecules are attracted to the straw molecules -When one water molecule moves closer to a the straw molecules the other water molecules (which are cohesively attracted to that water molecule) also move up into the straw Plants use capillary action to draw water from the ground into themselves Adhesion of water to the cell surfaces

2. Moderation of temperature High Specific Heat Index Water moderates air temperature by absorbing heat that is warmer and releasing air that is cooler. High specific heat- water takes longer to heat up and cool down because of the hydrogen bonds. H bonds resist breaking -takes much energy before bonds break, Ex. -Moderates temperature on Earth

Evaporative Cooling Hottest water molecules leave first, reducing the average temp. of those remaining Ex. -Cools tropical oceans -Keeps coastal areas more moderate

Insulation of bodies of water by floating ice Physical States Which is most dense: solid, liquid, gas? Usually solid, but in water, liquid is most dense Water forms a crystalline structure as a solid Breaking and reforming of H bonds slow down to form a lattice work of ice crystals

Ice and Liquid Water

Water as the “Universal Solvent” Water dissolves more substances than any other liquid. Solution -mix of 2 or more substances Solvent -dissolving agent Solute -what dissolves Polar substances dissolve in water(a polar substance) (like dissolves like)

Hydrophobic versus Hydrophilic Hydrophobic – Non ionic – non polar substances ( equal sharing of electrons ) seem to repel water “water fearing” Ex fats lipids Phospholipids of the cell membrane have a hydrophobic side separates cell from outside environment

Hydrophilic Hydrophilic – Polar and ionic substances have an affinity to water “water loving” Can dissolve or not in water colloids become suspended in water molecules without dissolving Ex. Cotton a large molecule of cellulose fibers with partial + and – ends attracts and absorbs water easily.

Hydrophilic substances that dissolve become dissociated in water. Ex. table salt NaCl dissociates in water

Water can dissociate into a weak acid (although very rare) H 2 O can separate (dissociate) into H+ and OH- The electron of the lost H is left behind with OH leaving a Hydrogen proton H+ and hydroxide ion OH- The H+ bonds to another water molecule to form H 3 O+ Normally OH- = H+ This is a neutral pH

Dissociation of a water molecule

Acids, Bases and the pH scale pH stands for “potential Hydrogen” The level of pH is measured on a scale of 0 to 14 with 0 representing the highest concentration of acid and 14 representing the most alkaline or strongest base Some solutes added to water disrupt the balance of H 2 O by increasing or decreasing H+

Acids More H+ Acid - any substance that increases H + ions in solution Strong acids dissociate completely in water Ex Hydrochloric acid HCl strong acid HCl  H + + Cl – Weak acids do not completely dissociate in water

Bases Base is any substance that Accepts H + ions Ex. Ammonia (relatively weak base) NH 3 accepts H + ions The Nitrogen attracts the H + to form NH 4 + OR reduces the amount of H + by contributing OH - Ex. NaOH (sodium hydroxide) dissociates into Na+ and OH -

pH scale is a logarithm pH is a logarithm scale in powers of 10 that means for every number you go up or down on the scale the concentration of Hydrogen ions increases or decreases x 10

Buffers Cell’s internal pH is close to 7 Changes in pH can be harmful to chemical processes which are sensitive to fluctuations in concentrations of H + and OH - Buffers maintain pH balance by minimizing changes in H + and OH - Blood must maintain a pH of 7.4

Buffers are weak acid-base pairs Buffers accept H+ ions when pH is too low or acidic Buffers donate H+ ions when pH is too high or basic (alkaline)

Carbonic acid-bicarbonate buffering system Weak acids and their corresponding base are used as buffers Weak acids reversibly release and accept back Hydrogen ions to maintain internal pH levels Carbonic acid H 2 CO 3 maintains pH equilibrium in biological fluids

Carbonic Acid Carbonic acid H 2 CO 3 dissociates to yield a bicarbonate ion HCO 3 - and a Hydrogen ion H +

Threats to Water Quality on Earth Acid precipitation refers to rain, snow, or fog with a pH lower than 5.6 Acid precipitation is caused mainly by the presence of sulfur oxides and nitrogen oxides which react with water to form strong acids, which fall to Earth as rain or snow Sources include burning of fossil feuls ( coal, oil, gas) from factories and automobiles. Copyright © 2008 Pearson Education, Inc., publishing as Pearson Benjamin Cummings

Winds carry pollutants away and acid rain can fall miles from its source Acid precipitation can damage life in lakes and streams Effects of acid precipitation on soil chemistry are contributing to the decline of some forests

Fig More acidic 0 Acid rain Acid rain Normal rain More basic