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Chapter 2: The Chemical Context of Life AP Biology Overview: A Chemical Connection to Biology • Biology is a multidisciplinary science • Living organisms are subject to basic laws of physics and chemistry • • Ex) The use of formic acid by ants to maintain “devil’s gardens,” stands of Duroia trees • Some patches of forest are almost completely dominated by this single plant species • Researchers showed that ants prevent other plant species from growing in these” gardens” by injecting intruders with the poisonous chemical formic acid This chapter focuses on the chemical components and their interactions that make up all matter • Elements and compounds • Atomic structure • Chemical bonding and the formation of molecules • Chemical reactions that make and break chemical bonds Concept 2.1: 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 • Ex) Rocks, metals, oils, gases, humans • Matter is made up of elements – An element is a substance that cannot be broken down to other substances by chemical reactions • Scientists recognize 92 naturally-occurring elements – Ex) Gold, copper, carbon, oxygen • Each element has a symbol, usually the first letter or two of its name – Some symbols are derived from Latin or German • Ex) Sodium is “Na” from the Latin word natrium • A compound is a substance consisting of two or more different elements in a fixed ratio – A compound has characteristics different from those of its elements Fig. 2-3 • Ex) Table salt (NaCl) – pure sodium is a metal; pure chlorine is a poisonous gas Sodium Chlorine Sodium chloride Essential Elements of Life • About 25 of the 92 elements are essential to life – • Only 4 of these elements together make up 96% of living matter: • Carbon (18.5%) • Hydrogen (9.5%) • Oxygen (65%) • Nitrogen (3.3%) Most of the remaining 4% consists of: – Calcium (1.5%) – Phosphorus (1.0%) – Potassium (0.4%) – Sulfur (0.3%) – Sodium (0.2%) – Chlorine (0.2%) – Magnesium (0.1%) Essential Elements of Life • Trace elements are those required by an organism in minute quantities – Some trace elements are needed by all forms of life • – Ex) Iron (Fe) Other trace elements are required by only certain species • Ex) Iodine (I) is an essential ingredient of a hormone produced by the thyroid gland in all vertebrates – Iodine deficiency causes a condition called goiter, in which the thyroid gland grows to an abnormal size Concept Check 2.1 • 1) Explain how table salt has emergent properties. • 2) Is a trace element an essential element? Explain. • 3) Iron (Fe) is a trace element required for the proper functioning of hemoglobin, the molecule that carries oxygen in RBCs. What might be the effects of an iron deficiency? Concept 2.2: An element’s properties depend on the structure of its atoms • Each element consists of unique atoms – An atom is the smallest unit of matter that still retains the properties of an element • Atoms are symbolized with the same abbreviation used for the element made up of those atoms – Ex) “C” represents both the element carbon and a single carbon atom • Atoms are composed of subatomic particles – Relevant subatomic particles include: • Neutrons (no electrical charge) • Protons (positive charge) • Electrons (negative charge) • Neutrons and protons are packed tightly together at the center of an atom to form the atomic nucleus • • Protons give the nucleus a positive charge Electrons form a cloud around the nucleus • Attraction between opposite charges keeps these electrons in the vicinity of the nucleus Fig. 2-5 Cloud of negative charge (2 electrons) Electrons Nucleus (a) (b) • Neutron mass and proton mass are almost identical (~1.7 x 10 –24 grams) • Because the mass of these particles is so miniscule, they are measured in daltons • The dalton is the same unit of measure as the atomic mass unit (amu) • Neutrons and protons both have masses close to 1 dalton • Electron mass can be ignored when computing atomic mass because it is only ~1/2000 that of a proton or neutron Atomic Number • Atoms of the various elements differ in number of subatomic particles – All atoms of a particular element have a unique number of protons, known as the atomic number, in their nucleus • The atomic number is written as a subscript to the left of the symbol for that element – Ex) 2He = Helium (2 protons) – Unless otherwise indicated, an atom contains the same number of protons and electrons, making it electrically neutral Atomic Mass • An element’s mass number is the sum of protons plus neutrons in the nucleus – The mass number is written as a superscript to the left of the element’s symbol – Because the atomic number indicates how many protons are present, the number of neutrons can be determined by subtracting the atomic number from the mass number • • Ex) = Helium (2 protons; 4 - 2 = 2 neutrons) Atomic mass, the atom’s total mass, can be approximated by the mass number – This is because the contribution of electrons to mass is negligible • Ex) Sodium ( ) is 22.9898, or ~23 daltons Isotopes • All atoms of an element have the same number of protons but may differ in number of neutrons – Isotopes are two atoms of an element that differ in number of neutrons • Because neutrons contribute to atomic weight, isotopes have a different mass than that of the element – Despite their different masses, isotopes behave identically