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LECTURE PRESENTATIONS For CAMPBELL BIOLOGY, NINTH EDITION Jane B. Reece, Lisa A. Urry, Michael L. Cain, Steven A. Wasserman, Peter V. Minorsky, Robert B. Jackson Chapter 2 The Chemical Context of Life Lectures by Erin Barley Kathleen Fitzpatrick © 2011 Pearson Education, Inc. Overview: A Chemical Connection to Biology • Biology is a multidisciplinary science • living organisms are subject to basic laws of physics and chemistry • this chapter deals with the relevance of chemistry to biology • one example are the “devil’s gardens” • • • • • stands of Duroia trees dominates its landscape not planted or maintained by humans called devil’s gardens by the locals who tends the garden??? Dead leaf tissue (cm2) after one day RESULTS 16 12 8 4 0 Inside, unprotected Inside, protected Outside, unprotected Outside, protected Cedrela saplings, inside and outside devil’s gardens • • • Stanford biologists studied the question one hypothesis was that ants living in the trees produce a poison that kills other plants field experiments – planted saplings of another tree Cedrela inside the “garden” – amongst the Duria trees • • • wanted to see if the pre-existing Duria trees killed the Cedrela transplants to see if the ants had any affect on viability - applied an insect barrier to some of the Cedrela saplings planted; the others were unprotected and vulnerable to the Duria trees to control for the location – the same Cedrela saplings also planted outside the garden Dead leaf tissue (cm2) after one day RESULTS 16 12 8 4 0 Inside, unprotected Inside, protected Outside, unprotected Outside, protected Cedrela saplings, inside and outside devil’s gardens • • • • • RESULTS: the unprotected Cedrela saplings inside the garden had the most damage CONCLUSIONS: 1. the location has no effect – inside protected and outside protected Cedrela trees had the same viability 2. BUT protection was important – unprotected Cedrela trees planted next to the Duria trees died • so the ants had an effect ants were found to kill non-host trees by injecting formic acid into their leaves Life: Levels of Organization •Atoms •Molecules •Macromolecules •Organelles •Cells •Tissues •Organs •Organ systems •Organism Concept: Matter consists of chemical elements in pure form and in combinations called compounds • organisms are composed of matter • matter is anything that takes up space and has mass • made up of elements • an element is a substance that cannot be broken down into other substances by chemical reactions • a compound is a substance consisting of two or more elements in a fixed ratio • a compound has characteristics different from those of its elements Sodium Chlorine Sodium chloride The Elements of Life • about 20–25% of the 92 natural elements are essential to life = essential elements • • • • • needed to live a healthy life and reproduce Table 2.1 of the total amount of natural elements on Earth - carbon, hydrogen, oxygen, and nitrogen make up 96% most of the remaining 4% consists of calcium, phosphorus, potassium, and sulfur trace elements are those required by an organism in minute quantities Concept: An element’s properties depend on the structure of its atoms Atom = smallest unit of an element that still retains the chemical & physical properties of that element i.e. really, really, really tiny thing! -composed of subatomic particles: 1. protons = one positive charge, 1 atomic mass unit (1.673x10-24g) 2. electrons = one negative charge, no mass (9.109x10-28g) 3. neutrons = no charge, 1 atomic mass unit (1.673x10-24g) • • • neutrons and protons form the atomic nucleus electrons form a cloud that “orbits” around the nucleus neutron mass and proton mass are almost identical and are measured in Daltons (Da) – but can be measured in grams PERIODIC TABLE OF ELEMENTS -elements are grouped on a Periodic Table of Elements -the elements are grouped according to physical and chemical characteristics -on the chart each element is associated with a letter, an atomic number & an atomic mass IA IIA http://oxford-labs.com/xray-fluorescence/theperiodic-table/ IIIA IVA VA VIA VIIA VIII Atomic Number and Atomic Mass • atoms of the various elements differ in number of subatomic particles • an element’s atomic number is the number of protons in its nucleus • • Atomic number = # protons this equals the number of electrons orbiting when the atom is electrically neutral • an element’s mass number is the sum of protons plus neutrons in the nucleus • Mass number = Protons + Neutrons • atomic mass = the atom’s total mass • • can be approximated by the mass number