1. 1 Chapter 2 The Chemical Context of Life

2. 2 Matter Takes up space and has mass Exists as elements (pure form) and in chemical combinations called compounds

3. 3 Elements Can’t be broken down into simpler substances by chemical reaction Composed of atoms Essential elements in living things include carbon C, hydrogen H, oxygen O, and nitrogen N making up 96% of an organism

4. 4 Other Elements A few other elements Make up the remaining 4% of living matter Table 2.1 Trace Elements

5. 5 Deficiencies (a) Nitrogen deficiency (b) Iodine deficiency (Goiter) If there is a deficiency of an essential element, disease results Figure 2.3

6. 6 Trace Elements Trace elements Are required by an organism in only minute quantities Minerals such as Fe and Zn are trace elements

7. 7 Compounds SodiumChloride Sodium Chloride + Are substances consisting of two or more elements combined in a fixed ratio Have characteristics different from those of their elements Figure 2.2

8. 8 Properties of Matter An element’s properties depend on the structure of its atoms Each element consists of a certain kind of atom that is different from those of other elements An atom is the smallest unit of matter that still retains the properties of an element

9. 9 Subatomic Particles Atoms of each element Are composed of even smaller parts called subatomic particles Neutrons, which have no electrical charge Protons, which are positively charged Electrons, which are negatively charged

10. 10 Subatomic Particle Location Protons and neutrons –Are found in the atomic nucleus Electrons –Surround the nucleus in a “cloud”

11. 11 Simplified models of an Atom Nucleus (a) (b) In this even more simplified model, the electrons are shown as two small blue spheres on a circle around the nucleus. Cloud of negative charge (2 electrons) Electrons This model represents the electrons as a cloud of negative charge, as if we had taken many snapshots of the 2 electrons over time, with each dot representing an electron‘s position at one point in time. Figure 2.4

12. 12 Atomic Number & Atomic Mass Atoms of the various elements Differ in their number of subatomic particles The number of protons in the nucleus = atomic number The number of protons + neutrons = atomic mass Neutral atoms have equal numbers of protons & electrons (+ and – charges)

13. 13 Atomic Number Is unique to each element and is used to arrange atoms on the Periodic table Carbon = 12 Oxygen = 16 Hydrogen = 1 Nitrogen = 17

14. 14 Atomic Mass Is an approximation of the atomic mass of an atom It is the average of the mass of all isotopes of that particular element Can be used to find the number of neutrons (Subtract atomic number from atomic mass)

15. 15 Isotopes Different forms of the same element Have the same number of protons, but different number of neutrons May be radioactive spontaneously giving off particles and energy May be used to date fossils or as medical tracers

16. 16 Energy Levels of Electrons An atom’s electrons Vary in the amount of energy they possess Electrons further from the nucleus have more energy Electron’s can absorb energy and become “excited” Excited electrons gain energy and move to higher energy levels or lose energy and move to lower levels

17. 17 Energy –Is defined as the capacity to cause change Potential energy - Is the energy that matter possesses because of its location or structure Kinetic Energy - Is the energy of motion

18. 18 Electrons and Energy The electrons of an atom –Differ in the amounts of potential energy they possess A ball bouncing down a flight of stairs provides an analogy for energy levels of electrons, because the ball can only rest on each step, not between steps. (a) Figure 2.7A

19. 19 Energy Levels Are represented by electron shells Third energy level (shell) Second energy level (shell) First energy level (shell) Energy absorbed Energy lost An electron can move from one level to another only if the energy it gains or loses is exactly equal to the difference in energy between the two levels. Arrows indicate some of the step-wise changes in potential energy that are possible. (b) Atomic nucleus Figure 2.7B

20. Thermodynamics and Biology First law of thermodynamics In any process, the total energy of the universe remains the same. It can also be defined as: for a thermodynamic cycle the sum of net heat supplied to the system and the net work done by the system is equal to zero. Second law of thermodynamics The entropy (useless energy) of an isolated system will tend to increase over time. In a simple manner, the second law states that "energy systems have a tendency to increase their entropy" rather than decrease it. 20

21. 21 Periodic table –Shows the electron distribution for all the elements Second shell Helium 2 He First shell Third shell Hydrogen 1 H 2 He 4.00 Atomic mass Atomic number Element symbol Electron-shell diagram Lithium 3 Li Beryllium 4 Be Boron 3 B Carbon 6 C Nitrogen 7 N Oxygen 8 O Fluorine 9 F Neon 10 Ne Sodium 11 Na Magnesium 12 Mg Aluminum 13 Al Silicon 14 Si Phosphorus 15 P Sulfur 16 S Chlorine 17 Cl Argon 18 Ar Figure 2.8

22. 22 Why do some elements react? Valence electrons –Are those in the outermost, or valence shell –Determine the chemical behavior of an atom

23. 23 Electron Orbitals An orbital –Is the three-dimensional space where an electron is found 90% of the time

