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Chemistry for BIOS 302 Life is Chemistry • We believe that life is just a large set of chemical reactions. Living organisms obey the same physical and chemical principles that non-living things do. – The scientific world has long since rejected the theory of vitalism: the idea that a nonmaterial, spiritual force is needed to convert non-living into living. • Chemistry of carbon compounds = organic chemistry • In aqueous (watery) environment, in a very narrow temperature range. • The chemical reactions of life are very complex and intertwined. • Mostly involve large molecules (macromolecules) composed of many subunits. Elements and Atoms • • • • All matter is composed of atoms. Elements such as carbon and oxygen are a group of atoms of the same type. There are 92 naturally occurring elements, plus about 25 artificially-created elements. Living things are mainly composed of the elements carbon, hydrogen, oxygen, and nitrogen. Almost as important are sulfur (found in 2 of the 20 amino acids) and phosphorus (an essential component of all nucleic acids). Another dozen or so elements are also used: iron, magnesium, calcium, sodium, potassium, chlorine, to name a few. Atoms • Atoms have 3 components: protons, neutrons, and electrons • protons and neutrons are in the nucleus • Electrons circle around the nucleus • NOTE! This nucleus is NOT the same as the cell’s nucleus Properties of Protons, Neutrons, and Electrons • Protons have a mass of 1 dalton (or 1 atomic mass unit). Protons also have an electrical charge of +1: protons are positively charged. – Daltons are defined as 1/12 the mass of a carbon-12 atom. This averages out the small weight difference between protons and neutrons. – Daltons are related to grams by Avogadro’s number: 6 x 1023 protons have a mass of 1 gram. • Neutrons also have a mass of 1 dalton. Most of the mass of an atom is in its protons and neutrons. Neutrons are neutral (not electrically charged). • Electrons weigh about 1/1800 of a dalton, very light. They have a -1 charge (negatively charged). Atomic Number and Weight • The atomic number of an atom is the number of protons it has. The atomic number defines the basic identity of the atom: what element it is. – For example. All hydrogen atoms have 1 proton (atomic number = 1), and all carbon atoms have 6 protons (atomic number = 6). All gold atoms have 79 protons; uranium atoms have 92 protons. • The atomic weight of an atom is the number of protons plus the number of neutrons. – For example, most carbon atoms have 6 protons and 6 neutrons, so the atomic weight is 12. – Electrons are very light, so they don’t add significantly to the atomic weight. Isotopes • • • • Isotopes are atoms with the same atomic number (same element) but different atomic weights: that is, isotopes have the same number of protons but different numbers of neutrons. Most elements have several naturally occurring isotopes. Different isotopes of the same element all have the same chemical properties. Chemical properties are determined by the number of protons only. Heavier isotopes (too many neutrons) are usually radioactive: the atomic nucleus is unstable and spontaneously changes into a different configuration of protons and neutrons. In the process, particles or electromagnetic radiation are released. Different isotopes are written like 12C or 14C. Molecular Weight, Moles and Molarity • The molecular weight of a chemical compound is the sum of the atomic weights of the atoms in the compound. – For example, carbon dioxide (CO2) has 1 carbon (12 daltons) plus 2 oxygens (2 * 16 = 32 daltons), so it has a molecular weight of 44 daltons. – Note that atomic weights are usually not whole numbers: the fraction comes from averaging the atomic weights of all the isotopes present in the Earth’s crust. • Molecular weight can also be thought of as grams per mole (g/mol). A mole of carbon dioxide has a mass of 44 grams. • How fast chemical reactions occur depends on the concentration of the reactants. Most biological chemistry occurs in aqueous solution, so we care about concentration in terms of moles per liter (molar, symbolized as M). – A 1 M solution contains 1 mole of the solute dissolved in a total volume of 1 liter. • For example, what is the concentration of 16 g of CO2 in 0.5 L? – 16 g / 44 g/mol = 0.36 moles. – 0.36 mol / 0.5 L = 0.72 mol/L = 0.72 M Electron Shells • • • • Most chemistry is caused by the electrons of the outermost shell: atoms have a “desire” to have a full outer (valence) shell with 