so an atom’s mass number is close to its atomic mass (weight) © 2011 Pearson Education, Inc. atomic symbol atomic mass (weight) 7 3 Li e.g. # protons (e-) = 3 # pr+3 + #No 4 = 7 atomic number 39 19 K e.g. # protons (e-) = 19 # pr+19 + #No 20 = 39 Isotopes • all atoms in an element have the same number of protons but may differ in number of neutrons • the number of protons defines what the element is • e.g. 6 protons = carbon and only carbon • isotopes are two atoms of an element that differ in number of neutrons • in nature an element is a mixture of its isotopes © 2011 Pearson Education, Inc. Radioactivity • the protons and neutrons are held together in the nucleus by a kind of nuclear “glue” • when the number of neutrons increase – the nucleus becomes unstable • the breakup of the nucleus releases particles with energy in the form of radioactivity • • also known as radioactive decay three different kind of particles released – each with different energy levels • • alpha (helium atom), beta and gamma the decay can eventually change the # protons – transform one atom into another pr+: e-: No: 12C 13C 6 6 6 6 6 7 14C 6 6 8** radioactive Radioactive isotope uses: 1. carbon dating - 14C decay 2. radioactive imaging - e.g. PET scanning -use of FDG – radioactive glucose tracer -18F radioactive isotope (2-fluoro-deoxy-glucose) 3. cancer treatment - 60Co, 131I Cancerous throat tissue The Energy Levels of Electrons • energy is the capacity to cause change • potential energy is the energy that matter has because of its location or structure • the electrons of an atom differ in their amounts of potential energy • differ in their location around the nucleus • an electron’s state of potential energy (i.e. its location around a nucleus) is called its energy level or electron shell © 2011 Pearson Education, Inc. Electron Configurations • “bed check” for electrons • description on how are electrons organized around the nucleus of protons and neutrons • Bohr model: Nils Bohr proposed electrons “orbit” around the atom’s nucleus in specific energy levels or orbits (electron shells) – these shells have a specific energy level – closer the electron is to the nucleus the less energy it needs to “orbit” – to move up to a new electron shell requires an input of energy – as the electrons returns to its correct shell (its “ground” state) – it releases this energy – his model only works for smaller atoms – larger atoms are described by quantum mechanics (a)A ball bouncing down a flight of stairs provides an analogy for energy levels of electrons. Third shell (highest energy level in this model) Second shell (higher energy level) First shell (lowest energy level) (b) Atomic nucleus Energy absorbed Energy lost • the Bohr model proposed that not only are there electron shells that surround the nucleus • but each shell is comprised of subshells called orbitals • an orbital is the three-dimensional space where an electron pair is found 90% of the time • each electron shell consists of a specific number of orbitals • s, p and d orbitals First shell Neon, with two filled Shells (10 electrons) Second shell (a) Electron distribution diagram First shell Second shell y x 1s orbital 2s orbital z Three 2p orbitals (b) Separate electron orbitals 1s, 2s, and 2p orbitals (c) Superimposed electron orbitals – – – – – – – – each orbital holds a pair of electrons s orbital = 2 electrons maximum p orbitals = 6 electrons maximum d orbitals = 10 electrons maximum 1st shell – closest to the nucleus - holds 2 electrons (1 s orbital only) 2nd shell can hold 8 (1 s and 3 p orbitals – 2 + 6 electrons) 3rd holds 18 (1 s, 3 p and 5 d orbitals – 2 + 6 + 10 electrons) 4th holds 18 (1 s, 3 p and 5 d orbitals – 2 + 6 + 10 electrons) Electron Distribution and Chemical Properties • the chemical behavior of an atom is determined by the distribution of electrons in electron shells • the reactivity of certain atoms results from the presence of unpaired electrons in one or more orbitals • an atom will always try to complete its outermost shell = valence shell • basis for bonding reactions • the number of electrons in the outer most electron shell involved in bonding reactions = valence electrons • chemists really only consider the electrons in the s and p orbitals as valence electrons • once an atom completes fills up its valence orbitals – it is chemically inert • unable to participate in bonding reactions © 2011 Pearson Education, Inc. Valence = 1 Valence = 4 Valence = 3 Valence = 2 VIII I Hydrogen 1H Electron distribution diagram First shell II Lithium 3Li Beryllium 4Be III Boron 5B IV Helium 2He V VI VII Carbon 6C