Lecture 1 Chemical Biology PDF

LECTURE 1 CHEMICAL BIOLOGY Key concepts Matter consists of chemical elements in pure form and in combinations called compounds. An element’s properties depend on the structure of its atoms. The formation and function of molecules depend on chemical bonding between atoms. Polar covalent...

LECTURE 1 CHEMICAL BIOLOGY Key concepts Matter consists of chemical elements in pure form and in combinations called compounds. An element’s properties depend on the structure of its atoms. The formation and function of molecules depend on chemical bonding between atoms. Polar covalent bonds in water molecules result in hydrogen bonding. Acidic and basic conditions affect living organisms. Carbon atoms can form diverse molecules by bonding to up to four other atoms. A few chemical groups are key to molecular function. 1.1. Atoms Elements Matter is made up of elements. Elements are substances that cannot be broken down to other substances by chemical reactions. About 20–25% of the 92 elements are essential to life. Four make up 96% of living matter (carbon, hydrogen, oxygen, and nitrogen). Most of remaining 4% of living matter consists of calcium, phosphorus, potassium, and sulphur. Trace elements are required by an organism in minute quantities (example: iron, iodine, etc.) Other living organisms will have differing amounts. Atoms Each element consists of unique atoms. An atom is the smallest unit of matter that retains the properties of an element. Atoms are composed of subatomic particles. Relevant subatomic particles include neutrons (no electrical charge), protons (positive charge), electrons (negative charge). Neutrons and protons form the atomic nucleus. Electrons form a cloud around the nucleus. Atomic number and mass number Atoms of the various elements differ in number of subatomic particles. An element’s atomic number is the number of protons in its nucleus. An element’s mass number is the sum of protons plus neutrons in the nucleus. Atomic mass, the atom’s total mass, can be approximated by mass number. Isotopes All atoms of an element have same number of protons, but may differ in the number of neutrons. Isotopes are two atoms of an element that differ in number of neutrons. Radioactive or unstable isotopes decay spontaneously, giving off particles and energy. Ions An ion is an atom or molecule in which the total number of electrons is not equal to the total number of protons, giving the atom a net positive or negative charge. The atom/molecule missing an electron (carrying the (+) charge) is called a cation. The atom/molecule that has an extra electron (carrying the (-) charge) is called an anion. Electrons’ energy Electrons of an atom differ in their amounts of potential energy. An electron’s state of potential energy is called its energy level, or electron shell. Atom’s chemical behaviour is determined by distribution of electrons in its electron shells. Periodic table The periodic table of the elements shows the electron distribution: Rows = periods = number of shells | Columns => addition of 1 electron (and proton) Valence electrons are those in outermost shell, or valence shell. Chemical behaviour of an atom is mostly determined by valence electrons. Elements with a full valence shell are chemically inert. Note that the first electron shell can only accommodate up to two electrons. 1.2. Chemical bonds Ø Covalent Bonds Ø Ionic Bonds Ø Hydrogen Bonds Covalent bonds Atoms with incomplete valence shells share or transfer valence electrons with other atoms. These interactions result in atoms being held by attractions called chemical bonds. Covalent bond is the sharing of a pair of valence electrons by two atoms. A molecule consists of two or more atoms held together by covalent bonds. Sharing of one pair of valence electrons results in a single bond. Double and triple bonds are also possible. Covalent bonds can form between atoms of the same element or of different elements. A compound is a combination of two or more different elements. All compounds are molecules but not all molecules are compounds. Valence Bonding capacity is called the atom’s valence. Example: Valence numbers of H, O, N, C is 1, 2, 3, 4 respectively. The situation is more complex for elements in the third row of the periodic table (P for example can have a valence of 3, but in some molecules it can form 3 single bonds and one double bond, having therefore a valence of 5. Electronegativity 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. Therefore, there are two different types of covalent bonds: Nonpolar covalent bonds where atoms share electrons equally (example: H-H) Polar covalent bonds where one atom is more electronegative and the atoms do not share electrons equally (example: water). Unequal sharing of electrons causes a partial positive or negative charge for each atom of a molecule. Ionic bonds Covalent bonds are among the strongest that link atoms to form the molecules found in a cell. Weak chemical bonds (e.g. ionic and hydrogen bonds) are also indispensable, as they regulate interactions of macromolecules in biological systems. 