in chemical reactions – In nature, an element occurs as a mixture of its isotopes – • Ex) There are 3 isotopes of carbon: carbon-12 (99%) and carbon-13 and carbon-14 (1%) The number usually given as the atomic mass of an element is actually an average of the atomic masses of all the element’s naturally occurring isotopes Isotopes • Some isotopes, known as radioactive isotopes, are unstable and have a tendency to lose particles – When this decay eventually leads to a change in the number of protons, the atom is transformed to an atom of a different element • – Ex) Radioactive carbon decays to form nitrogen Radioactive isotopes have many useful applications • Measurements of radioactivity in fossils can be used to date them Fig. 2-7 • Radioactive isotopes are useful as tracers to follow atoms through metabolism – Cells use radioactive atoms in the same manner as they would nonradioactive isotopes of the same element, but they can be easily detected • Cancerous throat tissue Ex) Biologists can detect cancer by using radioactive tracers to monitor the synthesis of DNA and looking for “hot spots” of activity The Energy Levels of Electrons • When two atoms approach each other during a chemical reaction, only their electrons are involved directly • An atom’s electrons vary in the amount of energy they posses • • Energy is the capacity to cause change The electrons of an atom have a type of energy called potential energy because of how they are arranged in relation to the nucleus • • Potential energy is the energy that matter has because of its location or structure The negatively-charged electrons are attracted to the positively-charged nucleus • As a result, it takes work to move an electron farther away from the nucleus • This means that more distant electrons have greater potential energy The Energy Levels of Electrons • Changes in the potential energy of electrons can occur only in steps of fixed amounts, similar to a ball on a staircase • The ball can have different amounts of potential energy, depending on which step it is on, but it cannot spend much time between steps • Similarly, an electron cannot exist in between energy levels (a) A ball bouncing down a flight of stairs provides an analogy for energy levels of electrons The Energy Levels of Electrons • An energy level, or electron shell, is an electron’s state of potential energy • An electron’s energy level or shell is correlated with its average distance from the nucleus • In diagrams, shells can be represented by concentric circles • Fig. 2-8 The first shell is closest to the anddown contains (a) Anucleus ball bouncing a flight electrons with the of stairs provides an analogy for energy levels of electrons lowest potential energy • Electrons in the higher shells have increasingly higher amounts of potential energy Third shell (highest energy level) Second shell (higher energy level) First shell (lowest energy level) (b) Atomic nucleus Energy absorbed Energy lost The Energy Levels of Electrons Fig. 2-8 • (a) A ball bouncing down a flight of stairs provides an analogy for energy levels of electrons An electron can change the shell it occupies by absorbing or losing an amount of energy equal to the difference in potential energy between its position in the old shell and that in the new shell Third shell (highest energy level) • When an electron absorbs energy, it moves to a shell farther out from the nucleus • • Second shell (higher energy level) First shell (lowest energy level) Ex) Light energy can excite an (b) electron to a higher energy level, as seen in the first step of photosynthesis Energy absorbed Energy lost Atomic nucleus When an electron loses energy, it “falls back” to a shell closer to the nucleus, with the lost energy usually released to the environment as heat • Ex) The hot temperature of a car’s surface on a sunny day is the result of sunlight exciting electrons at the car surface to higher energy levels, and then those electrons falling back to their original levels Electron Distribution and Chemical Properties • The chemical behavior of an atom is determined by the distribution of electrons in the atom’s electron shells • The periodic table of the elements shows the electron distribution for each element • The elements are arranged in rows, known as periods, that correspond to the number of electron shells in their atoms • • Ex) Nitrogen = 2nd row = 2 electron shells The left-to-right sequence of elements in each row corresponds to the sequential addition of electrons and protons • Ex) Oxygen has +1 electrons and protons than nitrogen • Electrons (like all matter) tend to exist in the lowest available state of potential energy – In an atom, the lowest such state is in the 1st shell • The 1st shell, however, can only hold 2 electrons – • Atoms with more than 2 electrons must therefore use higher shells because the 1st shell is full The 2nd shell holds a max of 8 electrons – Atoms with more than 10 electrons (2 in the 1st, 8 in the 2nd) must therefore begin filling the next energy level • The chemical behavior of an atom depends mostly on the number of electrons in its outermost shell – Electrons in the outermost shell are called valence electrons • – The outermost electron shell in called the valence shell • • Ex) Lithium’s valence shell is its 2nd shell Atoms with the same number of electrons in their valence