Nitrogen 7N Oxygen 8O Fluorine 9F Neon 10Ne Silicon 14Si Phosphorus 15P Sulfur 16S Chlorine 17Cl Argon 18Ar Second shell Sodium Magnesium Aluminum 11Na 12Mg 13Al Third shell • the periodic table of the elements tells you the electron distribution for each element • • • • • by its row and column position 1st row – valence electrons in the first electron shell 2nd row – valence electrons in the second electron shell 1st column – 1 valence electron 3rd column – 3 valence electrons Concept: The formation and function of molecules depend on chemical bonding between atoms • atoms with incomplete valence shells can share or transfer valence electrons with certain other atoms to form molecules • molecule - particle formed by the union of more than one atom • e.g. same kind of atom - O2 • e.g. different types of atoms - H20 • formed by held by attractions called chemical bonds • two types of chemical bonds © 2011 Pearson Education, Inc. Ionic Bonds • atoms sometimes strip electrons from their bonding partners • • • • • • an example is the transfer of an electron from sodium to chlorine after the transfer of an electron, both atoms have charges a charged atom (or molecule) is called an ion a cation is a positively charged ion an anion is a negatively charged ion an ionic bond is an attraction between an anion and a cation Na Sodium atom Cl Chlorine atom + – Na+ Sodium ion (a cation) Cl– Chloride ion (an anion) Sodium chloride (NaCl) • Compounds formed by ionic bonds are called ionic compounds, or salts • • salts, such as sodium chloride (table salt), are often found in nature as crystals HINT: elements in columns I and II form ionic bonds with the elements in column VII Na+ Cl– © 2011 Pearson Education, Inc. Covalent Bonds • if it isn’t favorable for an atom to gain or lose an electron - it will have to share it with another • covalent bond = bond in which atoms share electrons • atoms that like to form covalent bonds • oxygen • nitrogen • carbon -usually forms when one atom has to lose or gain three or more electrons Hydrogen atoms (2 H) e.g. carbon would have to gain 4 valence electrons to complete its outer shell, nitrogen would have to gain 3 valence electrons -can also form between two identical atoms e.g. nitrogen (N3), oxygen gas (O2), hydrogen gas (H2) Hydrogen molecule (H2) Covalent Bonds for the Biologist • • • • • • • covalent bonding by some common biological elements Hydrogen = valence 1; 1 electron needed; 1 covalent bond Oxygen = valence 2; 2 electrons needed; 2 covalent bonds Sulfur = valence 2; 2 electrons needed; 2, 4 or 6 covalent bonds Nitrogen = valence 3; 3 electrons needed; 3 or 4 covalent bonds Carbon = valence 4; 4 electrons needed; 4 covalent bonds Phosphorus = valence 3; 3 electrons needed; 5 covalent bonds © 2011 Pearson Education, Inc. • so we can now modify the definition of a molecule further • a molecule consists of two or more atoms held together by covalent bonds • a single covalent bond, or single bond, is the sharing of one pair of valence electrons • is notated in the structural formula as H-H • a double covalent bond, or double bond, is the sharing of two pairs of valence electrons • is notated in the structural formula as C=C © 2011 Pearson Education, Inc. Figure 2.12 Name and Molecular Formula (a) Hydrogen (H2) (b) Oxygen (O2) (c) Water (H2O) (d) Methane (CH4) Lewis Dot SpaceElectron Distribution Structure and Filling Structural Model Diagram Formula Polar and Non-Polar Covalent Bonds • atoms in a molecule attract electrons to varying degrees • electronegativity is an atom’s attraction for the electrons in a covalent bond • the more electronegative an atom the more strongly it pulls shared electrons toward itself • modifies the covalent bond • what are the two categories of covalent bond? Polar and Non-Polar Covalent Bonds • in a nonpolar covalent bond, the atoms share the electron equally • carbon frequently forms non-polar covalent bonds Polar and Non-Polar Covalent Bonds • in a polar covalent bond, one atom is more electronegative, and the atoms do not share the electron equally • unequal sharing of electrons causes a partial positive or negative charge for each atom or molecule • • annotated with the greek letter delta positive and negative ends are known as dipoles • oxygen is very electronegative – creates polar bonds • nitrogen is also very electronegative – • \ O + H H H2O + Weak Chemical Bonds • the strongest bonds in organisms are covalent bonds that form a cell’s molecules • BUT weak chemical bonds, such as ionic bonds and hydrogen bonds, are also important • weak chemical bonds reinforce shapes of large molecules • help molecules adhere to each other • weak non-covalent bonds are transient • constantly breaking and reforming at room temperature • but together multiple, weak non-covalent bonds can produce highly stable structures • most common weak chemical bonds? A Weak Bond: Hydrogen Bond • a hydrogen bond forms when a hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom • in living cells, the electronegative partners are usually oxygen or nitrogen atoms • • e.g. seen between the bases of the DNA double helix e.g. seen between two water molecules • not really a bond but an interaction between two dipoles • weaker than covalent and ionic © 2011 Pearson Education, Inc. + – Water (H2O) + Hydrogen bond – Ammonia (NH3) + + + Hydrogen bonds are intermolecular bonds Concept: 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 Hydrogen bond + + Polar covalent bonds + © 2011 Pearson Education, Inc. + Another Weak Bond: Van der Waals • if electrons are distributed asymmetrically in molecules or atoms, they can result in “hot spots” of positive or negative charge • Van der Waals interactions are attractions between molecules that are close together as a result of these charges • • named after the physicist Johannes Diedrick Van der Waals often used to describe the totality of intermolecular forces • collectively, these bonds can be strong • e.g. as between molecules of a gecko’s toe hairs and a wall surface van der Waals bonds are intermolecular bonds Molecular Shape and Function • a molecule’s shape is usually very important to its function • molecules such as H2 or O2 are always linear • but others like H2O have a 3D conformation s orbital • • a molecule’s shape is determined by the positions of its atoms’ valence orbitals when a covalent bond forms – the valence electrons undergo rearrangement • • • this creates specific molecular shapes for methane - creates a tetrahedron of four tear-drop like orbitals • • specifically the electrons in the s and p orbitals may hybridize with each other as they shift and rearrange within the molecule Four hybrid orbitals z Three p orbitals x y Tetrahedron (a) Hybridization of orbitals Space-Filling Model Ball-and-Stick Model Unbonded Electron pair Water (H2O) symmetrical shape for water – creates a V • assymetrical shape (angle of 104.5) Hybrid-Orbital Model (with ball-and-stick model superimposed) Methane (CH4) (b) Molecular-shape models Carbon Hydrogen Natural endorphin • • biological molecules recognize and interact with each other based on molecular shape creates specificity • • Nitrogen Sulfur Oxygen Morphine e.g. interaction of a hormone with its receptor e.g. interaction of an enzyme with its substrate (a) Structures of endorphin and morphine • • molecules with similar shapes can have similar biological effects IN OTHER WORDS – you can predict function by looking at 3D structure Natural endorphin Brain cell Morphine Endorphin receptors (b) Binding to endorphin receptors Concept: Chemical reactions make and break chemical bonds • Chemical reactions are the making and breaking of chemical bonds • The starting molecules of a chemical reaction are called reactants • The final molecules of a chemical reaction are called products 2 H2 + Reactants O2 2 H2O Reaction Products Chemical reactions: 3 types: 1. Synthesis - A + B 2. Decomposition - AB reactions) 3. Exchange - AB + CD AB (Anabolism reactions) A + B (Catabolism AD + BC -these equations must be balanced -Law of conservation of Mass or “chemical book keeping” -i.e. the number of atoms of each element is the same before and after a chemical reaction • All chemical reactions are reversible: products of the forward reaction become reactants for the reverse reaction • Chemical equilibrium is reached when the forward and reverse reaction rates are equal Water: The Solvent of Life • major component of blood plasma, interstitial fluid, CSF, cytosol etc… • hydrogen bonding allows the even distribution of dissolved substances throughout our system – so water is an excellent transport medium role in : transporting chemicals, waste products •solvent – substance that dissolves solutes • can assist in chemical reactions •water is a polar solvent + • dissolves polar solutes •water is the universal solvent for polar compounds – facilitates most chemical reactions in the body O H H + Concept: Four emergent properties of water contribute to its importance 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 #1: Cohesion of Water Molecules • collectively, hydrogen bonds hold multiple water molecules together in a well-ordered structure – this attractive phenomenon is called cohesion – at