Chemistry of Life Part 1 PDF

The Chemistry of Life Part 1 Learning objectives: Fundamentals of the chemistry of life: atoms, molecules, compounds, chemical reactions and chemical bonds. Inorganic compounds: water, salts, acids, bases, the pH scale … Organic compounds: polymerization reactions, the chemistr...

The Chemistry of Life Part 1 Learning objectives: Fundamentals of the chemistry of life: atoms, molecules, compounds, chemical reactions and chemical bonds. Inorganic compounds: water, salts, acids, bases, the pH scale … Organic compounds: polymerization reactions, the chemistry of macromolecules (sugars, lipids, nucleic acids, proteins). General intro to the functions of macromolecules in living matter The hierarchy of biological organization Hierarchical organization of life: Emergent properties Life’s Emergent properties Composition of Matter Atoms ▪ Building blocks of elements ▪ Atoms of elements differ from one another Subatomic Particles Nucleus ▪ Protons (p+) are positively charged ▪ Neutrons (n0) are uncharged or neutral Orbiting the nucleus ▪ Electrons (e–) are negatively charged Atoms are electrically neutral: ✓ Number of protons equals number of electrons in an atom ✓ Ions are atoms that have lost or gained electrons © 2015 Pearson Education, Limited. Figure 2.2 Atomic structure of the three smallest atoms. (a) Hydrogen (H) (b) Helium (He) (c) Lithium (Li) (1p+; 0n0; 1e–) (2p+; 2n0; 2e–) (3p+; 4n0; 3e–) KEY: Proton Neutron Electron  Molecules and Compounds Molecule—two or more atoms of the same elements combined chemically Example: H (atom) + H (atom) → H2 (molecule) (reactants) (product) Compound—two or more atoms of different elements combined chemically to form a molecule of a compound Example: 4H + C → CH4 (methane) (reactants) (product) © 2015 Pearson Education, Limited.  Chemical bonds and chemical reactions - Electrons are distributed around the dense core of atoms in shells (orbitals): - 1st shell takes only 2 electrons to fill up (stabilize). That is why atoms of helium He2 are stable (nonreactive or inert). - 2nd, 3rd, …7th shell: each takes 8 electrons to be complete. - Number of electrons in the last shell (valence shell) is key in determining the chemical bonding behavior of atoms  Chemical bonds and chemical reactions ▪ If valence shell of an atom has 2 (helium) or 8 (Neon, Argon, Xenon, …etc) electrons, the atom is stable (non-reactive or chemically inert) ▪ When the valence shell is less than 8, the atom tends to gain, lose, or share electrons with other atoms to complete their outermost orbital. ▪ Chemical bonds form. This will lead to: – Atoms reach a stable state – Bond formation produces a stable valence shell Figure 2.5a Chemically inert and reactive elements. (a) Chemically inert elements Outermost energy level (valence shell) complete 8e 2e 2e He Ne Helium (He) Neon (Ne) (2p+; 2n0; 2e–) (10p+; 10n0; 10e–) Figure 2.5b Chemically inert and reactive elements. (b) Chemically reactive elements Outermost energy level (valence shell) incomplete 4e 1e 2e H C Hydrogen (H) Carbon (C) (1p+; 0n0; 1e–) (6p+; 6n0; 6e–) 1e 6e 8e 2e 2e O Na Oxygen (O) (8p+; 8n0; 8e–) Sodium (Na) (11p+; 12n0; 11e–)  Chemical Bonds Ionic bonds – Form when electrons are completely transferred from one atom to another – Allow atoms to achieve stability through the transfer of electrons Ions – Result from the loss or gain of electrons Anions have negative charge due to gain of electron(s) Cations have positive charge due to loss of electron(s) – Tend to stay close together because opposite charges attract © 2015 Pearson Education, Limited. Figure 2.6 Formation of an ionic bond. Chemical Bonds Ionic bonds + – Na Cl Na Cl Sodium atom (Na) Chlorine atom (Cl) Sodium ion (Na+) Chlorine ion (Cl–) (11p+; 12n0; 11e–) (17p+; 18n0; 17e–) Sodium chloride (NaCl)  Chemical Bonds Covalent bonds – Atoms become stable through shared electrons – Electrons are shared in pairs – Single covalent bonds share one pair of electrons – Double covalent bonds share two pairs of electrons Examples of atoms that can make covalent bond: C and C, C and H, O and O, O and H, N and H, etc. © 2015 Pearson Education, Limited. Chemical Bonds Covalent bonds Figure 2.7c Formation of covalent bonds. Chemical Bonds Covalent bonds Reacting atoms Resulting molecules H H H C H C H or H H H Hydrogen atoms Carbon atom Molecule of methane gas (CH4) (c) Formation of four single covalent bonds  Chemical Bonds Covalent bonds Covalent bonds are either nonpolar or polar – Nonpolar Electrons are shared equally between the atoms of the molecule (e.g.: C-C and C-H) Electrically neutral as a molecule Example: carbon dioxide (a) Carbon dioxide (CO2) © 2015 Pearson Education, Limited.  Chemical Bonds Covalent bonds Covalent bonds are either nonpolar or polar – Polar Electrons are not shared equally between the atoms of the molecule (one of the atoms have higher affinity for the electrons than the other Molecule has a positive and negative side, or pole Example: water δ– δ+ δ+ (b) Water (H2O) © 2015 Pearson Education, Limited.  Chemical Bonds Hydrogen bonds – Forms between atoms in molecules with polar covalent bonds – Weak chemical bonds – Hydrogen is attracted to the negative portion of a polar molecule – Provides attraction between molecules – Responsible for the surface tension of water – Important for forming intramolecular bonds, as in protein structure © 2015 Pearson Education, Limited. Figure 2.9 Hydrogen bonding between polar water molecules. Hydrogen Bonds δ+ H H O δ– Hydrogen bonds δ+ δ+ δ– δ– δ– H H O O δ+ δ+ H H H δ+ O H δ– (a) (b)  Patterns of Chemical Reactions Synthesis reaction (A + B → AB) – Atoms or molecules combine – Energy is absorbed for bond formation – Underlies all anabolic activities in the body Decomposition reaction (AB → A + B) – Molecule is broken down – Chemical energy is released – Underlies all catabolic activities in the body Exchange reaction (AB + C → AC + B or AB + CD → AD + CB) – Involves both synthesis and decomposition reactions as bonds are both made and broken In general, most chemical reactions are reversible A+B AB © 2015 Pearson Education, Limited.  Chemical reactions Absorb energy Release energy Absorb & release energy Anabolic Catabolic Anabolic / catabolic (condensation) Table 2.4 Factors Increasing the Rate of Chemical Reactions. Factors influencing rate of chemical reactions  Biochemistry: Chemical Composition of Living Matter Chemicals found in the body are either Inorganic or Organic compounds - Inorganic compounds: mostly lack carbon and/or small; e.g. CO2, H2O, NH3, NaCl, Calcium, etc. - organic compounds: contain Carbon, tend to be larger in size than inorganic compounds; e.g. glucose C6H12O6, lipids, proteins, etc. Inorganic compounds in biological systems include: water, salts, acids, and bases Organic compounds in living matter include: proteins, lipids, sugars and nucleic acids among others  Inorganic Compounds Water : All properties are related to its ability to form hydrogen bonds High Heat capacity: amount of heat required to raise the temperature of one mole or one gram of a substance by one degree Celsius without change of phase. This means that water can absorb or release large amount of heat before changing temperature suddenly (sun exposure, winter, etc.) Polarity/solvent properties: because of its polarity, water is the most common solvent for infinite number of solutes; it is a Universal solvent. All chemical reactions occur in water Carrier of nutrients, gases and wastes  Inorganic Compounds Water properties (cont.) Capillary action: ability of water to flow in narrow spaces without the assistance of, and in opposition to external forces like gravity High Surface tension: increased cohesion of water molecules to each others. Think of the ability of some insects (e.g. water striders) to run on the water surface Chemical Reactivity: Water is an important reactant in many types of chemical reactions. Hydrolysis is the breakdown of large molecules due to addition of water into the reaction Cushioning /protection: fluids (water solutions) around heart, brain (CSF), developing fetus, etc.  