Biology 101_Chemistry of Life

The first lecture of the Biology 1 for health sciences students

1. University of Sharjah Faculty of health Sciences Department of Medical Laboratory Sciences Biology 1 Chapter2 ||| Lecture 1

2. Chapter 2 Chemistry of life Please note: Notes to follow are organized according to chap. 2 in the 10 th ed. of the textbook. These brief notes are meant to help you read and understand the text material and not to substitute for it. 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

3. The hierarchy of life At the most basic level, life is made of atoms  molecules  macromolecules  organelles  cells  may be unicellular (e.g. bacteria)  or multicellular (plants, animals, etc.) Multicellular life has higher levels of organization above cells For humans cells make up tissues  organs  organ systems  organisms (full human) Therefore, one can say that life is Matter

4. Matter: - It is what we and everything else in the universe is made of. - Technically = anything that occupies space and has mass . States of matter: Solid : has specific volume and shape Liquid : has a defined volume but conforms to the shape of the container it occupies Gas : No definite volume or shape

5. Matter is also energy, so what is energy? Energy: ability to do work, 2 types: Potential: stored energy like that in chemical bonds, water in tank above the roof, and so on Kinetic: energy of motion or energy displayed while the work is being done Forms : mechanical, chemical, heat/radiation, electrical, ...etc. It is possible to convert energy from one form to another. Example: we eat plants which photosynthesize, so: Sun (radiation) absorbed by plants  chemical energy ( glucose ) produced by plants  used to produce chemical (food) energy in animals  in animal cells it gets converted into chemical energy ( proteins and lipids, ATP and so on ) mechanical (movement of muscles), heat ( metabolism-related body heat) , electrical (nerve impulses), etc.

6. Atoms & elements An element is a unique substance that cannot be broken down into simpler substances by conventional approaches 112 elements are present in nature; 92 occur naturally and 20 are produced artificially. C, H, O, & N = 96% of body weight; The remaining 4% is made up of trace elements ( table 2-1) –Calcium ( Ca )  bone formation –Phosphorus (P)  DNA synthesis –Potassium (K)  cell signaling, nervous system –Sulfur (S) –Sodium (Na) –Chlorine (Na) –Magnesium (Mg) Trace elements are present in trace, or minute quantities, but are extremely important

7. What about the remaining 0.01%? –Iron (Fe) – hemoglobin, binds oxygen –Copper (Cu) – enzymes, electron transport –Fluorine (F) – prevents tooth decay, added to municipal water –Iodine (I) – thyroid enzymes (deficiency -> goiter); “iodized salt”

8. Each element is made up of small building blocks = atoms Atom from Greek for incapable of being divided Atoms are made up of small subatomic particles; main ones are protons, neutrons and electrons Subatomic particles differ in their mass, electrical charge and location within/around atoms  table 2-2

9. Main Subatomic particles

10. Models of atomic structure = orbital vs. planetary Planetary = protons + neutrons clustered in the core with electrons orbiting around in designated shells  something like a solar system Orbital = a dense core of protons + neutrons with electrons represented as a dense cloud of negative charge instead of fixed orbits

11. Identifying elements Atomic number: equivalent to number of protons inside the atom (= number of electrons) >> why? Atomic mass: Σ protons + neutrons [but not electron as weight of electrons (1/800 that of protons) is negligible] Atomic weight: same as atomic mass for atoms with non-variable atomic structure . For certain elements, the number of neutrons ( but not protons or electrons ) in the atom vary = isotopes . Hence the mass of the atom changes = atomic weight is used to denote a changing atomic mass of the element

12. Isotopes All isotopes of an element have the same atomic number but different atomic mass All isotopes exhibit the same chemical properties (fixed number of electrons) Atomic weight of an element = atomic mass of the most abundant isotope of the element Example: Hydrogen has an atomic number of 1 but it has 3 isotopes with atomic masses 1, 2 or 3. Atomic weight of H = 1.0079 which is ~ atomic mass of 1 H (the most abundant isotope of H). Heavier isotopes of an atom are less stable than lighter ones  decay or decompose fast so as to stabilize  radioisotopes

13. Radioisotopes are radioactive = emit (generate) electromagnetic particles (  or  = energy Particles can be detected, hence use of isotopes as biological tracers in research ( 125 I, 32 P, 35 S, 3 H) or in medical diagnostic imaging (PET scanning, iodine uptake activity of the thyroid gland, etc.). Some particles have high energy to destroy /change the structure of certain molecules like DNA  hence their use in radiotherapy  radioactive cobalt or radium in cancer therapy. **For more on the medical use of radioisotopes see pages 10-11 in textbook

14. Molecules and Compounds : When 2 or more atoms of the same element combine = molecules form  O 2 and H 2 are molecules When 2 or more atoms of different elements combine = compounds form  NaCl , H 2 O, CH 4 are compounds In practice, the two terms are interchangeable? When a compound forms it acquires a number of emergent properties that differ from those of the forming atoms (boiling point, freezing point, chemical behavior, etc ).

15. Chemical bonds and chemical reactions - Electrons are distributed around the dense core of atoms in shells ( orbitals ); 1 st shell takes only 2 electrons to fill up (stabilize) that is why atoms of helium He 2 are stable (nonreactive or inert), shells 2, 3, …. 7 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 If valence shell of an atom has 2 (helium) or 8 (Neon, Argon, Xenon, …etc) electrons = stable  atom non-reactive = inert substance

16. Reactive vs. nonreactive ( inert ) elements

17. Chemical Bonding

18. Ionic bonds: total transfer of an electron from one atom into another atom  atoms that loses the electron become positively charged ( cation ) and that which acquires the electron becomes negatively charged ( anion )

19. Covalent bonds: Form between similar or different atoms with valence shells lacking more than one electron or requiring more than one electron to complete their valence shells (except for hydrogen).   Sharing of a pair of electrons between 2 atoms Examples: C and C, C and H, O and O, O and H, N and H, etc. A covalent bond could be non-polar = both atoms have the same affinity for electrons or polar = one atom has high affinity for electrons than the other atom Covalent bonds that form between 2 identical atoms is always non-polar; example C-C C-H is also non-polar as the affinity of both C and H for electrons is the same

20. Hydrogen bonds: forms between atoms in molecules with polar covalent bonds Example: the covalent bond between oxygen and hydrogen is polar because oxygen loves electrons far more than hydrogen  the pair of electrons between the 2 atoms move more towards oxygen  more negative charge closer to oxygen and more positive charge closer to hydrogen  partially positive hydrogens in water attracts partially negative oxygens present on a different water molecule and vice versa

21. The hierarchy of life At the most basic level, life is made of atoms  molecules  macromolecules  organelles  cells  may be unicellular (e.g. bacteria)  or multicellular (plants, animals, etc.) Multicellular life has higher levels of organization above cells For humans cells make up tissues  organs  organ systems  organisms (full human)

22. Chemical reactions Absorb energy Release e nergy Absorb & release energy Anabolic Catabolic Anabolic / catabolic (condensation) (hydrolysis)

23. Factors that influence the rate of chemical reactions (table 2-4) Temperature Concentration of reactants Particle size Catalysts (e.g. enzymes)