any given time many of water’s molecules are linked by hydrogen bonds • cohesion: attraction between the same molecules • explains the capillary action of water we see in plants cohesive force helps the transport of water and its dissolved substances against gravity in plants - pulls water molecules up the plant following evaporation from the leaf surface Property #1: Cohesion of Water Molecules • DON’T CONFUSE COHESION WITH ADHESION!!! • adhesion is an attraction between different molecules – for example, between water and plant cell walls • cohesion can also be observed in other liquids – e.g. mercury – cohesive attraction between Hg molecules is stronger than the adhesive attraction to another surface • collectively – all the water molecules connected to each other creates a cohesive force • surface tension is a measure of how hard it is to stretch or break the surface of a liquid – water has a greater surface tension than most liquids • surface tension is related to cohesion – – at the water-air interface is a thin film of well-ordered water molecules – all hydrogen bonded to one another creates a strong “invisible film” at this interface = known as surface tension Property #2: 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 • so water is great at stabilizing temperature • BUT – what is heat and temperature? © 2011 Pearson Education, Inc. 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 • high motion = higher heat • temperature measures the intensity of heat due to the average kinetic energy of molecules – – – • the unit of heat expended in a reaction is a calorie (cal) – – • temperature on Earth is measured with the Celsius scale (ºC) but the true measure of temperature is in Kelvin absolute zero (0ºK) – when all molecular motion stops the amount of heat required to raise the temperature of 1 g of water by 1 ºC 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 © 2011 Pearson Education, Inc. Water’s High Specific Heat • the ability of water to stabilize temperature comes from its relatively high specific heat • specific heat = amount of heat that must be absorbed or lost for 1 g of that substance to change its temperature by 1ºC – – – – specific heat of water is 1 cal/g/ºC specific heat of ethanol is 0.6 cal/g/ºC specific heat of iron – ten times less than water so its easier and faster to change the temperature of iron vs. water • so because of this - water changes its temperature very little when it absorbs or gives off heat • think of specific heat as how well a substance resists a temperature change © 2011 Pearson Education, Inc. • water’s high specific heat can be traced to hydrogen bonding – heat must be absorbed for hydrogen bonds to break – heat is released when these hydrogen bonds reform • the high specific heat of water minimizes temperature fluctuations to within limits that permit life Burbank 90° Santa Barbara 73° Los Angeles (Airport) 75° 70s (°F) 80s San Bernardino 100° Riverside 96° Santa Ana 84° Palm Springs 106° Pacific Ocean 68° 90s 100s © 2011 Pearson Education, Inc. San Diego 72° 40 miles Water and Evaporative Cooling • molecules of a liquid stay close together due to cohesive attraction • if they move fast enough (due to increased kinetic energy) – they can overcome this attraction and be released – – i.e. increase in kinetic energy can result from an increase in heat molecules are released as a gas • transformation from liquid to gas occurs = evaporation • heat of vaporization is the heat a liquid must absorb for 1 g to be converted to gas – water has a high heat of vaporization due to the strength of its hydrogen bonding • as a liquid evaporates – the remaining surface cools - a process called evaporative cooling – the “hottest” molecules leave as a gas and the remaining liquid cools • evaporative cooling of water helps stabilize temperatures in organisms and bodies of water © 2011 Pearson Education, Inc. Property #3: Expansion of Ice Upon Freezing • water is one of the few substances that is less dense as a solid than as a liquid – water expands when heated and contracts as it cools • in other words ice floats in liquid water • below 4°C water begins to freeze – molecules move too slowly to break their hydrogen bonds – molecules become locked in place (crystalline formation) • distance between molecules when frozen is greater than in liquid form – makes ice about 10% less dense vs. liquid form • when ice melts – the hydrogen bonds are broken and the molecules ‘slip’ closer to each other Hydrogen bond Liquid water: Hydrogen bonds break and re-form Ice: Hydrogen bonds are stable *** if ice sank, all bodies of