Inorganic Compounds Salts: Ionic compounds that contain cations (other than H+) and anions (other than OH-) Salts of many metals are present in our bodies. e.g.: calcium salts (found in bone & teeth) and phosphorous (found in bone, teeth, nucleic acid) Easily dissolve (dissociate) in solvents like water Salts dissolved in body fluids to form ions (electrolytes = conductors of electricity). e.g.: Sodium Chloride is an example of electrolyte in blood (Na+ and Cl-) Electrolyte (ionic) balance is of utmost importance for the normal functioning of the body → loss of electrolyte balance is very harmful to. the body (e.g. kidney function), cell viability, etc  Inorganic Compounds Acids: Substances that release H+ (protons) Taste sour, very corrosive and damaging. Acids that ionize completely and release their protons are called strong acids (HCl) HCl → H+ + Cl- (complete dissociation) (proton) (anion) Acids that dissolve partially are called weak acids (e.g. carbonic acid). H2CO3  H+ + HCO3- + H2CO3 (partial dissociation) (proton) (anion) Weak acids are good biological buffers.  Inorganic Compounds Bases: Substances that accept H+ (protons) ➔ release OH- to accept H+ and form water Taste bitter and have a slippery feeling to them, NaOH → Na+ + OH- cation hydroxyl ion Bases that dissociate completely and release all their OH- are considered strong bases (NaOH) HCO3- can function as a weak base as it accepts H+ -------------------------------------------------------------------------- Neutralization Reaction: Type of exchange reaction in which acids and bases react to form water and a salt – Acid + base → H2O + salt – NaOH + HCL → H2O + NaCl  Inorganic Compounds The pH concept - The concentration of protons [H+] in solutions (in blood, body fluids, etc) reflects the degree of acidity or alkalinity of the solution. - It is not easy to directly measure proton [H+] concentration in a solution. - But as protons are ions (conductors of electricity), we can indirectly measure their relative concentration by measuring their degree of conductivity in the solution. - The output (reading) is called the pH = which is a measure of proton activity ([H+]) in a solution and is usually expressed as – Log10 [H+] Example 1 : if [H+] = 0.00001 = 10-5, then pH = -Log [10-5] ➔ -(-5) = 5 Example 2: If pH = 3, then–log H+ = 3 ➔ [H+] = 10-3 = 0.003 A difference of 1 unit = 10x up- or down-change in [H+]; A change in pH from 4 to 5 = 10-fold decrease in [H+] A change in pH from 6 to 5 = 10-fold increase in [H+]  Inorganic Compounds The pH concept (cont.) pOH is also a measure of proton activity in a solution. Example: if [OH-] = 10-7, then pOH = -Log [OH-] ➔ - (-7) = 7 Hence, pH = 14 – 7 = 7 (neutral solution) Range of pH scale = Zero (very acidic) to 14 (very basic) pH 7 is basic or alkaline The pH scale and pH values of representative substances  Inorganic Compounds Buffers In biological systems, range of [H+] (or proton activity) is narrow. Increased or decreased [H+] can disrupt or damage biological molecules, halt metabolism or kill cells. Example: [H+] in blood is around = 10-7 M; any sudden change in blood [H+] could be fatal → needs to be regulated Buffering systems exist to prevent sudden or dramatic changes in [H+] Weak acids and weak bases are good buffers