H-B Woodlawn Biology - Jim’s Discussion Notes

THE CHEMICAL CONTEXT OF LIFE

Introduction

· Nature is not neatly packaged into individual sciences.

· While biologists specialize in the study of life, organisms and the world they live in are natural systems to which the basic concepts of chemistry and physics apply.

· An appreciation of chemistry is important to both microbiology and macrobiology studied by high school students.

· Life can be organized into a hierarchy of structural levels.

· At each successive level additional emergent properties appear.

A. Chemical Elements and Compounds

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.

· An element is a substance that cannot be broken down into other substances by chemical reactions.

· There are 92 naturally-occurring elements.

· Each element has a unique symbol, usually from the first one or two letters of the name, often from Latin or German.

• A compound is a substance consisting of two or more elements in a fixed ratio.

· Table salt (sodium chloride or NaCl) is a compound with equal numbers of chlorine and sodium atoms.

· While pure sodium is a metal and chlorine is a gas, their combination forms an edible compound, an emergent property.

2. Life requires about 25 chemical elements

· About 25 of the 92 natural elements are known to be essential for life.

· Four elements - carbon (C), oxygen (O), hydrogen (H), and nitrogen (N) - make up 96% of living matter.

· Most of the remaining 4% of an organism’s weight consists of phosphorus (P), sulfur (S), calcium (Ca), and potassium (K).

· Trace elements are required by an organism, but only in minute quantities.

· Some trace elements, like iron (Fe), are required by all organisms.

· Other trace elements are required only by some species.

· For example, a daily intake of 0.15 milligrams of iodine is required for normal activity of the human thyroid gland.

B. Atoms and Molecules

1. Atomic structure determines the behavior of an element

· Each element consists of unique atoms.

· An atom is the smallest unit of matter that still retains the properties of an element.

· Atoms are composed of even smaller parts, called subatomic particles.

· Two of these, neutrons and protons, are packed together to form a dense core, the atomic nucleus, at the center of an atom.

· Electrons form a cloud around the nucleus.

· Each electron has one unit of negative charge.

· Each proton has one unit of positive charge.

· Neutrons are electrically neutral.

· The attractions between the positive charges in the nucleus and the negative charges of the electrons keep the electrons in the vicinity of the nucleus.

· A neutron and a proton are almost identical in mass, about 1.7 x 10-²⁴ gram per particle.

· For convenience, an alternative unit of measure, the dalton, is used to measure the of mass subatomic particles, atoms or molecules.

· The mass of a neutron or a proton is close to 1 dalton.

· The mass of an electron is about 1/200th that of a neutron or proton.

· Therefore, we typically ignore the contribution of electrons when determining the total mass of an atom.

· All atoms of a particular element have the same number of protons in their nuclei.

· Each element has a unique number of protons, its unique atomic number.

· The atomic number is written as a subscript before the symbol for the element (for example, ₂He).

· Unless otherwise indicated, atoms have equal numbers of protons and electrons - no net charge.

· Therefore, the atomic number tells us the number of protons and the number of electrons that are found in a neutral atom of a specific element.

· The mass number is the sum of the number of protons and neutrons in the nucleus of an atom.

· Therefore, we can determine the number of neutrons in an atom by subtracting the number of protons (the atomic number) from the mass number.

· The mass number is written as a superscript before an element’s symbol (for example, ⁴He).

· The atomic weight of an atom, a measure of its mass, can be approximated by the mass number.

· For example, ⁴He has a mass number of 4 and an estimated atomic weight of 4 daltons.

· More precisely, its atomic weight is 4.003 daltons.

· While all atoms of a given element have the same number of protons, they may differ in the number of neutrons.

· Two atoms of the same element that differ in the number of neutrons are called isotopes.

· In nature, an element occurs as a mixture of isotopes.

· For example, 99% of carbon atoms have 6 neutrons (¹²C).

· Most of the remaining 1% of carbon atoms have 7 neutrons (¹³C) while the rarest isotope, with 8 neutrons, is ¹⁴C.

· Most isotopes are stable; they do not tend to lose particles.

· Both ¹²C and ¹³C are stable isotopes.

· The nuclei of some isotopes are unstable and decay spontaneously, emitting particles and energy.

· ¹⁴C is one of these unstable or radioactive isotopes.

· When ¹⁴C decays, a neutron is converted to a proton and an electron.