shells exhibit similar chemical behavior – • Ex) Lithium has 1 valence electron Ex) Fluorine and chlorine both have 7 valence electrons; both form compounds when combined with the element sodium An atom with a completed valence shell is chemically unreactive, or inert – Such atoms are found at the far right of the periodic table (He, Ne, Ar) Electron Orbitals • • The electron shells of atoms are often symbolized as 2-dimensional concentric-circle diagrams – Each concentric circle, however, represents only the average distance between an electron in that shell and the nucleus – Thus, the exact location of an electron in an atom can never be known We can instead describe the space in which an electron spends most of its time – An orbital is the threedimensional space where an electron is found 90% of the time Electron Orbitals • Orbitals can be thought of as components of an electron shell – There are a specific number of orbitals of distinctive shapes and orientations in each electron shell • The 1st electron shell has only 1 spherical s orbital, called 1s • The 2nd shell has 4 orbitals: – One large spherical s orbital, called 2s – • Three dumbbellshaped p orbitals, called 2p orbitals The 3rd shell and other higher electron shells also have s and p orbitals, along with orbitals of more complex shapes Electron Orbitals • • No more than 2 electrons can occupy a single orbital – The 1st electron shell (1 orbital) can therefore only hold 2 electrons – The 2nd electron shell (4 orbitals) can hold up to 8 electrons The reactivity of atoms arises from the presence of unpaired electrons in one or more orbitals of their valence shells – Atoms interact in such a way that completes their valence shells, which involves finding partners for their unpaired electrons Concept Check 2.2 • 1) A lithium atom has 3 protons and 4 neutrons. What is its atomic mass in daltons? • 2) A nitrogen atom has 7 protons, and the most common isotope of nitrogen has 7 neutrons. A radioactive isotope of nitrogen has 8 neutrons. Write the atomic number and mass number of this radioactive nitrogen. • 3) How many electrons does fluorine have? How many electron shells? Name the orbitals that are occupied. How many electrons are needed to fill the valence shell? • 4) If 2 or more elements are in the same row, what do they have in common? If 2 or more elements are in the same column, what do they have in common (see Figure 2.9, pp. 36)? Concept 2.3: The formation and function of molecules depend on chemical bonding between atoms • Atoms with incomplete valence shells can interact with other atoms in such a way that completes their valence shell, either by: • Sharing valence electrons • Transferring valence electrons • These interactions usually result in atoms staying close together, held by attractions called chemical bonds • The strongest kinds of chemical bonds are: • Covalent bonds • Ionic bonds Covalent Bonds • A covalent bond is the sharing of a pair of valence electrons by two atoms – When 2 atoms approach one another, their valence orbitals overlap and allow them to share electrons – In a covalent bond, the shared electrons count as part of each atom’s valence shell – Two or more atoms held together by covalent bonds form a molecule • Ex) H2: the 1s orbitals (both containing 1 electron) overlap to complete the valence shell with 2 electrons • Molecules can be represented in different ways: – Lewis dot structure: uses element symbols, with dots representing the outermost valence electrons • Ex) H:H – Electron distribution diagram: shows the locations of all electrons within their appropriate orbitals, along with overlapping valence orbitals containing shared electrons – Space-filling model: 3-D model that comes closest to representing the actual shape of a molecule • Molecules can be represented in different ways: – Structural formula: uses element symbols, with lines to represent the sharing of pairs of electrons – – Ex) H-H • One line represents a single bond, which is the sharing of one pair of valence electrons • Two lines represent a double bond, which is the sharing of two pairs of valence electrons Molecular formula: indicates the type and number of atoms present in a molecule but not electron sharing • Ex) H2 • Atoms that can share valence electrons have a bonding capacity that corresponds to the number of covalent bonds they can form – This bonding capacity is called an atom’s valence – It usually equals the number of unpaired electrons required to complete the atom’s outermost (valence) shell • Ex) Oxygen, with 6 electrons in its outermost shell, has a valence of 2 • Covalent bonds can form between atoms of the same element or atoms of different elements – Molecules consisting of the same types of atoms are known as pure elements • Ex) H2 – A compound is a combination of two or more different elements • Ex) H2O • The attraction of a particular kind of atom for the electrons of a covalent bond is called its electronegativity – The more electronegative an atom, the more strongly it pulls shared electrons toward itself • In a covalent bond between 2 atoms of the same element, both atoms have equal electronegativity and thus share electrons equally – This type of bond in known as a nonpolar covalent bond • Ex) O2 • In compounds where one atom is bonded to a more electronegative atom, the electrons are not shared equally – This type of bond in known as a polar covalent bond • Ex) H2O • Unequal sharing of electrons causes a partial positive or negative charge for each atom or molecule – Electrons spend more time near the more electronegative atom, leading to a partial negative charge (δ-) – As a result, the less electronegative atom has a partial positive charge (δ+) Ionic Bonds • In some cases, 2 atoms are so unequal in electronegativity that the more electronegative atom strips a valence electron completely away from its partner – • Ex) the transfer of an electron from sodium (11Na) to chlorine (17Cl) After the transfer of an electron, both atoms have charges – A charged atom (or molecule) is called an ion • • Depending on its charge, an ion can be classified as: – A cation: a positively charged ion – An anion: a negatively charged ion An ionic bond is an attraction between an anion and a cation – The transfer of an electron is not the formation of a bond itself; rather it simply allows an ionic bond to form – Ions participating in an ionic bond with each other need not have acquired their charge by electron transfer with each other Animation: Ionic Bonds • Compounds formed by ionic bonds are called ionic compounds, or salts – Salts (ex: sodium chloride (table salt)) are often found in nature as crystals of various sizes and shapes – Each salt crystal is an aggregate of many cations and anions bonded by their electrical attraction and arranged in a 3-D lattice – Ionic compounds do not consist of molecules having a definite size and number of atoms • The formula for an ionic compound (ex: NaCl) indicates only the ratio of elements in a crystal of the salt • Not all salts have equal numbers of cations and anions – Ex) MgCl2 has 2 chloride ions for each magnesium ion; magnesium must lose 2 valence electrons to obtain a complete valence shell, making it a cation with a net charge of 2+ • – One magnesium cation can therefore form ionic bonds with 2 chloride ions The term ion also applies to entire molecules that are electrically charged • Ex) In the salt ammonium chloride (NH4Cl), an ion with an electrical charge of 1+, the anion is a single chloride ion (Cl-) and the cation is ammonium (NH4+) Weak Chemical Bonds • In organisms, most of the strongest chemical bonds are covalent bonds – • The bonds link atoms to form cell’s molecules Weaker bonding within and between molecules is also essential to the cell • – Ex) Ionic and hydrogen bonds Most important large biological molecules are held in their functional form by weak bonds – In addition, when 2 molecules in a cell make contact, they may adhere temporarily by weak bonds • This reversibility of weak bonding is advantageous because 2 molecules can come together, respond to one another in some way, and then separate Hydrogen Bonds • 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 Van der Waals Interactions • Even molecules with nonpolar covalent bonds may have positively and negatively charged regions – Electrons are not always evenly distributed in such molecules • At any instant, they may accumulate by chance in one part of the molecule or another – This results in continually changing regions of positive and negative charge that allow all atoms and molecules to stick together • These are called van der Waals interactions Van der Waals Interactions • Van der Waals interactions are weak and occur only when atoms or molecules are very close together – Despite their weakness, they have recently been shown to be responsible for ability of gecko lizard to walk up a wall • Each gecko has 100s of 1000s tiny hairs with multiple projections at hair’s tip that increase surface area • These hairs are so numerous that, despite their individual weakness, together they can support the gecko’s body weight Molecular Shape and Function • Together, van der Waals interactions, hydrogen bonds, ionic bonds in water, and other weak bonds may form not only between molecules but also between different regions of a single large molecule, such as a protein – Though these bonds are individually weak, their cumulative effect is to reinforce 3-D shape of a large molecule – The characteristic shape and size of a molecule is very important to its function • A molecule’s shape is determined by the positions of its atoms’ valence orbitals – When an atom forms covalent bonds, the orbitals in the valence shell rearrange • Molecules made up of only 2 atoms (H2, O2) are always linear • Molecules with more than 2 atoms have more complicated shapes because the s and p orbitals may hybridize – For atoms with valence electrons in both s and p orbitals, the single s and 3 p orbitals hybridize to form 4 new hybrid orbitals shaped like identical teardrops extending from the atomic nucleus • If we connect the larger ends of the teardrops with lines, we have the outline of a geometric shape called a tetrahedron (similar to a pyramid) • 3 models representing molecular shape are shown for water and methane – Water: 2 of the hybrid orbitals in the oxygen’s valence shell are with hydrogen atoms • – • This results in a molecule shaped roughly like a V, with its 2 covalent bonds spread apart at an angle of 104.5 degrees Methane: has shape of a completed tetrahedron because all 4 hybrid orbitals of carbon are shared with hydrogen atoms • Nucleus of carbon is at the center • 4 covalent bonds radiate to