The Chemistry of Life

1. The Chemistry of Life Chapter 3

2. Atoms • Organisms are chemical machines • one must know chemistry in order to understand biology • Any substance in the universe that has mass and occupies space is comprised of matter • MATTER • Anything that takes up space and has mass • Can exist as a liquid, solid, or gas • All matter is made up of atoms

3. Atoms • Atom = smallest particle a substance can be divided into that can retain its properties • All atoms have the same structure • at the core is a dense nucleus comprised of two types of subatomic particles • protons (positively charged) • neutrons (no associated charge) • orbiting the nucleus is a cloud of another subatomic particles • electrons (negatively charged)

4. An atom can be characterized by: Atomic Number the number of protons in the nucleus atoms with the same atomic number exhibit the same chemical properties and are considered to belong to the same element e.g. Carbon = C: atomic number = 6 Atomic Mass (mass number) the number of protons plus neutrons in the nucleus electrons have negligible mass (1/1840 dalton) e.g. C: atomic mass = 12.011 Atoms

5. Atomic Number andMass Note: Hydrogen is unique in that it has 1 proton, but 0 neutrons!

6. Atomic Symbol Atomic Mass = Number of Protons + Number of Neutrons Atomic Number = The Number of Protons in the Nucleus

7. Atoms • Electrons determine the chemical behavior of atoms • These subatomic components are the parts of the atom that come close enough to each other in nature to interact • Same charges repel each other • Opposite charges attract each other • Electrons can be shared

8. Atoms • Electrons are associated with energy • Potential energy: energy of position • e.g. Rollercoaster at top of peak • e.g. As electrons move away from core, they increase potential energy • The field of energy around an atom is arranged as levels called electron shells • Orbitals are the location electrons are most likely to be found within this volume of space

9. Electron Arrangement of Atoms 1 electron 6 electrons 7 electrons 8 electrons

10. Atoms • Electron shells have specific numbers of orbitals that may be filled with electrons • atoms that have incomplete electron orbitals tend to be more reactive • atoms will lose, gain, or share electrons in order to fill completely their outermost electron shell • these actions are the basis of chemical bonding

11. Atoms • How many electron shells? • How many electrons in first shell? • How many electrons in second shell? • How many electrons in third shell? • How many total electrons? • What is Atomic Number? • What is Atomic Mass? Sodium atom = Na

12. Basic building block of matter 92 naturally occurring elements Only 6 elements make up most of the body weight of organisms C Carbon H Hydrogen N Nitrogen O Oxygen P Phosphorus S Sulfur Elements

14. Ions • Ions – atoms that have gained or lost one or more electrons • Gaining an electron makes gives a negative charge • Losing an electron gives a positive charge • For Example: • Sodium ion has 11 protons, 10 electrons • Does this sodium ion have a positive or negative charge? • A negative ion could also form if an extra electron were added

15. Sodium Ion

16. Isotopes • Isotopes – atoms that have the same number of protons but different numbers of neutrons • most elements in nature exist as mixtures of different isotopes • C-12: 6 protons, 6 neutrons, 6 electrons • C-14: 6 protons, 8 neutrons, 6 electrons • Some isotopes are unstable and break up into particles with lower atomic numbers • this process is known as radioactive decay • Radioactive isotopes can be used in nuclear medicine and for dating fossils

17. Figure3.5 Isotopes of the element carbon

18. Molecules • A molecule is a group of atoms held together by energy • e.g. water (H2O), sodium chloride (NaCl), oxygen (O2) • The energy holding two atoms together is called a chemical bond • Atoms can interact in 3 ways: • 1. Share one or more electrons • 2. Accept extra electrons • 3. Donate electrons to another atom • There are 3 principal types of chemical bonds • Ionic • Covalent • Hydrogen

19. IONIC BONDS • Ionic bonds involve the attraction of opposite electrical charges • Transferof electrons from one atom to another • Molecules comprised of these bonds are often most stable as crystals • Remember: • An ION is an atom that with a charge Fig. 3.8(a)The formation of the ionic bond in table salt

20. COVALENT BONDS • – Covalent bonds form between two atoms when they share electrons • The number of electrons shared varies depending on how many the atom needs to fill its outermost electron shell • Covalent bonds are stronger than ionic bonds