8 electrons in it, so they share electrons to accomplish this. • • Atoms are usually neutral in electrical charge, which means that they have the same number of electrons (- charge) as protons (+ charge). Electrons circle the nucleus at defined positions called shells. The innermost shell of every atom holds 2 electrons. The next two shells hold up to 8 electrons. The first three shells thus hold 2 + 8 + 8 = 18 electrons, which corresponds to argon. – Things get more complex above this point, but nearly all the biologically important elements use only the first 3 electron shells. – Big exception: metals such as iron, zinc, manganese, copper. Electrons fill in from the nucleus out, so the inner shells are always full (except hydrogen, which only has 1 electron). In most atoms, the outermost shell (the valence shell) is not full. Ions • The number of electrons in an atom is usually the same as the number of protons. Since electrons have a -1 charge and protons have a +1 charge, atoms are electrically neutral. • Ions are atoms where the number of electrons is different from the number of protons. • Ions have an electrical charge: positive charge if more protons than electrons, and negative charge if more electrons than protons. • Cations are positively charged: Na+ and Mg2+ are examples. • Anions are negatively charged: Cl- for example. – The prefix “an-” means “not” or “negative”: anonymous = no name; anarchy = no government; anhydrous = no water. For example, sodium atoms have 11 protons. Neutral sodium atoms also have 11 electrons, but sodium ions have only 10 electrons: 11 protons plus 10 electrons, giving a total charge of +1. Chemical Bonds • Atoms can combine with each other to form molecules. Very few elements exist naturally in an uncombined state: mostly they are joined into molecules. • A molecule is a defined number of atoms grouped into a defined spatial relationship. For example, water, H2O, is 2 hydrogen atoms connected to an oxygen atom. The oxygen is in the middle, and the hydrogens are attached at an angle to it. The bonds have specific lengths and angles relative to each other. • A large group of the same molecule is called a compound (just as a large group of the same atom is called an element). • Molecules are held together by chemical bonds. • Chemical bonds are the result of 2 forces: • 1. The octet rule, which means that atoms want to have 8 electrons in their outer shell (2 in the case of hydrogen). • 2. The attraction between atoms of opposite electrical charge. • The three main types of chemical bond are: ionic bond, covalent bond, and hydrogen bond. Outer Shell Electrons • Hydrogen has 1 electron in its outer shell, and needs a total of 2 electrons to be stable. • All other biologically relevant atoms need 8 electrons in their outer shell – Thus, carbon has 4 electrons and wants 4 more, nitrogen has 5 electrons and wants 3 more, and oxygen has 6 and wants 2 more. Ionic Bonds • In an ionic bond, one atom gives an electron to another atom. This makes both atoms ions, and they are held together because their opposite charges attract each other. • In sodium chloride (table salt), sodium starts out with 1 electron in its outer shell. The next shell down has 8 electrons, so by giving 1 electron away, the sodium atom gets a full outer shell. It then has a +1 charge. • Chlorine starts out with 7 electrons in its outer shell. By gaining one more electron, it gets 8 in the outer shell, and a -1 charge. • The + charged sodium and the – charged chlorine attract each other, and they pack together in salt crystals. Ionic Bonds in Biology • Ionic bonds are very strong in many environments. But in an aqueous environment, they are greatly weakened by interactions with water molecules, and are much weaker than covalent bonds. – However, unlike covalent bonds, ionic bonds don’t have a fixed geometry. They are equally strong regardless of how the 2 atoms are oriented. • • In biology, ionic bonds are mainly used to stabilize the three-dimensional shapes of molecules and to help molecules interact with each other. We usually speak of electrostatic interactions between charged molecules (both attraction and repulsion), rather than use the term “ionic bond”. – Hydrogen bonds are another example of electrostatic interactions that we will discuss soon. Covalent Bonds • Covalent bonds occur when 2 atoms share a pair of electrons. The electrons spend part of their time with both atoms, so the octet rule is satisfied sufficiently. • A molecule of hydrogen gas, H2, has 2 hydrogen atoms. Each atom