Biochemistry Basics

1. Biochemistry Basics Section 1.1

3. Subatomic Particles and the Atom • Protons (+ charge) and neutrons (neutral) • found in the nucleus • Electrons (- charge) • Surround the nucleus in a “cloud” or orbital • Orbital • the 3D space where an electron is found 90% of the time • Each orbital can fit only 2 electrons

4. Hydrogen atoms (2 H) In each hydrogen atom, the single electron is held in its orbital by its attraction to the proton in the nucleus. + + 2 3 1 When two hydrogen atoms approach each other, the electron of each atom is also attracted to the proton in the other nucleus. + + The two electrons become shared in a covalent bond, forming an H2 molecule. + + Hydrogen molecule (H2) Bonding – Covalent Bonds • Atoms bond through interaction of their valence (outer orbital) electrons • Covalent bond • electrons are shared between atoms and the valence orbitals overlap

5. Name (molecular formula) Electron- shell diagram Space- filling model Structural formula Water (H2O). Two hydrogen atoms and one oxygen atom are joined by covalent bonds to produce a molecule of water. H O H Methane (CH4). Four hydrogen atoms can satisfy the valence of one carbon atom, forming methane. H H H C H

6. Ionic Bonds • In some cases, atoms strip electrons away from their bonding partners • Ionic bond – electrons are transferred from one atom to the other, resulting in a negative ion (anion) and a positive ion (cation), which are electrostatically attracted to each other

7. Each resulting ion has a completed valence shell. An ionic bond can form between the oppositely charged ions. The lone valence electron of a sodium atom is transferred to join the 7 valence electrons of a chlorine atom. – + Cl Na Na Cl Cl– Chloride ion (an anion) Na+ Sodium on (a cation) Na Sodium atom (an uncharged atom) Cl Chlorine atom (an uncharged atom) Sodium chloride (NaCl)

8. Covalent bonds are stronger than ionic bonds • Covalent and Ionic bonds are intramolecular forces of attraction because they are within molecules

9. Polarity • Electronegativity • Is the attraction of an atom for electrons • The more electronegative an atom • The more strongly it pulls electrons toward itself • The smaller the atom • the more electronegative

10. to determine the type of bond between two atoms, calculate the difference between their electronegativity values =0 covalent strong electrons shared equally electrons 0 < x < 1.7 polar covalent partially shared >= 1.7 ionic weak electrons not (extreme polarity) shared • the greater their difference in electronegativity, the greater the polarity of that substance

11. Polar Covalent Bond – electrons are shared unequally between atoms of different electronegativity; electrons are closer to the atom with the higher value Because oxygen (O) is more electronegative than hydrogen (H), shared electrons are pulled more toward oxygen. This results in a partial negative charge on the oxygen and a partial positive charge on the hydrogens. d– O H H d+ d+ H2O

12. Intermolecular Forces • intermolecular forces of attraction exist between molecules • London forces • form when the electrons of one molecule are attracted to the positive nuclei of neighbouring molecules; holds large nonpolar molecules together; very weak

13. H Water (H2O) O A hydrogen bond results from the attraction between the partial positive charge on the hydrogen atom of water and the partial negative charge on the nitrogen atom of ammonia. H  +  – Ammonia (NH3) N H H d+ + H Figure 2.15 • hydrogen bonds • form when the slightly negative O or N that is bonded to a slightly positive H is attracted to the slightly positive H of a neighbouring molecule; strongest  +  –

14. dipole-dipole forces • form when the slightly negative end of a polar molecule is attracted to the slightly positive end of a neighbouring polar molecule; stronger • Occurs because electrons are in constant motion and may accumulate by chance on one part of the molecule. The result is “hot spots” of positive and negative charge.

15. – Hydrogenbonds + H – + H + –  – + Figure 3.2 Water • highly polar because of asymmetrical shape and polar covalent bond • The polarity of water molecules results in hydrogen boding

16. “Like Dissolves Like” • ionic compounds dissolve in water because the ions separate

17. However, molecules do not need to be ionic to dissolve in water • Smaller polar covalent molecules (eg: sugars, alcohols) can dissolve in water, but large nonpolar molecules (eg: oils) do not • small nonpolar molecules (eg: O2, CO2) are slightly soluble and need soluble protein molecules to carry them (eg: hemoglobin transports oxygen through the blood)

18. hydrophilic – “water-loving;” dissolves in water • e.g. polar or ionic molecules, carbohydrates, salts • hydrophobic – “water-fearing;” does not dissolve in water • e.g. non-polar molecules, lipids

19. Acids and Bases • acid – donates H+ to water; pH 0-7 • base –donates OH- to water (or H3O); pH 7-14 • neutralization reaction – the reaction of an acid and a base to produce water and a salt (ionic compound)

20. Strong and Weak Acids/Bases • strong acids and bases – ionize completely when dissolved in water • HCl(aq) (100% H3O+(aq)) • NaOH(aq) (100% OH-(aq)) • weak acids and bases – ionize only partially when dissolved in water • CH3COOH(aq) (1.3%  H3O+(aq)) • NH3(aq) (10%  OH-(aq))

21. Buffers • The internal pH of most living cells must remain close to pH 7 • Buffers • Are substances that minimize changes in the concentrations of hydrogen and hydroxide ions in a solution • Can donate H+ ions or remove H+ ions when required • E.g. carbonic acid creates bicarbonate ions (base) and hydrogen ions (acid) (reversible reaction)

22. To Do • Section 1.1 Questions • Pg. 23 #1, 2, 4, 6-8, 12, 14, 15