Chapter 2

1. Chapter 2 The Chemical Context of Life 1

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

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 3

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

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

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

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

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 8

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 9

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

11. Simplified models of an Atom Cloud of negative charge (2 electrons) Electrons Nucleus Figure 2.4 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. (a) In this even more simplified model, the electrons are shown as two small blue spheres on a circle around the nucleus. (b) 11

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) 12

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 13

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) 14

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 15

16. Radioactive tracers for diagnostic purposes 16

17. Other uses ·Can be used in medicine to treat tumors Cancerous throat tissue Figure 2.6 18

18. 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 19

19. Energy ·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 20

20. Electrons and Energy ·The electrons of an atom ·Differ in the amounts of potential energy they possess (a) 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. Figure 2.7A 21

21. Energy Levels ·Are represented by electron shells Third energy level (shell) Second energy level (shell) Energy absorbed First energy level (shell) Energy lost Atomic nucleus (b) 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. Figure 2.7B 22

22. Electron Configuration and Chemical Properties ·The chemical behavior of an atom ·Is defined by its electron configuration and distribution ·K (2e-) ·L-M (8e-) 23

23. Periodic table ·Shows the electron distribution for all the elements Helium 2He Atomic number 2 He 4.00 Hydrogen 1H Element symbol Atomic mass First shell Electron-shell diagram Lithium 3Li Beryllium 4Be Boron 3B Carbon 6C Nitrogen 7N Oxygen 8O Fluorine 9F Neon 10Ne Second shell Aluminum 13Al Chlorine 17Cl Argon 18Ar Sodium 11Na Magnesium 12Mg Silicon 14Si Phosphorus 15P Sulfur 16S Third shell Figure 2.8 25

24. Why do some elements react? ·Valence electrons ·Are those in the outermost, or valence shell ·Determine the chemical behavior of an atom 26

25. Covalent Bonds ·Sharing of a pair of valence electrons ·Examples: H2 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. 1 + + 2 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. 3 + + Hydrogen molecule (H2) Figure 2.10 30

26. 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 31

27. Multiple Covalent Bonds Name (molecular formula) Electron- shell diagram Space- filling model Structural formula (a) Hydrogen (H2). Two hydrogen atoms can form a single bond. H H (b) Oxygen (O2). Two oxygen atoms share two pairs of electrons to form a double bond. O O Figure 2.11 A, B 32

28. Compounds & Covalent Bonds Name (molecular formula) Electron- shell diagram Space- filling model Structural formula (c) Water (H2O). Two hydrogen atoms and one oxygen atom are joined by covalent bonds to produce a molecule of water. H O H (d) Methane (CH4). Four hydrogen atoms can satisfy the valence of one carbon atom, forming methane. H H H C H Figure 2.11 C, D 33

29. Covalent Bonding ·Electronegativity ·Is the attraction of a particular kind of atom for the electrons in a covalent bond ·The more electronegative an atom ·The more strongly it pulls shared electrons toward itself 34

30. Covalent Bonding ·In a nonpolar covalent bond ·The atoms have similar electronegativities ·Share the electron equally 35

31. Covalent Bonding ·In a polar covalent bond ·The atoms have differing electronegativities ·Share the electrons unequally Because oxygen (O) is more electronegative than hydrogen (H), shared electrons are pulled more toward oxygen. d This results in a partial negative charge on the oxygen and a partial positive charge on the hydrogens. O H H Figure 2.12 d+ d+ H2O 36

32. 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 37

33. Ions ·An anion ·Is negatively charged ions ·A cation ·Is positively charged 38

34. Ionic Bonding ·An ionic bond ·Is an attraction between anions and cations 2 1 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) Figure 2.13 Sodium chloride (NaCl) 39

35. Ionic Substances ·Ionic compounds ·Are often called salts, which may form crystals Na+ Cl– Figure 2.14 40

36. Weak Chemical Bonds ·Several types of weak chemical bonds are important in living systems 41

37. Hydrogen Bonds ·A hydrogen bond ·Forms when a hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom d d + d + 42

38. Van der Waals Interactions ·Van der Waals interactions ·Occur when transiently positive and negative regions of molecules attract each other 43

39. Weak Bonds ·Weak chemical bonds ·Reinforce the shapes of large molecules ·Help molecules adhere to each other 44

40. 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 45

41. Orbitals & Covalent Bonds Hybrid-orbital model (with ball-and-stick model superimposed) Ball-and-stick model Space-filling model Unbonded Electron pair O O H H H H 104.5° Water (H2O) H H C C H H H H H H Methane (CH4) (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) 47

42. Shape and Function ·Molecular shape ·Determines how biological molecules recognize and respond to one another with specificity 48

43. Nitrogen Carbon Hydrogen Sulfur Oxygen Natural endorphin Morphine (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. Natural endorphin Morphine Endorphin receptors Brain cell (b) Binding to endorphin receptors. Endorphin receptors on the surface of a brain cell recognize and can bind to both endorphin and morphine. Figure 2.17 49

44. 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 50

45. Chemical Reactions ·Chemical reactions ·Convert reactants to products + 2 H2O 2 H2 + O2 Reactants Reaction Product 51

46. Chemical Reactions ·Photosynthesis ·Is an example of a chemical reaction Figure 2.18 52

47. Chemical Reactions ·Chemical equilibrium ·Is reached when the forward and reverse reaction rates are equal 53

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