University:
Kentucky State University
Course:
CHE 109 | Chemistry in Context
Academic year:

2024
Views:

69

Pages:

63
Author:

Andrioliwbf

Chapter 3 Water and Life Lectures by Erin Barley Kathleen Fitzpatrick Overview: The Molecule That Supports All of Life • Water is the biological medium on Earth • All living organisms require water more than any other substance • Most cells are surrounded by water, and cells themselves are about 70–95% water • The abundance of water is the main reason the Earth is habitable Figure 3.1 Concept 3.1: Polar covalent bonds in water molecules result in hydrogen bonding • The water molecule is a polar molecule: the opposite ends have opposite charges • Polarity allows water molecules to form hydrogen bonds with each other Animation: Water Structure Figure 3.2 − Hydrogen bond + + Polar covalent bonds − − + + − Figure 3.UN01 Concept 3.2: Four emergent properties of water contribute to Earth’s suitability for life • Four of water’s properties that facilitate an environment for life are – Cohesive behavior – Ability to moderate temperature – Expansion upon freezing – Versatility as a solvent Cohesion of Water Molecules • Collectively, hydrogen bonds hold water molecules together, a phenomenon called cohesion • Cohesion helps the transport of water against gravity in plants • Adhesion is an attraction between different substances, for example, between water and plant cell walls Animation: Water Transport Figure 3.3 Adhesion Two types of water-conducting cells Cohesion Direction of water movement 300 m Figure 3.3a Two types of water-conducting cells 300 m • Surface tension is a measure of how hard it is to break the surface of a liquid • Surface tension is related to cohesion Figure 3.4 Moderation of Temperature by Water • Water absorbs heat from warmer air and releases stored heat to cooler air • Water can absorb or release a large amount of heat with only a slight change in its own temperature Heat and Temperature • Kinetic energy is the energy of motion • Heat is a measure of the total amount of kinetic energy due to molecular motion • Temperature measures the intensity of heat due to the average kinetic energy of molecules • The Celsius scale is a measure of temperature using Celsius degrees (°C) • A calorie (cal) is the amount of heat required to raise the temperature of 1 g of water by 1°C • The “calories” on food packages are actually kilocalories (kcal), where 1 kcal = 1,000 cal • The joule (J) is another unit of energy where 1 J = 0.239 cal, or 1 cal = 4.184 J Water’s 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 1ºC • The specific heat of water is 1 cal/g/ºC • Water resists changing its temperature because of its high specific heat • Water’s high specific heat can be traced to hydrogen bonding – Heat is absorbed when hydrogen bonds break – Heat is released when hydrogen bonds form • The high specific heat of water minimizes temperature fluctuations to within limits that permit life Figure 3.5 Los Angeles (Airport) 75° 70s (°F) 80s San Bernardino 100° Riverside 96° Santa Ana Palm Springs 84° 106° Burbank 90° Santa Barbara 73° Pacific Ocean 68° 90s 100s San Diego 72° 40 miles Evaporative Cooling • Evaporation is transformation of a substance from liquid to gas • Heat of vaporization is the heat a liquid must absorb for 1 g to be converted to gas • As a liquid evaporates, its remaining surface cools, a process called evaporative cooling • Evaporative cooling of water helps stabilize temperatures in organisms and bodies of water Floating of Ice on Liquid Water • Ice floats in liquid water because hydrogen bonds in ice are more “ordered,” making ice less dense • Water reaches its greatest density at 4°C • If ice sank, all bodies of water would eventually freeze solid, making life impossible on Earth Figure 3.6 Hydrogen bond Ice: Hydrogen bonds are stable Liquid water: Hydrogen bonds break and re-form Figure 3.6a Water: The Solvent of Life • A solution is a liquid that is a homogeneous mixture of substances • A 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 • Water is a versatile solvent due to its polarity, which allows it to form hydrogen bonds easily • When an ionic compound is dissolved in water, each ion is surrounded by a sphere of water molecules called a hydration shell Figure 3.7 − Na+ + − + − Na+ + Cl− − + − Cl− + − + − + − − − + − • Water can also dissolve compounds made of nonionic polar molecules • Even large polar molecules such as proteins can dissolve in water if they have ionic and polar regions Figure 3.8 + − − + Hydrophilic and Hydrophobic Substances • A hydrophilic substance is one that has an affinity for water • A hydrophobic substance is one that does not have an affinity for water • Oil molecules are hydrophobic because they have relatively nonpolar bonds • A colloid is a stable suspension of fine particles in a liquid Solute Concentration in Aqueous Solutions • Most