24. 24 Covalent Bonds Figure 2.10 Sharing of a pair of valence electrons Examples: H 2 Hydrogen atoms (2 H) Hydrogen molecule (H 2 ) + + + + ++ In each hydrogen atom, the single electron is held in its orbital by its attraction to the proton in the nucleus. 1 When two hydrogen atoms approach each other, the electron of each atom is also attracted to the proton in the other nucleus. 2 The two electrons become shared in a covalent bond, forming an H 2 molecule. 3

25. 25 Covalent Bonding A molecule –Consists of two or more atoms held together by covalent bonds A single bond –Is the sharing of one pair of valence electrons A double bond –Is the sharing of two pairs of valence electrons Name (molecular formula) Electron- shell diagram Structural formula Space- filling model Hydrogen (H 2 ). Two hydrogen atoms can form a single bond. Oxygen (O 2 ). Two oxygen atoms share two pairs of electrons to form a double bond. HH O O Figure 2.11 A, B

26. 26 Compounds & Covalent Bonds Name (molecular formula) Electron- shell diagram Structural formula Space- filling model (c) Methane (CH 4 ). Four hydrogen atoms can satisfy the valence of one carbon atom, forming methane. Water (H 2 O). Two hydrogen atoms and one oxygen atom are joined by covalent bonds to produce a molecule of water. (d) H O H HH H H C Figure 2.11 C, D

27. 27 Covalent Bonding In a nonpolar covalent bond –The atoms have similar electronegativities –Share the electron equally

28. 28 Figure 2.12 This results in a partial negative charge on the oxygen and a partial positive charge on the hydrogens. H2OH2O –– O H H ++ ++ Because oxygen (O) is more electronegative than hydrogen (H), shared electrons are pulled more toward oxygen. In a polar covalent bond –The atoms have differing electronegativities –Share the electrons unequally Covalent Bonding

29. 29 Ionic Bonds In some cases, atoms strip electrons away from their bonding partners Electron transfer between two atoms creates ions Ions –Are atoms with more or fewer electrons than usual –Are charged atoms

30. 30 Ions An anion –Is negatively charged ions A cation –Is positively charged

31. 31 Ionic Bonding Cl – Chloride ion (an anion) – The lone valence electron of a sodium atom is transferred to join the 7 valence electrons of a chlorine atom. 1 Each resulting ion has a completed valence shell. An ionic bond can form between the oppositely charged ions. 2 Na Cl + Na Sodium atom (an uncharged atom) Cl Chlorine atom (an uncharged atom) Na + Sodium on (a cation) Sodium chloride (NaCl) Figure 2.13 An ionic bond –Is an attraction between anions and cations

32. 32 Ionic Substances Na + Cl – Figure 2.14 Ionic compounds –Are often called salts, which may form crystals

33. 33 Weak Chemical Bonds Several types of weak chemical bonds are important in living systems

34. 34 Hydrogen Bonds  – –  + +  + + Water (H 2 O) Ammonia (NH 3 ) O H H  + +  – – N H H H 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. ++ ++ Figure 2.15 A hydrogen bond –Forms when a hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom

35. 35 Van der Waals Interactions Van der Waals interactions –Occur when transiently positive and negative regions of molecules attract each other

36. 36 Weak Bonds Weak chemical bonds –Reinforce the shapes of large molecules –Help molecules adhere to each other

37. 37 Molecular Shape and Function Structure determines Function! The precise shape of a molecule –Is usually very important to its function in the living cell –Is determined by the positions of its atoms’ valence orbitals

38. 38 Orbitals & Covalent Bonds Space-filling model Hybrid-orbital model (with ball-and-stick model superimposed) Unbonded Electron pair 104.5° O H Water (H 2 O) Methane (CH 4 ) H H H H C O H H H C Ball-and-stick model H H H H (b) Molecular shape models. Three models representing molecular shape are shown for two examples; water and methane. The positions of the hybrid orbital determine the shapes of the molecules Figure 2.16 (b)

39. 39 Shape and Function Molecular shape –Determines how biological molecules recognize and respond to one another with specificity

40. 40 Morphine Carbon Hydrogen Nitrogen Sulfur Oxygen Natural endorphin (a) Structures of endorphin and morphine. The boxed portion of the endorphin molecule (left) binds to receptor molecules on target cells in the brain. The boxed portion of the morphine molecule is a close match. (b) Binding to endorphin receptors. Endorphin receptors on the surface of a brain cell recognize and can bind to both endorphin and morphine. Natural endorphin Endorphin receptors Morphine Brain cell Figure 2.17

41. 41 Chemical Reactions Chemical reactions make and break chemical bonds A Chemical reaction –Is the making and breaking of chemical bonds –Leads to changes in the composition of matter

42. 42 Chemical Reactions ReactantsReactionProduct 2 H 2 O2O2 2H2O2H2O + + Chemical reactions –Convert reactants to products

43. 43 Chemical Reactions Photosynthesis –Is an example of a chemical reaction Figure 2.18

44. 44