Chapter 3: Water and Life PDF

Chapter 3: Water and Life Objectives 1. You will define of polarity and how it is associated with hydrogen bonds. 2. You will describe the 4 properties of water, what is responsible for them, and why they are important for life. 3. You will know why water makes a good solvent 4. You will...

Chapter 3: Water and Life Objectives 1. You will define of polarity and how it is associated with hydrogen bonds. 2. You will describe the 4 properties of water, what is responsible for them, and why they are important for life. 3. You will know why water makes a good solvent 4. You will define of molarity and how to make solutions. 5. You will know the pH scale and how the pH of water and other substances are calculated (not the equation but what that number means). 6. You will define buffers, how they function and why they might be important in biology Atoms Ionic bonds In the formation of an ionic compound, metals lose electrons and nonmetals gain electrons to achieve an octet. Covalent bonds Two or more atoms may bond with each other to form a molecule. When two hydrogens and an oxygen share electrons via covalent bonds, a water molecule is formed. Water and Life  Three-quarters of Earth’s surface is covered by water  Most in liquid form  Some exists as a solid (ice) and a gas (water vapor)  Only common substance on Earth that exists in the environment as all three states of matter  The unique properties of water are due to polar covalent bonds and hydrogen bonds Polar covalent bonds  A water molecule consists of an oxygen atom and two hydrogen atoms (two single covalent bonds)  Oxygen is more electronegative than hydrogen so electrons spend more time closer to oxygen  The result of these polar covalent bonds are partial charges Hydrogen bonds Partially + hydrogen atoms form bonds to partially negative oxygen atoms Emergent properties of water  Four of water’s properties that facilitate an environment for life:  Cohesive behavior  Ability to moderate temperature  Expansion upon freezing  Versatility as a solvent Cohesion of Water Molecules  Hydrogen bonds hold water molecules together (Cohesion)  Cohesion helps the transport of water against gravity in plants  Water from the roots reaches the leaves as water evaporates from the leaves  Hydrogen bonds “pull” water up  Adhesion (clinging of water to another substance) allows water to cling to the cell walls as it moves upward Cohesion of Water Molecules  The cohesion of water is also responsible for surface tension  Water molecules stick together (through hydrogen bonds) on the surface of water and forms a film Moderation of Temperature by Water  Water absorbs heat from warmer air and releases stored heat to cooler air  For example: Living next to a body of water  Water resists changing its temperature  This is due to its high specific heat  The specific heat of a substance is the amount of heat that must be absorbed or lost for 1 g of that substance to change its temperature by 1°C Moderation of Temperature by Water  Calorie is a measure of heat/energy  The specific heat of water is 1 calorie/(g  °C)  Water has a high specific heat compared to other substances  Ethanol has a specific heat of 0.6 calorie/(g  °C) Moderation of Temperature by Water  It takes a lot of heat to move water molecules which means it can absorb more heat than other substances before it changes temperature  Water also requires a lot of heat to be converted to a gas (heat of vaporization)  For example, alcohols evaporates faster than water Moderation of Temperature by Water  When water eventually evaporates, the surface of water that remains cools down (evaporative cooling)  Some of the water molecules absorb heat and evaporate which means there’s less heat left in the remaining water (or on surface)  For example, sweating  What about sweating in high or low humidity? Floating of Ice on Liquid Water  Water is less dense as a solid than as liquid which means it floats on water  Water molecules move when hydrogen bonds are broken (due to heat)  As water cools, hydrogen bonds do not break as often  In ice, hydrogen bonds are unbroken Structure of liquid and ice Floating of Ice on Liquid Water  Floating ice is important for life  If ice sank, all ponds, lakes, and oceans would freeze solid (instead of just the surface)  Floating ice also insulates the liquid water below  Allows life to exist under the sheet of ice  Also allows life to exist over the sheet of ice  This is one concern about global warming Water: The Solvent of Life  A solution is a liquid that is a completely homogeneous mixture of substances  The solvent is the dissolving agent of a solution  The solute is the substance that is dissolved  An aqueous solution is one in which water is the solvent  Polar covalent bonds make water a good solvent Water: The Solvent of Life  Water can also dissolve compounds made of nonionic polar molecules (such as proteins)  Or any other hydrophilic (have an affinity for water) substances Water: The Solvent of Life When table salt (NaCl) is mixed in water, spheres of hydration are formed around the ions. Water: The Solvent of Life  Water can also dissolve compounds made of nonionic polar molecules (such as proteins)  Or any other hydrophilic (have an affinity for water) substances  Non polar molecules (hydrophobic) do not dissolve in water  This also supports life  Cell membranes are composed of non polar molecules Solute concentration in solutions  The unit of concentration used most often by scientists is Molarity (M)  Concentrations are sometimes written in brackets, [NaCl]  Molarity is the number moles of solute per liter of solution  Moles (mol) is 6.02 x 1023 molecules of any substance (aka Avogadro’s number)  Although the number of moles is the same for any substance, the molecular mass is not Solute concentration in solutions  For example, 1M solution of sodium chloride (NaCl) in water  1 mol of sodium chloride in 1 liter of water  Molecules can not be counted but they can be weighed  The weight of 1 mol of any substance is the molecular mass of that substance Solute concentration in solutions  Molecular mass is the sum of the mass numbers of all the atoms in the molecule 11 17 Na Cl 23 35  Molecular mass of sodium chloride is 58  You need 58 grams of sodium chloride in 1 liter of water to make 1 M solution. Acidic and basic conditions affect living organisms  pH and the pH scale is used to measure how acidic or basic a solution is  On the pH scale, water is neutral (establishes a baseline for all other acids and bases) Dissociation of water molecules  Hydronium ion can also be represented by a hydrogen ion (H+)  In water, this rarely occurs  The concentration of H+ in water is 0.0000001M or 10-7 M Acids and Bases  An acid is any substance that increases the concentration H+ in a solution For example, HCl ⇨ H+ + Cl- an acid can also decrease [OH-] Acids and Bases  A base is any substance that increases the concentration OH- in a solution For example, NaOH ⇨ Na+ + OH- a base can also decrease [H+] The pH scale  pH = -log [H+]  The concentration of H+ in water is 10-7 M  pH of water is 7 (neutral solution)   A solution with a pH less than 7 means it’s acidic  A solution with a pH more than 7 means it’s basic The concentration H+ in water is 10-7 and the concentration of H+ in urine is 10-6 Urine is 10 times more acidic than water (10x more H+) Coffee is 100 times (10x10) Buffers  Internal pH of most living cells is close to 7  Biological fluids fall somewhere between 6-8  H+ and OH- ions can be harmful to living things  Living things require buffers that minimize changes in the concentration of H+ and OH- in biological fluids Buffers  Buffers can accept hydrogen ions when they are in excess or donate them when they are needed  The pH of buffers remain constant when you add acids or bases to them  Typically a weak acid (which means it easily accepts or donate hydrogens)  For example: Carbonic acid is found in human blood Buffers  Carbonic acid (H2CO3) is formed when carbon dioxide (CO2) reacts with water in blood and can dissociate to bicarbonate ion and a hydrogen ion Carbonic acid lowers pH by donating hydrogen Bicarbonate raises pH by accepting hydrogen