biochemical reactions occur in water • Chemical reactions depend on collisions of molecules and therefore on the concentration of solutes in an aqueous solution • Molecular mass is the sum of all masses of all atoms in a molecule • Numbers of molecules are usually measured in moles, where 1 mole (mol) = 6.02 x 1023 molecules • Avogadro’s number and the unit dalton were defined such that 6.02 x 1023 daltons = 1 g • Molarity (M) is the number of moles of solute per liter of solution Possible Evolution of Life on Other Planets with Water • The remarkable properties of water support life on Earth in many ways • Astrobiologists seeking life on other planets are concentrating their search on planets with water • To date, more than 200 planets have been found outside our solar system; one or two of them contain water • In our solar system, Mars has been found to have water Figure 3.9 Concept 3.3: Acidic and basic conditions affect living organisms • A hydrogen atom in a hydrogen bond between two water molecules can shift from one to the other – The hydrogen atom leaves its electron behind and is transferred as a proton, or hydrogen ion (H+) – The molecule with the extra proton is now a hydronium ion (H3O+), though it is often represented as H+ – The molecule that lost the proton is now a hydroxide ion (OH–) • Water is in a state of dynamic equilibrium in which water molecules dissociate at the same rate at which they are being reformed Figure 3.UN02 + 2 H 2O Hydronium ion (H3O+) − Hydroxide ion (OH−) • Though statistically rare, the dissociation of water molecules has a great effect on organisms • Changes in concentrations of H+ and OH– can drastically affect the chemistry of a cell • Concentrations of H+ and OH– are equal in pure water • Adding certain solutes, called acids and bases, modifies the concentrations of H+ and OH– • Biologists use something called the pH scale to describe whether a solution is acidic or basic (the opposite of acidic) Acids and Bases • An acid is any substance that increases the H+ concentration of a solution • A base is any substance that reduces the H+ concentration of a solution The pH Scale • In any aqueous solution at 25°C the product of H+ and OH– is constant and can be written as [H+][OH–] = 10–14 • The pH of a solution is defined by the negative logarithm of H+ concentration, written as + pH = –log [H ] • For a neutra[H la lu–ti7 o)n = 7 +]qiu se1o0u–s7 so = –( • Acidic solutions have pH values less than 7 • Basic solutions have pH values greater than 7 • Most biological fluids have pH values in the range of 6 to 8 Figure 3.10 H+ H+ − H+ H+ OH + − OH H H+ + H H+ Acidic solution Increasingly Acidic [H+] > [OH−] pH Scale 0 1 Battery acid 2 Gastric juice, lemon juice 3 Vinegar, wine, cola 4 Tomato juice Beer Black coffee 5 6 Neutral solution OH− OH− OH− H+ OH− − OH− OH OH− + H Basic solution Neutral [H+] = [OH−] 7 8 Increasingly Basic [H+] < [OH−] OH− OH− H+ H+ OH− − OH− OH + H+ H+ H Rainwater Urine Saliva Pure water Human blood, tears Seawater Inside of small intestine 9 10 Milk of magnesia 11 Household ammonia 12 13 Household bleach Oven cleaner 14 Figure 3.10a Figure 3.10b Figure 3.10c Figure 3.10d Buffers • The internal pH of most living cells must remain close to pH 7 • Buffers are substances that minimize changes in concentrations of H+ and OH– in a solution • Most buffers consist of an acid-base pair that reversibly combines with H+ Acidification: A Threat to Water Quality • Human activities such as burning fossil fuels threaten water quality • CO2 is the main product of fossil fuel combustion • About 25% of human-generated CO2 is absorbed by the oceans • CO2 dissolved in sea water forms carbonic acid; this process is called ocean acidification Figure 3.11 CO2 CO2 + H2O H2CO3 H2CO3 H+ + HCO3− H+ + CO32− CO32− + Ca2+ HCO3− CaCO3 • As seawater acidifies, H+ ions combine with carbonate ions to produce bicarbonate • Carbonate is required for calcification (production of calcium carbonate) by many marine organism, including reef-building corals Figure 3.12 (a) (b) (c) Figure 3.12a (a) Figure 3.12b (b) Figure 3.12c (c) • The burning of fossil fuels is also a major source of sulfur oxides and nitrogen oxides • These compounds react with water in the air to form strong acids that fall in rain or snow • Acid precipitation is rain, fog, or snow with a ph lower than 5.2 • Acid precipitation damages life in lakes and streams and changes soil chemistry on land Figure 3.UN03 − + + − − + + − Figure 3.UN04 Ice: stable hydrogen bonds Liquid water: transient hydrogen bonds Figure 3.UN05 0 Acidic [H+] > [OH−] Neutral [H+] = [OH−] Basic [H+] < [OH−] Acids donate H+ in aqueous solutions. 7 Bases donate OH− or accept H+ in aqueous solutions 14 Calcification rate (mmol CaCO3/m2 • day) Figure 3.UN06 40 20 0 200 250 [CO32−] (mol/kg) Figure 3.UN07 Figure 3.UN08 Figure 3.UN09

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