provides 1 electron, so in the bond each atom shares 2, a complete shell for hydrogen. – The bond is symbolized as a line connecting the 2 H’s: H-H • Covalent bonds are what hold most biological molecules together. Covalent Bond Orientation • The 8 electrons of the outer electron shell arrange themselves into 4 groups of two, which stay as far apart as possible. This produces a tetrahedron shape (triangular pyramid). Covalently bonded molecules mostly fit this pattern. • In water (H2O), the oxygen has 6 electrons in its outer shell, and it shares one with each of the 2 hydrogens, giving 8 shared electrons for oxygen and 2 for each hydrogen. The hydrogens form 2 corners of the tetrahedron, and the two pairs of unbonded electrons form the other 2 corners. Single, Double, and Triple Bonds • In a single bond, a pair of electrons (2 electrons) is shared. H2 gas and water are examples of this. Most covalent bonds are single bonds. • In a double bond, 2 pairs of electrons (total of 4 electrons shared) are shared. In oxygen gas (O2), each atom has 6 electrons, a total of 12. Each atom contributes 2 electrons to the bond, so 4 are shared. Thus each atom has 4 unshared and 4 shared electrons, satisfying the octet rule. Carbon dioxide is another example. • In a triple bond, 3 pairs of electrons (6 electrons) are shared. There are very few triple bonds in biological molecules: nitrogen gas and cyanide (CN) are examples. Double Bonds Give Planar Configuration • Double bonds are shorter and stronger than single bonds. – • • Chemical groups attached to single bonds can rotate freely around the bond axis. Groups attached to double bonds are forced into a planar configuration and can’t rotate. – – – • Triple bonds are even shorter and stronger. This gives rise to cis and trans isomers for compounds with double bonds. Cis isomers have both side groups coming out of the same side of the double bond; trans isomers have the side groups coming out opposite sides. Here, cis-2-butene has different chemical properties than trans-2-butene. This is important with fatty acids, where artificially created trans isomers have been shown to be unhealthy relative to the naturally occurring cis isomers. Carbon dioxide is non-polar because the two double bonds between C and O create a linear molecule. The partial negative charges on the oxygens cancel out. Polar Covalent Bonds • Sometimes the electrons in a covalent bond aren’t shared equally, because one atom attracts electrons more strongly than the other. – Formally: the atom that attracts electrons more strongly has a greater electronegatvity. • When this happens, the electrons spend more time with one atom, and that atom becomes slightly negatively charged. The other atom becomes slightly positively charged. This is a polar covalent bond, because the atoms form positive and negative poles. • Rule: Carbon and hydrogen share electrons equally. Oxygen and nitrogen also share equally. But, oxygen and nitrogen attract electrons more strongly than carbon or hydrogen. – That is, oxygen and nitrogen are more electronegative than carbon and hydrogen. – Or: in a covalent bond, an oxygen or nitrogen will have a slightly negative charge, and carbon or hydrogen will have a slightly positive charge. • Water is a polar compound, because the oxygen is slightly negative and the hydrogens slightly positive. • – Note that the total charge on the molecule is balanced, same number of electrons as protons, but within the molecule the charges are slightly separated. Bonds where the electrons are shared equally are called non-polar. The C-H bond is non-polar. Hydrogen Bonds • The slight + and – charges in polar bonds attract each other. In biological molecules, it is common for the partial + charge on a hydrogen to attract the partial – charge on a nearby oxygen or nitrogen. This attraction is called a hydrogen bond. • Hydrogen bonds are very weak compared to covalent bonds, but large numbers of them can add up to a strong bond. The strands of DNA are held together by hydrogen bonds. • Hydrogen bonds also form between different parts of the same molecule, and between water and other molecules. • NaCl dissolves as H-bonds form between water molecules and the Na+ and Cl- ions. Water molecules surround each ion completely, separating the ions from the salt crystal and dispersing them throughout the water. More on Bonds • Covalent bonds are by far the strongest bonds for biological molecules. • Ionic bonds are really just extreme cases of polar bonds: the sharing is so unequal that one atom never