BASIC CHEMISTRY REVIEW Intro to Biochemistry

1. BASIC CHEMISTRY REVIEWIntro to Biochemistry

2. Chemistry Fundamentals • All living things are made up of matter • Matter has mass, occupies space and has many forms • The atom is composed of a tiny nucleus containing protons (+), and neutrons (neutral) surrounded by (-) electrons in orbit

3. Chemistry Fundamentals • Mass number = # of protons and neutrons • Atomic Number = # of protons

4. Isotopes • Atoms that have different number of neutrons are called isotopes. Same atomic number, but different mass. • Carbon has 3 isotopes • 99% of Carbon found in nature is C-12 • It was discovered that the nucleus of some isotopes break apart (decay) over time • These are called radioisotopes

5. Radioisotopes • Radioisotopes are radioactive, that is, they decay into smaller atoms, subatomic particles and energy • All Radioactive isotopes have a half life-the time it takes for one half of the atoms in a sample to decay. • Half life for different radioisotopes vary while rate of decay of a particular isotope is constant.

6. Radioisotopes • Benefits-radiometric dating, radioactive tracers • Tracers can follow chemicals through chemical reactions and their path as they move through cells and the body • Radio labelled molecules can also be used

7. Carbon DatingMeasuring the ratio of C-12 to C-14 in a dead/fossilized organism allows one to calculate the time that has elapsed since the organisms death. C-12 remains constant while C-14 decays

8. Hazards • Radiation from decaying radioisotopes is harmful to living tissues and cells (mutations) • Dosimeters are used by physicians/scientists to monitor radiation levels

9. Chemical Bonding • Electrons are arranged in energy levels in spaces around the nuclei called orbitals. The further away from the nucleus, the more potential energy electrons have • An orbital can only accommodate 2 electrons and has either a spherical (s) shape or a dumbbell shape (p) • n=1 (1st energy level) 1s orbital • n=2 (2nd energy level) 2s orbital, 2p orbitals (x, y, z) • n=1 (max. 2 electrons) • n=2 (max. 8 electrons)

10. Orbital shapes

11. Valence electrons • Located on outermost “s” and “p” orbitals and determine an atom’s chemical behaviour • Noble gases have full valence orbitals and are thus stable (don’t gain, lose, or share electrons) He, Ne, Ar... • Other elements attempt to gain, lose, share electrons to become stable. It is these interactions that cause chemical bonds and jump start reactions.

12. Lewis dot diagrams

13. Periodic Table • Vertical columns=group (family) • Horizontal rows=periods

14. Ionic vs. Covalent bonding (video) • Compounds are stable combinations of atoms of different elements held together by chemical bonds (intramolecular forces of attraction) • When atoms lose electrons they become positively charged (cations), when they gain electrons they become negatively charged (anions) • Ionic bond=attraction b/w cations and anions (eg. NaCl) • Covalent bond=2 atoms share one or more pairs of valence electrons (can be single-H2O, double-O2(g) or triple N2(g) bonds. • Covalent bonds are usually stronger than ionic bonds

15. Electronegativity and Polarity • Electronegativity is a measure of an atom’s ability to attract a shared electron pair (in a covalent bond) • The larger the electronegativity #, the stronger the atom attracts the electrons. • Atoms that attract stronger are assigned a negative charge (-δ) • Atoms that attract weaker are assigned a positive charge (+δ) • “δ” denotes the term “partial” as electrons are shared • The difference in electron attraction forms a polar covalent bond

16. Table of electronegativities (p.14)

17. Electronegativity difference • If ΔEn is zero that means the electron pair are sharing equally (non-polar covalent bond) • If ΔEn is greater than zero but less than 1.7, the bond is polar covalent • If ΔEn is greater than 1.7 the bond is considered ionic. • Atoms from groups 1, 2 and 16, 17 generally form ionic compounds

18. Molecular Shape • A molecule’s function is determined by its bonds b/w atoms, shape and polarity. • When covalent bonds are formed, hybridization occurs (change in the orientation of the valence electrons) • The Valence Shell Electron Pair Repulsion (VSEPR) Theory can predict molecular shape

19. VSEPR • B/c electrons are negative, valence electron pairs repel each other • 4 valence pairs (CH4)-tetrahedral • 3 valence pairs (NH3)-pyramidal • 2 valence pairs (H2O)-angular • 1 valence pair (HCl)-linear

20. Molecular Polarity • Covalent bonds can be polar or non-polar…but the polarity of a molecule as a whole depends on bond polarity and molecular shape. • Symmetrical structure + polar/non-polar bonds = non-polar molecules • Asymmetrical structure + non-polar bonds = non-polar molecules • Asymmetrical structure + polar bond(s) = polar molecules

21. Water • Life would not exist without water! • H2O has polar covalent bonds and asymmetrical shape, making it a highly polar molecule. • Bonds b/w molecules are called intermolecular and help determine physical states (solid, liquid, gas) • 3 types; London forces, dipole-dipole forces and Hydrogen bonds (collectively called van der Waals forces)

22. Intermolecular bonds • London dispersion forces are the weakest. Exist b/w all atoms and molecules (e.g. noble gases , non-polar molecules) • Formed by the temporary unequal distribution of electrons as they move around the nuclei • This is why He and CH4 are gases at room temperature

23. Intermolecular bonds • Dipole-dipole forces hold polar molecules together. Stronger than London forces • H-bonds are the strongest force of attraction. • Water molecules combine by H-bonds • Water is considered the universal solvent due to its polarity (+ve and –ve charges attract other polar molecules and ions) • Disassociation = ionic bonds broken • Salt in water

24. Water • Many substances dissolve in water (soluble-sugar), while others do not (insoluble-chalk) • Miscible-liquids that dissolve into one another (ethanol is miscible in water) • Immiscible-liquids that form separate layers (oil) • The subscript “aq”, aqueous means that it is dissolved in water

25. Water • Some small non-polar molecules can’t form H-bonds with water (O2, CO2) are only slightly soluble. • That is why hemoglobin (protein carrier) is needed to transport oxygen • Non-polar molecules = hydrophobic • Polar molecules = hydrophilic

26. Acids, Base and Buffers • Acids-sour taste, conduct electricity, turn blue litmus red, pH less than 7. Increase the concentration of H3O+(aq) • Bases-bitter taste, slippery feel, conduct electricity, turn red litmus blue, pH greater than 7. Increase the concentration of OH-(aq) • Pure water is neutral, pH of 7 (H3O+(aq) = OH-(aq)) • When acids are mixed with bases it is considered a neutralization reaction (water and salt are produced)

27. Strong Acids and Bases • Dependent on the degree they ionize when dissolved in water • Strong acids (HCl) and bases (NaOH) ionize completely

28. Weak Acids and Bases • Weak acids (CH3COOH(aq) - acetic acid) and bases (NH3 (aq) – ammonia) ionize partially in water. • They are reversible reactions in a state of equilibrium

29. Acid-Base Buffers • Living systems are sensitive to pH levels • Living cells use buffers to resist significant changes in pH • Carbonic acid-bicarbonate buffer, most common. • When carbon dioxide and water react, they form carbonic acid, which then ionizes to form bicarbonate and H+ ions. • When H+(aq)ions enter the blood (acidic food),HCO3 (aq) reacts with it to produce H2CO3(aq). • Together they help maintain the pH of blood around 7.4