· This converts ¹⁴C to ¹⁴N, changing the identity of that atom.

· Radioactive isotopes have many applications in biological research.

· Radioactive decay rates can be used to date fossils.

· Radioactive isotopes can be used to trace atoms in metabolism.

· Radioactive isotopes are also used to diagnose medical disorders.

· For example, the rate of excretion in the urine can be measured after injection into the blood of known quantity of radioactive isotope.

· Also, radioactive tracers can be used with imaging instruments to monitor chemical processes in the body.

· While useful in research and medicine, the energy emitted in radioactive decay is hazardous to life.

· This energy can destroy cellular molecules.

· The severity of damage depends on the type and amount of energy that an organism absorbs.

· To gain an accurate perspective of the relative proportions of an atom, if the nucleus was the size of a golf ball, the electrons would be moving about 1 kilometer from the nucleus.

· Atoms are mostly empty space.

· When two elements interact during a chemical reaction, it is their electrons that are actually involved.

· The nuclei do not come close enough to interact.

· The electrons of an atom may vary in the amount of energy that they possess.

· Energy is the ability to do work.

· Potential energy is the energy that matter stores because of its position or location.

· Water stored behind a dam has potential energy that can be used to do work turning electric generators.

· Because potential energy has been expended, the water stores less energy at the bottom of the dam than it did in the reservoir.

· Electrons have potential energy because of their position relative to the nucleus.

· The negatively charged electrons are attracted to the positively charged nucleus.

· The farther electrons are from the nucleus, the more potential energy they have.

· However, electrons cannot occupy just any location away from the nucleus.

· Changes in potential energy can only occur in steps of a fixed amount, moving the electron to a fixed location.

· An electron cannot exist between these fixed locations.

· The different states of potential energy that the electrons of an atom can have are called energy levels or electron shells.

· The first shell, closest to the nucleus, has the lowest potential energy.

· Electrons in outer shells have more potential energy.

· Electrons can only change their position if they absorb or release a quantity of energy that matches the difference in potential energy between the two levels.

· The chemical behavior of an atom is determined by its electron configuration - the distribution of electrons in its electron shells.

· The first 18 elements, including those most important in biological processes, can be arranged in 8 columns and 3 rows.

· Elements in the same row use the same shells.

· Moving from left to right, each element has a sequential addition of electrons (and protons).

· The first electron shell can hold only 2 electrons.

· The two electrons of Helium fill the first shell.

· Atoms with more than two electrons must place the extra electrons in higher shells.

· For example, Lithium with three electrons has two in the first shell and one in the second shell.

· The second shell can hold up to 8 electrons.

· Neon, with 10 total electrons, has two in the first shell and eight in the second, filling both shells.

· The chemical behavior of an atom depends mostly on the number of electrons in its outermost shell, the valence shell.

· Electrons in the valence shell are known as valence electrons.

· Atoms with the same number of valence electrons have similar chemical behavior.

· An atom with a completed valence shell is unreactive.

· All other atoms are chemically reactive because they have incomplete valence shells.

· The paths of electrons are often visualized as concentric paths, like planets orbiting the sun.

· In reality, an electron occupies a more complex three-dimensional space, an orbital.

· The first shell has room for a single spherical orbital for its pair of electrons.

· The second shell can pack pairs of electrons into a spherical orbital and three p orbitals (dumbbell-shaped).

· The reactivity of atoms arises from the presence of unpaired electrons in one or more orbitals of their valence shells.

· Electrons preferentially occupy separate orbitals within the valence shell until forced to share orbitals.

· The four valence electrons of carbon each occupy separate orbitals, but the five valence electrons of nitrogen are distributed into three unshared orbitals and one shared orbital.

· When atoms interact to complete their valence shells, it is the unpaired electrons that are involved.

2. Atoms combine by chemical bonding to form molecules

· Atoms with incomplete valence shells interact by either sharing or transferring valence electrons.

· These interactions typically result in the atoms remaining close together, held by an attractions called chemical bonds.

· The strongest chemical bonds are covalent bonds and ionic bonds.

· A covalent bond is the sharing of a pair of valence electrons by two atoms.

· If two atoms come close enough that their unshared orbitals overlap, each atom can count both electrons toward its goal of filling the valence shell.

· For example, if two hydrogen atoms come close enough that their 1s orbitals overlap, then they can share the single electrons that each contributes.