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 (they become ions). An ionic bond is an attraction between an anion and a cation. Ionic compounds Compounds formed by ionic bonds are called ionic compounds, or salts. Unlike a covalent compound, which consist of molecules having a definite size and number of atoms, the ionic compounds does not consist of molecules, but organize in crystal lattices. The formula (NaCl) only represents the ratio of ions in the lattice. For example, MgCl2 means a lattice with 2 chloride ions per magnesium ion. Hydrogen bonds A hydrogen bond forms when a hydrogen atom covalently bound to an electronegative atom is also attracted to another electronegative atom. In living cells, electronegative partners are usually oxygen or nitrogen. In this example, water = H-bond donor, and ammonia = H-bond acceptor. Van der Waals interactions Asymmetrically distributed electrons in molecules or atoms 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. Collectively, such interactions can be strong, as between molecules of a gecko’s toe hairs and a wall surface. Chemical reactions 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. Note: Matter is conserved in chemical reactions: Atoms do not break! Only their connections do. Most chemical reactions are reversible: products of the forward reaction become reactants for the reverse reaction. Chemical equilibrium is reached when the forward and reverse reactions occur at the same rate. The concentrations of reactants and products are not necessarily the same at equilibrium. 1.3. Water and pH Water is a polar molecule Water is the biological medium on Earth: abundance of water is an important reason Earth is habitable. All living organisms require water more than any other substance: most cells are surrounded by water, and are 70–95% water. Water is a polar molecule due to highly electronegative oxygen atom: the overall charge is unevenly distributed. Polarity allows water molecules to form hydrogen bonds with each other. Aqueous solutions An aqueous solution is one in which water is the solvent (the dissolving agent). Water is a versatile solvent due to its polarity: it forms hydrogen bonds readily. When an ionic compound is dissolved in water, each ion is surrounded by water molecules (hydration shell). Dissolution of nonionic molecules Water can also dissolve nonionic polar molecules or compounds. Large polar molecules such as proteins can dissolve in water if they have polar or ionic regions. Hydrophilic substance has an affinity for water. Hydrophobic substance does not have an affinity for water. Oil molecules are hydrophobic because they have relatively nonpolar bonds. They get together to form another type of chemical interaction: hydrophobic interactions. Water dissociation A hydrogen atom participating in a hydrogen bond between two water molecules can shift from one to the other. The hydrogen atom leaves its electron behind and is transferred as a proton, or hydrogen ion (H+). The molecule that lost the proton is now a hydroxide ion (OH−). The molecule with the extra proton is now a hydronium ion (H30+), though it is often represented as H+. pH Water is in a state of dynamic equilibrium: water molecules dissociate at same rate they are being reformed. In any aqueous solution at 25°C product of H+ and OH− is constant and written as [H+][OH−] = 10−14 The pH of a solution is defined by negative logarithm of H+ concentration, written as: pH = −log [H+] Concentrations of H+ and OH− are equal in pure water. Therefore, for a neutral aqueous solution, [H+] is 10−7, so pH = −(−7) = 7. pH scale Changes in concentrations of H+ and OH− can drastically affect the chemistry of a cell. Adding acids and bases, modifies concentrations of H+ and OH−. Acid: any substance that increases the H+ concentration of a solution HCl => H+ + Cl− Base: any substance that reduces the H+ concentration of a solution NH3 + H+ = NH4+ (ammonium ion) NaOH => Na+ + OH− The pH scale is used to describe whether a solution is acidic or basic. 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. Buffers The internal pH of most living cells must remain close to pH 7. Buffers minimize changes in concentrations of H+ and OH− in a solution. Buffers contain a weak acid and corresponding base, which combine reversibly with H+ ions. Example: Carbonic acid = Bicarbonate + Proton. 1.4. Organic molecules Life is carbon-based Living organisms consist mostly of carbon-based compounds. Carbon is able to form large, complex, and diverse molecules: proteins, DNA, carbohydrates in living matter are all composed of carbon compounds. Organic compounds range from simple molecules to colossal ones. Most contain hydrogen in addition to carbon Versatility of carbon The electron configuration of carbon gives it covalent compatibility with many different elements. The valences of carbon, hydrogen, oxygen, and nitrogen are the “building code” of living molecules. How many electron pairs do hydrogen, nitrogen, oxygen and carbon need share in order to complete their valence shell? 1 pair for hydrogen, 2 pairs for oxygen, 3 pairs for nitrogen, 4 pairs for carbon. The versatility of carbon makes possible the great diversity of organic molecules. Carbon skeleton Carbon chains form the skeletons of most organic molecules. Carbon chains vary in length and shape. Distinctive properties of organic molecules depend on: the carbon skeleton, and the molecular components attached to it. Hydrocarbons are organic molecules consisting of only carbon and hydrogen. Importance of functional groups Characteristic groups can replace hydrogens attached to skeletons of organic molecules. These groups are called functional groups are components of organic molecules involved in chemical reactions The number and arrangement of functional groups give molecules unique properties. Estradiol and testosterone are both steroids with a common carbon skeleton, in the form of four fused rings. These sex hormones differ only in the chemical groups attached to the rings of the carbon skeleton. Hydroxyl and carbonyl groups Carboxyl and amino groups Phosphate and methyl groups