gets the electron. • Hydrogen bonds are the most common form of polar bonds. • Van der Waals forces are very weak interactions that cause atoms to weakly attract each other in the absence of all other forces. Water • Water, H2O, is a polar compound. The 2 hydrogens are held at an angle to each other, and so the oxygen end of the molecule is partially negative and the hydrogen end is partially positive. • Water forms many hydrogen bonds with other water molecules and with other polar substances. This causes cohesion: water molecules sticking together. Also adhesion: water sticking to other (polar) things. Hydrophilic and Hydrophobic • Polar (and ionic) substances dissolve in water, because water forms hydrogen bonds with the polar molecules. Thus, polar substances are called hydrophilic or “water-loving”. • Non-polar substances don’t dissolve in water because they can’t form hydrogen bonds, so they are called hydrophobic or ‘water-fearing”. Oils and fats are examples of non-polar substances. – Non-polar substances dissolve in non-polar solvents. Hexane and toluene are examples. • “Like dissolves like”. • Lipids are the main hydrophobic molecules in the cell. Membranes are composed of lipids. The inside and outside of the cell are polar, aqueous environments, but they are separated by a hydrophobic barrier that prevents them from mixing freely. Hydrophobic molecules tend to clump together, to minimize their interactions with water molecules. This is a major driving force in the assembly of membranes and in protein folding. Acids and Bases • Water molecules tend to dissociate into H+ and OH- ions. The H+ ion is just a proton. – It’s an equilibrium: water molecules are constantly dissociating and associating. • In practice, free protons are not found in water. Instead, a proton combines with another water molecule, generating H3O+ (hydronium ion) and OH- (hydroxyl ion). – In pure water, the concentrations of hydronium and hydroxyl are each 10-7 M. This is a pH of 7. • Acids produce H+ ions when dissolved in water. For example, HCl (hydrochloric acid). • Bases absorb H+ ions when dissolved in water. A substance that releases OH- ions is a base because the OH- combine with the H+ in the water. • Water is neutral, neither acid nor base, because it always has equal numbers of H+ ions and OH- ions. Strong vs. Weak • Strong acids and bases completely dissociate as soon as they are dissolved in water. Examples: HCl and NaOH. • Weak acids and bases do not completely dissociate: some molecules retain the H+ or OH-. The percentage that are dissociated depends on the dissociation constant of the acid or base. • Most biological acids and bases are weak. • Biological acids contain the –COOH group (carboxylic acid) • Biological bases mostly contain the –NH2 group (amino) – The amine group removes a H+ (proton) from water, generating OH- Dissociation Constants • The dissociation constant is an equilibrium constant: it describes the relative amounts of the different molecules at equilibrium. • We write the relevant reaction, then put the concentrations of the products multiplied together in the numerator, and the concentration of the original compound in the denominator. – We ignore the concentration of water. – For acids, we use the acid dissociation constant, Ka. This is easily mistaken for the association constant. – For bases, we use the base dissociation constant, Kb, which is exactly the same thing as Ka except that [OH-] is used instead of [H+] is used. More Dissociation Constants • Dissociation constants are often written as pKa, the negative logarithm of the Ka. This is the same as pH, which is the negative log of the hydrogen ion concentration. • For a single H+ or OH-, the pKa is the pH at which half the molecules are dissociated. pH Scale • Acidity is measured on the pH scale, which indicates the amount of H+ ions present. The scale runs from 0 (very acidic) to 14 (very basic). • pH is defined as the negative logarithm of the H+ concentration. – Water, which is neutral, has a pH of 7 because its H+ concentration is 10-7 • Acids have lower pHs: the hydrochloric acid in your stomach has a pH of about 2. Eating food stimulates your stomach to secrete more acid. Antacids (like Tums or Rolaids) neutralize some of this acid. • Bases have a higher pH: the lye in oven cleaner has a pH of about 14. • Body fluids are slightly basic, pH 7.4 Buffers and Salts • Too much acid or base is harmful. The body needs to protect itself against large pH shifts. It uses buffers, pairs of weak acids and weak bases to