hydrogen nuclei at corners of tetrahedron -Larger molecules containing multiple carbon atoms have a more complex overall shape – However, the tetrahedral shape of carbon bonded to 4 other atoms is a common theme in many other the molecules that make up living matter • Molecular shape is important in biology because it determines how biological molecules recognize and respond to one another with specificity – Only molecules with complementary shapes can form weak bonds with one another – Molecules with similar shapes can have similar biological effects • Ex) We can see this specificity in the effects of opiates (drugs derived from opium) • Opiates relieve pain and alter mood by binding to specific receptor molecules on surface of brain cells (Why would brain cells carry receptors for opiates, compounds NOT made by our bodies?) – The discovery of endorphins answered this question in 1975 • Endorphins are signaling molecules made by pituitary gland that bind to receptors, relieving pain and producing euphoria during times of stress (intense exercise) – Opiates have shapes similar to endorphins (see boxed portions) and mimic them by binding to endorphin receptors in brain • This is why opiates and endorphins have similar effects Concept Check 2.3 • 1) Why does the following structure fail to make sense chemically? H-C=C-H • 2) Explain what holds together the atoms in a crystal of magnesium chloride (MgCl2). • 3) If you were a pharmaceutical researcher, why would you want to learn the three-dimensional shapes of naturally occurring signaling molecules? Concept 2.4: Chemical reactions make and break chemical bonds • The making and breaking of chemical bonds, which leads to changes in the composition of matter, are called chemical reactions – When we write a chemical reaction, we use an arrow to indicate the conversion of the starting material (reactants) to the final molecules (products) • -Coefficients indicate the number of molecules involved – Ex) Reaction between hydrogen and oxygen that forms water • This reaction breaks covalent bonds of H2 and O2 and forms the new bonds of H2O – The coefficient 2 in front of H2 means that the reaction starts with 2 molecules of hydrogen • Notice that all the atoms of the reactants (amounts) must be accounted for in the products (matter is conserved in a chemical reaction) – Stoichiometry – Reactions cannot create or destroy matter – they can only rearrange it! • Photosynthesis is another important chemical reactions – The raw materials of photosynthesis are carbon dioxide (taken from air) and water (absorbed from soil) • Within plant cells, sunlight powers the conversion of these ingredients to sugar (glucose) and oxygen molecules (released as by-product to surroundings) – Although photosynthesis is actually a sequence of many chemical reactions, we still end up with the same number and kinds of atoms we had at the start • Matter has simply been rearranged, with an input of energy provided by sunlight 6 CO2 + 6 H20 → C6H12O6 + 6 O2 • All chemical reactions are reversible, with products of forward reaction becoming reactants for reverse reaction – Ex) Hydrogen and nitrogen molecules can combine to form ammonia (NH3), but ammonia can also decompose to regenerate hydrogen and nitrogen • 2 opposite-headed arrows indicate that the reaction is reversible: 3H2 + N2 – 2 NH3 One factors affecting the rate of a reaction is the concentration of reactants • Greater concentration = more frequent collisions of reactant molecules with one another = more opportunities to react and form products – As products accumulate, however, collisions resulting in reverse reaction become more frequent • Eventually, forward and reverse reactions occur at same rate, so relative concentrations of products and reactants stop changing , a condition known as chemical equilibrium • Chemical equilibrium is dynamic – Reactions are still going on, but with no NET effect on concentrations of reactants and products – Equilibrium does NOT mean that reactants and products are equal in concentration, but that their concentrations have stabilized at a particular ratio • In some chemical reactions, equilibrium point may lie so far to the right (toward products), that these reactions go essentially to completion (virtually all reactants are converted to products) Concept Check 2.4 • 1) Refer to the reaction between hydrogen and oxygen that forms water, shown with ball-and-stick models on pp. 42. Draw the Lewis dot structures representing this reaction. • 2) Which types of chemical reactions occur fastest at equilibrium: the formation of products from reactants or reactants from products? • 3) Write an equation that uses the products of photosynthesis as reactants and uses the reactants as products. Add energy as another product. This new equation describes a process that occurs in your cells. Describe this equation in words. How does this equation relate to breathing? You should now be able to: 1. Identify the four major elements 2. Distinguish between the following pairs of terms: neutron and proton, atomic number and mass number, atomic weight and mass number 3. Distinguish between and discuss the biological importance of the following: nonpolar covalent bonds, polar covalent bonds, ionic bonds, hydrogen bonds, and van der Waals interactions Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings