Chapter 2 Atoms and Molecules: The Chemical Basis of Life

1. Chapter 2Atoms and Molecules: The Chemical Basis of Life

2. Chemistry and Life • All organisms share fundamental similarities in their chemical composition and basic metabolic processes • The structure of atoms determines the way they form chemical bonds to produce complex compounds • molecular biology • Chemistry and physics of the molecules that constitute living things

3. Inorganic and Organic Compounds • inorganic compounds • Small, simple substances • Biologically important groups include water, simple acids and bases, and simple salts • organic compounds • Generally large, complex carbon-containing compounds • Typically, two or more carbon atoms are bonded to each other to form the backbone, or skeleton, of the molecule

4. 2.1 ELEMENTS AND ATOMS LEARNING OBJECTIVES: • Name the principal chemical elements in living things and provide an important function of each • Compare the physical properties (mass and charge) and locations of electrons, protons, and neutrons. Distinguish between the atomic number and the mass number of an atom • Define the terms orbital and electron shell; relate electron shells to principal energy levels

5. Elements • elements • Substances that can’t be broken down into simpler substances by ordinary chemical reactions • Each element has a chemical symbol (Example: C for carbon) • Four elements (oxygen, carbon, hydrogen, and nitrogen) make up more than 96% of the mass of most organisms • Calcium, phosphorus, potassium, and magnesium, are present in smaller quantities • Iodine and copper are trace elements

6. Functions of Elements in Organisms Table 2-1, p. 27

7. Atoms and Matter • atom • Smallest unit of an element that retains that element’s chemical properties • Made up of tiny subatomic particles of matter • matter • Anything that has mass and takes up space

8. Subatomic Particles • There are three basic types of subatomic particles: • Anelectroncarries a unit of negative electric charge • Aprotoncarries a unit of positive charge • Aneutronis an uncharged particle • Protons and neutrons compose the atomic nucleus • Electronsmove rapidly around the atomic nucleus • In an electrically neutral atom, the number of electrons equals the number of protons

9. Atomic Number and the Periodic Table • Every element has a fixed number of protons in the atomic nucleus (atomic number) which determines an atom’s identity and defines the element • The periodic tableis a chart of the elements arranged in order by atomic number and chemical behavior • Bohr models represent the electron configurations of elements as a series of concentric rings

10. The Periodic Table

11. Chemical symbol Atomic number Chemical name Number of e– in each energy level Fig. 2-1, p. 28

12. Atomic Mass • The mass of a subatomic particle is expressed in terms of the atomic mass unit (amu)ordalton • One amu equals the approximate mass of a single proton or a single neutron; an electron is about 1/1800 amu • The atomic mass of an atom equals the total number of protons and neutrons, expressed in amus or daltons

13. Characteristics of Subatomic Particles Particle Charge ~Mass Location Proton Positive 1 amu Nucleus Neutron Neutral 1 amu Nucleus Electron Negative ~1/1800 amu Outside nucleus

14. Isotopes • Most elements consist of a mixture of atoms with different numbers of neutrons and different masses • isotopes • Atoms of the same element (having the same number of protons and electrons) with varying numbers of neutrons • The mass of an element is expressed as an average of the masses of its isotopes

15. Isotopes of Carbon

16. Radioisotopes • Some isotopes are unstable and tend to break down (decay) to a more stable isotope (usually a different element) • radioisotope • Unstable isotope that emits radiation as it decays • Example:14C decays to 14N when a neutron decomposes to form a proton and a fast-moving electron • Radioactive decay can be detected by autoradiography, on photographic film

17. Radioisotopes in Biology • Radioisotopes such as 3H (tritium), 14C, and 32P can replace normal molecules and are used as tracers in research • In medicine, radioisotopes are used for both diagnosis (such as thyroid function or blood flow) and treatment (such as cancer)

18. Autoradiography • Tritium (3H) incorporated into the DNA of a fruit fly

19. Atomic Orbitals and Energy • Electrons move through characteristic regions of 3-D space (orbitals), each containing a maximum of 2 electrons • The energy of an electron depends on the orbital it occupies • Electrons in orbitals with similar energies (the same principal energy level) make up an electron shell • Electrons farther from the nucleus generally have greater energy than those closer to the nucleus

20. Valence Electrons • The most energetic electrons (valence electrons) occupy the valence shell, represented as the outermost concentric ring in a Bohr model • Valence electrons participate in chemical reactions • An electron can move to a higher orbital by receiving more energy, or give up energy and sink to a lower orbital • Changes in electron energy levels are important in energy conversions in organisms

21. Atomic Orbitals

22. Nucleus (a) The first principal energy level contains a maximum of 2 electrons, occupying a single spherical orbital (designated 1s). The electrons depicted in the diagram could be present anywhere in the blue area. 1s Fig. 2-4a, p. 30