· Two or more atoms held together by covalent bonds constitute a molecule.

· We can abbreviate the structure of this molecule by substituting a line for each pair of shared electrons, drawing the structural formula.

· H-H is the structural formula for the covalent bond between two hydrogen atoms.

· The molecular formula indicates the number and types of atoms present in a single molecule.

· H₂ is the molecular formula for hydrogen gas.

· Oxygen needs to add 2 electrons to the 6 already present to complete its valence shell.

· Two oxygen atoms can form a molecule by sharing two pairs of valence electrons.

· These atoms have formed a double covalent bond.

· Every atom has a characteristic total number of covalent bonds that it can form - an atom’s valence.

· The valence of hydrogen is 1.

· Oxygen is 2.

· Nitrogen is 3.

· Carbon is 4.

· Phosphorus should have a valence of 3, based on its three unpaired electrons, but in biological molecules it generally has a valence of 5, forming three single covalent bonds and one double bond.

· Covalent bonds can form between atoms of the same element or atoms of different elements.

· While both types are molecules, the latter are also compounds.

· Water, H₂O, is a compound in which two hydrogen atoms form single covalent bonds with an oxygen atom.

· This satisfies the valences of both elements.

· Methane, CH₄, satisfies the valences of both C and H.

· The attraction of an atom for the electrons of a covalent bond is called its electronegativity.

· Strongly electronegative atoms attempt to pull the shared electrons toward themselves.

· If electrons in a covalent bond are shared equally, then this is a nonpolar covalent bond.

· A covalent bond between two atoms of the same element is always nonpolar.

· A covalent bond between atoms that have similar electronegativities is also nonpolar.

· Because carbon and hydrogen do not differ greatly in electronegativities, the bonds of CH₄ are nonpolar.

· If the electrons in a covalent bond are not shared equally by the two atoms, then this is a polar covalent bond.

· The bonds between oxygen and hydrogen in water are polar covalent because oxygen has a much higher electronegativity than does hydrogen.

· Compounds with a polar covalent bond have regions that have a partial negative charge near the strongly electronegative atom and a partial positive charge near the weakly electronegative atom.

· An ionic bond can form if two atoms are so unequal in their attraction for valence electrons that one atom strips an electron completely from the other.

· For example, sodium with one valence electron in its third shell transfers this electron to chlorine with 7 valence electrons in its third shell.

· Now, sodium has a full valence shell (the second) and chlorine has a full valence shell (the third).

· After the transfer, both atoms are no longer neutral, but have charges and are called ions.

· Sodium has one more proton than electrons and has a net positive charge.

· Atoms with positive charges are cations.

· Chlorine has one more electron than protons and has a net negative charge.

· Atoms with negative charges are anions.

· Because of differences in charge, cations and anions are attracted to each other to form an ionic bond.

· Atoms in an ionic bonds need not have acquired their charge by electrons transferred with each other.

· Compounds formed by ionic bonds are ionic compounds or salts, like NaCl or table salt.

· The formula for an ionic compound indicates the ratio of elements in a crystal of that salt.

· Atoms in a crystal do not form molecules with a definitive size and number of atoms as in covalent bonds.

· Ionic compounds can have ratios of elements different from 1:1.

· For example, the ionic compound magnesium chloride (MgCl₂) has 2 chloride atoms per magnesium atom.

· Magnesium needs to lose 2 electrons to drop to a full outer shell, each chlorine needs to gain 1.

· Entire molecules that have full electrical charges are also called ions.

· In the salt ammonium chloride (NH₄Cl), the anion is Cl^- and the cation is NH₄ ⁺.

· The strength of ionic bonds depends on environmental conditions.

3. Weak chemical bonds play important roles in the chemistry of life

· Within a cell, weak, brief bonds between molecules are important to a variety of processes.

· For example, signal molecules from one neuron use weak bonds to bind briefly to receptor molecules on the surface of a receiving neuron.

· This triggers a momentary response by the recipient.

· Weak interactions include ionic bonds (weak in water), hydrogen bonds, and van der Waals interactions.

· Hydrogen bonds form when a hydrogen atom already covalently bonded to a strongly electronegative atom is attracted to another strongly electronegative atom.

· These strongly electronegative atoms are typically nitrogen or oxygen.