absorb excess H+ and OH- and keep the body’s pH near neutral. • The main buffer in the body is carbonic acid, which dissociates into H+ and HCO3-. If H+ is added by an acid, it gets converted into the neutral H2CO3. Similarly, excess OH- combines with the H+, leaving the much less basic HCO3-. These opposing reactions keep the pH at the proper level. Buffers and Salts Cl- Na+ • Salts are ionic compounds that don’t release H+ or OH- when they dissolve. Thus, sodium chloride is a salt because it dissolves to form Na+ and Cl- , while hydrogen chloride (hydrochloric acid, HCl) is an acid because it dissolves to form H+ and Cl-. • Acids usually have an associated salt. An example of this is glutamic acid, an important component of proteins, and monosodium glutamate (MSG), which is used as a flavor enhancer in food. They are the same thing chemically, except that MSG has a sodium where glutamic acid has a hydrogen. Simple Organic Chemistry • Carbon forms 4 bonds, because it has only 4 electrons in its outer shell and needs 4 more to become stable. This means carbon can form many different compounds. – Nitrogen forms 3 bonds, oxygen forms 2 bonds, hydrogen and halogens (F, Cl, Br, I) form 1 bond. • Carbon bonds are quite stable at typical temperatures. • When drawing organic molecules, the hydrogens are often left out completely • Also, the lines connecting carbons are drawn, but not the actual “C”s. Other atoms are represented by their letter symbols. More Organic Chemistry • Since the 4 single bonds at a carbon atom are formed in a tetrahedron, chains of carbons are often drawn as a squiggly line. Hydrogens are left off: their existence is implied by the need for each C to have 4 bonds. • Many ring-shaped molecules have alternating single and double bonds, which are actually all identical, and thus can be drawn as a circle inside the carbon ring. The simplest example is benzene. Molecules of this type are called aromatic, and they have special properties. – Hydrocarbons that don’t have aromatic rings are called aliphatic. Functional Groups • Many organic compounds have functional groups attached to them. Functional groups are groups of atoms that have specific functions and properties. • The –OH group is called hydroxyl, and molecules with it are called alcohols, with names that end in –ol. • The C=O group is called carbonyl, and molecules containing it are aldehydes or ketones. – A carbonyl group that has another carbon attached to one side is called an acyl group. • Also important: -CH3 (methyl group) • Also –PO4 (phosphate group). At biological pHs, phosphates carry a -2 charge, or a -1 charge if they are connected to 2 other groups. Stereoisomers • If a compound contains a carbon atom is attached to 4 different groups, there will be a left handed version and a right handed version, mirror images of each other. – These molecules are called chiral molecules, and the carbon atom is called an asymmetric carbon atom or a chiral center. • These are stereoisomers, molecules with the same chemical and structural formula but which are not superimposable. – Stereoisomers don’t exist for atoms that join with 3 or fewer others, or with carbons that have 2 or more identical side groups. • It is possible for a molecule to have several chiral centers, which give rise to a whole series of stereoisomers. – Stereoisomers that have all chiral centers in the opposite configuration as each other are called enantiomers. – Stereoisomers that are not enantiomers are called diastereomers. Erythrose has D- and L- forms (enantiomers). Erythrose is a diastereomer with threose: they are not mirror images, but they have the same chemical and structural formulas. More Stereoisomers • In biology, the amino acids used in proteins are all left-handed (L-amino acids: L stands for “levo”). Also, the sugars used in the cell are right handed (D-sugars: D stands for “dextro”); this includes the ribose and deoxyribose used in nucleic acids. – If a beam of light is shined on a solution of a stereoisomer, the plane of polarization is rotated, the to left or to the right. For this reason, stereoisomers are considered to be optically active. – However, the D and L refer to a compound’s chemical relationship to the original stereoisomers of glyceraldehyde, not how the beam of is actually rotated. • Living systems usually use just one stereoisomer. On the other hand, when a chiral compound is synthesized in the laboratory from optically inactive components, it usually appears as a mixture of D and L forms. This is called a racemic mixture.