1. CHAPTER 2: THE CHEMICAL CONTEXT OF LIFE

2. MULTIDISCIPLINARY SCIENCE Biological systems follow the same laws that are dictated by physics and chemistry. Emergent properties – biological functions are made possible by the interaction between atoms and molecules that make up an organism. Discuss: What is the definition of life? What is required for something to be considered alive? At what point do we cross the line between nonliving and living at the molecular level of cell biology? Do you consider a virus to be “alive?” Bombardier Beetle (chemistry) - https://www.youtube.com/watch?v=Pib9qT-pccIhttps://www.youtube.com/watch?v=Pib9qT-pccI Pistol Shrimp (physics) - https://www.youtube.com/watch?v=QXK2G2AzMTUhttps://www.youtube.com/watch?v=QXK2G2AzMTU

3. THE BASICS - IMPORTANT TERMS Matter – Anything that has mass and occupies volume. Element – A substance that cannot be broken down into other substances through chemical reactions. Compound – A substance that consists of two or more different elements in a fixed ratio. Atom – The smallest unit of an element that maintains the chemical properties of the element. Molecule – The smallest unit of a compound that maintains the chemical properties of the compound,

4. LE 2-2 SodiumChlorineSodium chloride

5. THE BASICS – PERIODIC TABLE

6. ELEMENTS ESSENTIAL TO LIFE 92 natural known elements 25 elements are essential to life 4 elements make up 96% of living matter Carbon (C) Oxygen (O) Hydrogen (H) Nitrogen (N) Remaining 4% Phosphorous (P) Sulfur (S) Calcium (Ca) Potassium (K) Less that 0.01% of living matter is made up of trace elements

7. THE STRUCTURE OF ATOMS Subatomic particles Neutrons - ~1 amu (1.7 x10ˆ-24 grams); no charge Protons - ~1 amu; +1e (1.602 x 10ˆ-19 coulombs) Electrons - 5.4858 x 10ˆ-4 amu (or 1/1840 amu); -1e (-1.602 x 10ˆ-19 coulombs) What makes up the atomic nucleus? 1 amu=1 dalton Atomic Weight - # of Protons + # of Neutrons Mass of electrons is negligible AKA Atomic Mass; Mass Number Atomic Number - # of Protons Atomic Number Atomic Weight Atomic NumberAtomic Weight

8. LE 2-4 Nucleus Electrons Cloud of negative charge (2 electrons)

9. Atoms of the same element that have different atomic weights Always have the same number of protons Differ in the number of neutrons Atomic Weight = Average Atomic Weight of all isotopes Some isotopes are stable Most isotopes are unstable (radioactive) ISOTOPES

10. ISOTOPES – APPLICATIONS IN LIFE SCIENCES Dating techniques for biological materials and fossils Tracers used to follow metabolic processes (Rate of DNA synthesis) Medical Diagnostics PET Scans Risk associated with using radioactive isotopes Severity depends on type and amount of radiation

11. ISOTOPES - PRACTICE How many protons, neutrons, and electrons are present in the following isotopes? Technetium-99 Flourine-18 Iodine-131 Cobalt-60 What is the atomic mass of an element, given the following isotopes? Oxygen-16, oxygen-17, oxygen-18, oxygen-12, oxygen-24, and oxygen-15

12. ELECTRONS – ENERGY LEVELS Energy is the capacity to cause a change; the capacity to do work (exert a force on an object). The energy level of an electron is the potential energy of that electron. Average distance of an electron from the nucleus Represented by electron shells Shells further from the nucleus represent higher levels of energy Energy input (light, etc.) required to move an electron to a higher energy level (further shell) Energy released (usually heat) when electron moves back to original shell (ground state)

13. The higher up the ball is, the more potential energy is has

14. ELECTRONS – CONFIGURATION/CHEMICAL PROPERTIES Chemical properties of an element depend on the number of valence electrons (outermost shell). Elements with the same number of valence electrons exhibit similar chemical properties

15. ELECTRONS – ORBITALS Electron Orbital – The area within an atom that electrons have the highest probability of being found. Each shell has a specific number of orbitals. Each orbital can contain up to 2 electrons. Different orbitals represented by different shapes Types of orbitals – s orbitals (spherical) – 1 per shell p orbitals (dumbbell-shaped) – 3 per shell d orbitals – 5 per shell

16. GROUND STATE ELECTRON CONFIGURATION The ground state electron configuration for oxygen is… 1s 2 2s 2 2p 4 For aluminum… 1s 2 2s 2 2p 6 3s 2 3p 1 What is the ground state electron configuration for… Sodium (Na) Sulfur (S) Potassium (K) Scandium (Sc)

17. CHEMICAL BONDING Most stable atoms are those with completely filled valence shells (usually eight valence electrons). Atoms interact with each other to form bonds in order to fill their valence shells. Types of chemical bonding Covalent Ionic Metallic

18. COVALENT BONDS Atoms fill their valence shells by sharing electrons with other atoms Occur between nonmetals Generally have a low to intermediate difference in electronegativity Polar covalent bonds – electrons are shared unequally; one atom has a more negative charge (δ-) and the other has a more positive charge (δ+). Nonpolar covalent bonds – electrons are shared equally and result in a net charge of zero. No difference in electronegativity. Bond order – ½(# of bonding electrons); single, double, or triple bonds. Bonding capacity = atom’s valence (number of unpaired valence electrons). The structure of a molecule depends on the type and number of covalent bonds present, as well as the number of paired electrons. Tetrahedral Bent

19. IONIC BONDS Atoms that have a high difference in electronegativity Atom with higher electronegativity strips electron away from atom with lower electronegativity to create two ions. Cation – positively charged ion Anion – negatively charged ion Resulting ions are attracted to each other, forming an ionic bond. Metals form cations; nonmetals form anions

20. IONIC COMPOUNDS Also referred to as salts Most commonly found in nature as crystals Formula only represents ratio of atoms (NaCl); individual molecules do not form Ionic compounds form lattice structures Cl – Na +

21. WEAK CHEMICAL BONDS Hydrogen bonds – Hydrogen atoms in a polar covalent bond are attracted to another electronegative atom Van der Waals Interactions – Due to delocalized electrons, molecules form temporary “hotspots” of charge that enable molecules to stick together.

22. MOLECULAR SHAPE AND FUNCTION Molecular shape depends on the positions of its atoms’ valence orbitals. In covalent bonds, overlapping of orbitals results in orbital hybridization. Specific shapes are associated with specific hybrid orbitals.

23. MOLECULAR SHAPE AND FUNCTION Molecular function depends on molecular shape Molecules interact with each other depending on shape Enzymes – Fit into specific receptor cites based on shape

24. CHEMICAL REACTIONS Molecular bonds break  Atoms rearrange themselves  New molecular bonds form

25. CHEMICAL REACTIONS Chemical Equilibrium – The rate of a chemical reaction is the same as the rate of the reverse reaction Concentration of reactants and products is NOT necessarily equal when chemical equilibrium is reached. Reactions still occur after equilibrium is reached, but there is no net change in the concentration of either reactants or products.

26. EXAM PRACTICE The atomic number of sulfur is 16. Sulfur combines with hydrogen by covalent bonding to form a compound, hydrogen sulfide. Based on the electron configuration of sulfur, we can predict that the molecular formula of the compound will be; A.) HSb.) HS 2 c.) H 2 Sd.)H 3 S 2 E.) H 4 S

27. EXAM PRACTICE Draw the Lewis Structure for the following molecules… HCN H 2 S CH 2 Br 2

28. EXAM PRACTICE Why is it important that we understand chemistry when we are studying biology? What coefficients must be placed in the blanks so that all atoms are accounted for in the products? ___C 6 H 12 O 6 → ___C 2 H 6 O + ___CO 2 ___S 8 + ___O 2 → ___SO 2 ___Al 2 (SO 4 ) 3 + ___Ca(OH) 2 → ___Al(OH) 3 + ___CaSO 4