water would eventually freeze solid, making life impossible on Earth**** Property #4: Water - The Solvent of Life • a solution is a liquid that is a homogeneous mixture of 2 or more solutes – 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 • water is a universal solvent because it dissolves just about everything – except non-polar compounds – e.g. oil and water • water is a versatile solvent due to its polarity – allows it to form hydrogen bonds easily with other charged molecules • when an ionic compound is dissolved in water - each ion becomes surrounded by water molecules - known as a solvation or hydration shell -water + salt: the electronegative O- of water attracts the +ve sodium in the salt crystal and pulls it away - the electropositive H+ of water attracts the –ve chlorine and pulls it away -the crystal lattice of salt is eventually broken up and each Na+ and Clbecomes surrounded by water molecules • water can also dissolve compounds made of polar molecules • even large polar molecules such as proteins can dissolve in water if they have ionic and polar regions + + 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 are relatively nonpolar in terms of their bonds Solutions, Colloids and Suspensions • mixture = two or more types of elements or molecules physically blended together in a solvent without the formation of physical bonds between them • 1. solution = homogenous (same) mixture of substances dissolved in a solvent – – – • e.g. sugar + water the mixture is the same no matter where you sample it substances are very small and are not acted upon by gravity – remain suspended in the solution 2. colloid = solution of larger components called dispersed-phase particles – – – their particles are larger than that of solutions have the potential to settle out due to gravity BUT: these particles all carry the same charge (repel each other) – so they remain suspended in the solvent • • e.g. plasma proteins within the blood 3. suspension = solution of even larger components that are solid and can settle out – – larger particles than that of colloid if left undisturbed these particles will settle out to form a solid • e.g. red blood cells within blood 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 • the number of molecules is usually measured in moles, where 1 mole (mol) = 6.02 x 1023 molecules – known as Avogadro’s number – the unit Dalton (Da) is defined such that 6.02 x 1023 daltons = 1 g • Molarity (M) is the number of moles of solute per liter of solution • to calculate the number of moles – you need to know molecular mass = total mass of all atoms in a molecule Concept: Acidic and basic conditions affect living organisms • between two water molecules something unique can happen • a hydrogen atom in a hydrogen bond 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–) + 2 H 2O © 2011 Pearson Education, Inc. Hydronium ion (H3O+) Hydroxide ion (OH) 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 • 1. release H+ Acids e.g. HCl H+ + Cl2. release ions to combine with H+ Bases e.g. NaOH Na+ + OH3. acids + bases Salts e.g. HCl + NaOH H20 + NaCl © 2011 Pearson Education, Inc. The pH Scale • biologists use something called the pH scale to describe whether a solution is acidic or basic (the opposite of acidic) • in any aqueous solution at 25C 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 neutral aqueous solution, [H+] is 10–7, so pH = –log(–7) = 7 H+ H+ H+ H+ OH + OH H H+ + H H+ Acidic solution 1 Battery acid 2 Gastric juice, lemon juice 3 Vinegar, wine, cola 4 Tomato juice Beer Black coffee 5 6 OH OH H+ H+ OH OH OH + + H H H+ Neutral solution OH OH OH H+ OH OH OH H+ OH Basic solution Neutral [H+] = [OH] 7 8 Increasingly Basic [H+] < [OH] • 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 Increasingly Acidic [H+] > [OH] pH Scale 0 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 Buffers: • biological fluids need to remain relatively neutral • chemical or compound that keeps the pH of a solution within a normal range • resists pH change by taking up excess H+ or OH- ions H20 H2CO3 e.g. blood = pH 7.4 “Bicarbonate buffering system” H+ + HCO3- carbonic acid excess OH- © 2011 Pearson Education, Inc. excess 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 humangenerated CO2 is absorbed by the oceans • CO2 dissolved in sea water forms carbonic acid; this process is called ocean acidification © 2011 Pearson Education, Inc. 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 organisms, including reef-building corals • BUT 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 (a) © 2011 Pearson Education, Inc. (b) (c)