21. COVALENT BONDING • A covalent bond • Each Hydrogen has 1 electron in shell • Sharing 2 electrons fills the shell, increasing stability • A double covalent bond • Sharing 2 pairs of electrons • Oxygen has total of 8 electrons (2 in inner shell, 6 in outer shell) • Sharing 2 more fills its outer shell, increasing stability

22. HYDROGEN BONDS • Hydrogen bonds form from covalent bonds created by an atom’s electronegativity • Createpartial charges in atoms that are unequally sharing electrons • Are weak bonds • Electronegativity is the tendency of one atom’s nucleus to better attract the shared electrons from another nucleus

23. Hydrogen bonding of Water Molecules

24. Water molecules contain two covalent bonds

25. HYDROGEN BONDS • Hydrogen bonds form in association with polar molecules • each atom with a partial charge acts like a magnet to bond weakly to another polar atom with an opposite charge • the additive effects of many hydrogen bonding interactions can add collective strength to the bonds Figure 3.10 Hydrogen bonding water molecules

26. Hydrogen Bonds Give Water Unique Properties • Water is essential for life • The chemistry of life is water chemistry! • Water is a polar molecule • The partial charges of hydrogen bonds creates polarity • Water can form hydrogen bonds • Hydrogen bonding confers on water many different special properties

27. Hydrogen Bonds Give Water Unique Properties • Heat Storage • Water temperature changes slowly and holds temperature well • This is due to the large number of H bonds many water molecules will form with each other, it takes a lot of energy to break them (and raise temperature) • Ice Formation • Few hydrogen bonds break at low temperatures • Water becomes less dense as it freezes because hydrogen bonds stabilize and hold water molecules farther apart • High Heat of Vaporization • At high temperatures, hydrogen bonds can be broken • water requires tremendous energy to vaporize because of all the hydrogen bonds that must be broken

28. Hydrogen Bonds Give Water Unique Properties • Water molecules are sticky • Cohesion– when one polar water molecule is attracted to another polar water molecule • Adhesion – when OTHER polar molecules water are attracted to a water molecule Figure 3.12

29. Hydrogen Bonds Give Water Unique Properties • Water is highly polar • in solution, water molecules tend to form the maximum number of hydrogen bonds • Hydrophilicmolecules are attracted to water and dissolve easily in it • these molecules are also polar and can form hydrogen bonds • Hydrophobic molecules are repelled by water and do not dissolve • these molecules are non-polar and do not form hydrogen bonds

30. Acids and Bases • When water ionizes, it releases an equal number of hydrogen ions (H+) and hydroxide ions (OH-).

31. Water Ionizes • The amount of ionized hydrogen from water in a solution can be measured as pH • The pH scale is logarithmic, which means that a pH scale difference of 1 unit actually represents a 10-fold change in hydrogen ion concentration • e.g. pH of 4 has 10x greater H+ concentration than pH of 5 • e.g. pH of 4 has 100x greater H+ concentration than pH of 6 • pH difference of 2; 10 x 10 = 100 pH = -log[H+]

32. The pH scale

33. Water Ionizes • Pure water has a pH of 7 • there are equal amounts of [H+] relative to [OH-] • Acid – any substance that dissociates in water and increases the [H+] • acidic solutions have pH values below 7 • Base – any substance that combines with [H+] when dissolved in water • basic solutions have pH values above 7

34. Water Ionizes • The pH in most living cells and their environments is fairly close to 7 • proteins involved in metabolism are sensitive to any pH changes • Acids and bases are routinely encountered by living organisms • from metabolic activities (i.e., chemical reactions) • from dietary intake and processing • Organisms use buffers to minimize pH disturbances

35. Water Ionizes • Buffer– a chemical substance that takes up or releases hydrogen ions • Buffers don’t remove the acid or the base affecting pH, but minimize their effect on it • Most buffers are pairs of substances, one an acid and one a base

36. Buffer Example • Carbonic acid and bicarbonate in human blood • Interact in a pair of reversible reactions • CO2 + H2O  H2CO3 • H2CO3 H+ + HCO3- • If H+ is added, HCO3- can pick up H+ added to form H2CO3 • If H+ is removed, it disassociates to release more H+ into blood