Basic Chemistry

1. Basic Chemistry

2. Matter and Energy • Matter—anything that occupies space and has mass (weight) • Energy—the ability to do work • Chemical – in the bonds of chemicals • Electrical – movement of charged particles • Mechanical – moving matter • Radiant – travels in waves, ie. xrays, light • Recall that chemical energy in the body is in the form of ATP, which is transformed into electrical and mechanical energy

3. Composition of Matter • Elements—fundamental units of matter ; made of one type of atom only • 96% of the body is made from four elements • Carbon (C) • Oxygen (O) • Hydrogen (H) • Nitrogen (N) • Atoms—building blocks of elements; smallest unit of an element that still has all the physical & chemical properties of that element

4. Atomic Structure • Nucleus • Protons (p+) • Neutrons (n0) • Outside of nucleus • Electrons (e-) Figure 2.1

5. Atomic Structure of Smallest Atoms Recall that, in a neutralatom, # protons = # electrons Figure 2.2

6. Identifying Elements • Atomic Number—equal to the number of protons that the atom contains • Atomic Mass Number— sum of the protons and neutrons

7. Isotopes and Atomic Weight • Isotopes • Atoms of an element that have the same number of protons and electrons • Have the same chemical properties • Vary in number of neutrons Figure 2.3

8. Isotopes and Atomic Weight • Atomic weight (aka “Average Atomic Mass”) • Close to mass number of most abundant isotope • Atomic weight reflects the natural isotope variation of an element

9. Radioactivity • Radioisotopes • Tends to be the heavier isotope(s) of an element • Usually unstable due to competing nuclear forces • Spontaneously decompose over time to more stable isotope, a process known asradioactive decay • Particles and/or energy that are released during radioactive decay may damage to living cells • Damage to cells is generally a function of amount of exposure over time • Minute amounts of radioisotopes can be used for diagnostic purposes and treatment of cancer

10. Molecules and Compounds • Molecule—two or more like atoms combined chemically • Compound—two or more different atoms combined chemically Figure 2.4

11. Chemical Bonds & Reactions • Atoms are united by chemical bonds when forming molecules and compounds • Atoms dissociate from other atoms when chemical bonds are broken

12. Electrons and Bonding • Electrons occupy regions in the electron cloud known as “energy levels” or energy shells • Electrons closest to the nucleus are held tightly, electrons further away are more easily released • Each level has distinct properties… • The number of electrons has an upper limit Level 1 = 2 Level 2 = 8, Level 3 = 18 Level 4 = 32… • Levels closest to the nucleus fill first, according to the electron configuration pattern shown in the Periodic Table (1s2 2s2 2p6 …)

13. Electrons and Bonding • Bonding involves interactions between electrons in the valence level (aka outer level) to reach 8 electrons (octet rule) and chemical stability. • Full valence shells do not form bonds, so elements are chemically inert

14. Chemical Reactions • Chemically active atoms will gain, lose, or share electrons to complete their outermost orbitals and reach a stable state

15. Chemical Bonds • Ionic Bonds • Form when electrons are completely transferred from one atom to another • Ions • Atoms that have gained or lost electrons, aka “chargedparticles” • Anions are negative, because electrons are gained • Cations are positive, because electrons are lost • Positive ions attract negative ions to form chemicalsalts PLAY Ionic Bonds

16. Ionic Bond Formation Cl Na Sodium atom (Na)(11p+; 12n0; 11e–) Chlorine atom (Cl)(17p+; 18n0; 17e–) Figure 2.6, step 1

17. Ionic Bonds Cl Na Sodium atom (Na)(11p+; 12n0; 11e–) Chlorine atom (Cl)(17p+; 18n0; 17e–) Figure 2.6, step 2

18. Ionic Bond Formation + – Cl Na Cl Na Sodium atom (Na)(11p+; 12n0; 11e–) Chlorine atom (Cl)(17p+; 18n0; 17e–) Sodium ion (Na+) Chloride ion (Cl–) Sodium chloride (NaCl) Figure 2.6

19. Chemical Bonds • Covalent bonds • Atoms become stable through shared electrons • Single covalent bonds share one pair of electrons: PLAY Covalent Bonds

20. Examples of Covalent Bonds • Double covalent bonds share two pairs of electrons Figure 2.7b

21. Examples of Covalent Bonds Figure 2.7c

22. Polarity • Occurs in covalently bonded molecules • Some are non-polar • Electrically neutral as a molecule • Some are polar • Have a positive and negative area • Polar molecules attract other polar molecules, including water, and are involved in cellular reactions. Figure 2.8

23. Chemical Bonds • Hydrogen Bonds • Weak chemical bonds that form between the hydrogen of one molecule and the nitrogen, oxygen or fluorine of another molecule • H bonds within protein molecules create the 3Dstructure that helps determine function.

24. Hydrogen Bonds • Also provides attraction between water molecules and links DNA molecules together across the double helix. Figure 2.9

25. Patterns of Chemical Reactions • Synthesis reaction (A + B  AB) • Atoms or molecules combine into a larger molecule • Energy is absorbed for bond formation • Used in reactions involved in cellular growth and tissue repair (protein construction) • Decomposition reaction (AB  A + B) • Molecules are broken down • Chemical energy is released • Used in digestion of food to release glucose molecules PLAY Disaccharides

26. Synthesis and Decomposition Reactions Figure 2.10a

27. Synthesis and Decomposition Reactions Figure 2.10b

28. Patterns of Chemical Reactions • Single & Double Replacement (Exchange) Reactions • (AB + C  AC + B) & (AB + CD  AD + CB) • Involves both synthesis and decomposition reactions • Switch is made between molecule parts and different molecules are made • Involved in the ATP  ADP cycle for cellular energy production

29. Patterns of Chemical Reactions Figure 2.10c

30. Factors Affecting Rate of Chemical Reactions • Temperature • Concentration of Particles • Particle Size • Use of Catalysts • Changes in any of these can effect the rate at which reactions occur in the body