· Typically, these bonds result because the polar covalent bond with hydrogen leaves the hydrogen atom with a partial positive charge and the other atom with a partial negative charge.

· The partially positive charged hydrogen atom is attracted to negatively charged (partial or full) molecules, atoms, or even regions of the same large molecule.

· For example, ammonia molecules and water molecules link together with weak hydrogen bonds.

· In the ammonia molecule, the hydrogen atoms have partial positive charges and the more electronegative nitrogen atom has a partial positive charge.

· In the water molecule, the hydrogen atoms also have partial positive charges and the oxygen atom has a partial negative charge.

· Areas with opposite charges are attracted.

· Even molecules with nonpolar covalent bonds can have partially negative and positive regions.

· Because electrons are constantly in motion, there can be periods when they accumulate by chance in one area of a molecule.

· This creates ever-changing regions of negative and positive charge within a molecule.

· Molecules or atoms in close proximity can be attracted by these fleeting charge differences, creating van der Waals interactions.

· While individual bonds (ionic, hydrogen, van der Waals) are weak, collectively they have strength.

4. A molecule’s biological function is related to its shape

· The three-dimensional shape of a molecule is an important determinant of its function in a cell.

· The shape of a molecule is determined by the arrangement of electron orbitals that are shared by the atoms involved in the bond.

· When covalent bonds form, the orbitals in the valence shell rearrange.

· A molecule with two atoms is always linear.

· However, a molecule with more than two atoms has a more complex shape.

· For atoms with electrons in both s and p orbitals, the formation of a covalent bonds leads to hybridization of the orbitals to four new orbitals in a tetrahedron shape.

· In a water molecule the hybrid orbitals that oxygen shares with hydrogen atoms are spread in a V shape, at an angle of 104.5^o.

· A methane molecule (CH₄) has all four hybrid orbitals shared and has hydrogen nuclei at the corners of the tetrahedron.

· In larger molecules the tetrahedral shape of carbon bonded to four other atoms is often a repeating motif.

· Biological molecules recognize and interact with one another based on molecular shape.

· For example, signal molecules from a transmitting brain cell have specific shapes that fit together with the shapes of receptor molecules on the surface of the receiving cell.

· The temporary attachment of the receptor and signal molecule stimulates activity in the receptor cell.

· Molecules with similar shapes can interact in similar ways.

· For example, morphine, heroin, and other opiate drugs are similar enough in shape that they can bind to the same receptors as natural signal molecules, called endorphins.

· Binding to the receptors produces euphoria and relieves pain.

5. Chemical reactions make and break chemical bonds

· In chemical reactions, chemical bonds are broken and reformed, leading to new arrangements of atoms.

· The starting molecules in the process are called reactants and the end molecules are called products.

· In a chemical reaction, all of the atoms in the reactants must be accounted for in the products.

· The reactions must be “balanced.”

· For example, we can recombine the covalent bonds of H₂ and O₂ to form the new bonds of H₂O.

· In this reaction, two molecules of H₂ combine with one molecule of O₂ to form two molecules of H₂O.

· The ratios of molecules are indicated by coefficients.

· Photosynthesis is an important chemical reaction.

· Green plants combine carbon dioxide (CO₂) from the air and water (H₂O) from the soil to create sugar molecules and molecular oxygen (O₂), a byproduct.

· This chemical reaction is powered by sunlight.

· Humans and other animals depend on photosynthesis for food and oxygen.

· The overall process of photosynthesis is

· 6CO₂ + 6H₂O -> C₆H₁₂O₆ + 6O₂

· This process occurs in a sequence of individual chemical reactions.

· Some chemical reactions go to completion; that is, all the reactants are converted to products.

· Most chemical reactions are reversible, the products in the forward reaction becoming the reactants for the reverse reaction.

· For example in this reaction: 3H₂ + N₂ <=> 2NH₃ hydrogen and nitrogen molecules combine to form ammonia, but ammonia can decompose to hydrogen and nitrogen molecules.

· Initially, when reactant concentrations are high, they frequently collide to create products.

· As products accumulate, they collide to reform reactants.

· Eventually, the rate of formation of products is the same as the rate of breakdown of products (formation of reactants) and the system is at chemical equilibrium.

· At equilibrium, products and reactants are continually being formed, but there is no net change in the concentrations of reactants and products.

· At equilibrium, the concentrations of reactants and products are typically not equal, but their concentrations have stabilized.