Chapter 2

1. Chapter 2 The Chemical Context of Life

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

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

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

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

7. Compounds + Sodium Chloride 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. 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. 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. Subatomic Particle Location • Protons and neutrons • Are found in the atomic nucleus • Electrons • Surround the nucleus in a “cloud”

11. Simplified models of an Atom Cloud of negative charge (2 electrons) Electrons Nucleus 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) Figure 2.4

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. 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. 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. 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. 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. 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

18. Electrons and Energy (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 • The electrons of an atom • Differ in the amounts of potential energy they possess

19. Energy Levels 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 • Are represented by electron shells

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.

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

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

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

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

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

26. 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

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

28. Covalent Bonding 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 d+ d+ H2O • In a polar covalent bond • The atoms have differing electronegativities • Share the electrons unequally Figure 2.12

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. Ions • An anion • Is negatively charged ions • A cation • Is positively charged

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

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

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

34. Hydrogen Bonds H Water (H2O) O 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. H  +  – Ammonia (NH3) N H H d+ + H 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. Van der Waals Interactions • Van der Waals interactions • Occur when transiently positive and negative regions of molecules attract each other

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

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. 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 (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 H H Methane (CH4) Figure 2.16 (b)

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

40. Nitrogen Carbon Hydrogen Sulfur Oxygen Natural endorphin Morphine (a) Structures of endorphin and morphine. The boxed portion of the endorphin molecule (left) binds toreceptor 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

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

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