23. 2s 2py 2px 2pz (b) The second principal energy level includes four orbitals, each with a maximum of 2 electrons: one spherical (2s) and three dumbbell-shaped (2p) orbitals at right angles to one another. Fig. 2-4b, p. 30

24. z 1s 2s y 2py 2px x 2pz (c) Orbitals of the first and second principal energy levels of a neon atom are shown superimposed. Note that the single 2s orbital plus three 2p orbitals make up neon's full valence shell of 8 electrons. Compare this more realistic view of the atomic orbitals with the Bohr model of a neon atom at right. Fig. 2-4c, p. 30

25. (d) Neon atom (Bohr model) Fig. 2-4d, p. 30

26. ANIMATION: The shell model of electron distribution To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE

27. Key Concepts 2.1 • Carbon, hydrogen, oxygen, and nitrogen are the most abundant elements in living things

28. ANIMATION: Atomic number, mass number To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE

29. ANIMATION: Electron arrangement in atoms To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE

30. Animation: Electron distribution

31. ANIMATION: Subatomic particles To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE

32. 2.2 CHEMICAL REACTIONS LEARNING OBJECTIVES: • Explain how the number of valence electrons of an atom is related to its chemical properties • Distinguish among simplest, molecular, and structural chemical formulas • Explain why the mole concept is so useful to chemists

33. Valence Electrons • Chemical behavior of an atom is determined by the number and arrangement of its valence electrons • Atoms with full valence shells are unreactive • When the valence shell is not full, an atom tends to lose, gain, or share electrons to achieve a full outer shell • Elements in the same vertical column (group) of the periodic table have similar chemical properties

34. Compounds and Molecules • Two or more atoms may combine chemically • A chemical compound consists of atoms of two or more different elements combined in a fixed ratio • Two or more atoms joined very strongly form a stable molecule • Example: H20 (water) is a molecular compound

35. Chemical Formulas • A chemical formula is a shorthand expression that describes the chemical composition of a substance • In a simplest formula (empirical formula), subscripts give the smallest ratios for atoms in a compound (e.g. NH2) • A molecular formulagives the actual numbers of each type of atom per molecule (e.g. N2H4) • A structural formula shows the arrangement of atoms in a molecule (e.g. water, H—O—H)

36. The Mole • The molecular mass of a compound equals the sum of the atomic masses of the component atoms of a single molecule • The amount of a compound whose mass in grams is equivalent to its molecular mass is 1 mole (mol) • Example: • Molecular mass of water (H2O) is (hydrogen: 2 × 1 amu) + (oxygen: 1 × 16 amu) = 18 amu • 1 mol of water is 18 grams (g)

37. The Mole (cont.) • 1 mol of any substance has exactly the same number of atoms or molecules: 6.02 × 1023 (Avogadro’s number) • Avogadro’s number allows scientists to calculate the number of atoms or molecules in sample simply by weighing it • A 1 molar solution (1 M) contains 1 mol of a substance dissolved in a total volume of 1 liter (L)

38. Chemical Reactions • Chemical reactions, such as the reaction between glucose and oxygen, are described by chemical equations: C6H12O6 + 6 O2 →6 CO2 + 6 H2O + energy • Substances that participate in the reaction (reactants)are written on the left side of the arrow • Substances formed by the reaction (products)are written on the right side

39. Chemical Reactions (cont.) • Many reactions proceed forward and reverse simultaneously • At dynamic equilibrium, the rates of forward and reverse reactions are equal CO2 + H2O ↔ H2CO3 • When this reaction reaches equilibrium, there will be more reactants (CO2 and H2O) than product (H2CO3)

40. Key Concepts 2.2 • The chemical properties of an atom are determined by its highest-energy electrons, known as valence electrons

41. ANIMATION: Chemical bookkeeping To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE

42. Animation: Covalent bonds

43. 2.3 CHEMICAL BONDS LEARNING OBJECTIVE: • Distinguish among covalent bonds, ionic bonds, hydrogen bonds, and van der Waals interactions • Compare them in terms of the mechanisms by which they form and their relative strengths

44. Chemical Bonds • Atoms can be held together by chemical bonds • Valence electrons dictate how many bonds an atom can form • bond energy • Energy necessary to break a chemical bond • Two types of strong chemical bonds: covalent and ionic

45. Covalent Bonds • Covalent bonds involve sharing electrons between atoms in a way that fills each atom’s valence shell • A molecule consists of atoms joined by covalent bonds • Example: hydrogen gas (H2) • Unlike atoms linked by covalent bonds form a covalent compound

46. Lewis Structure • A simple way of representing valence electrons is to use dots placed around the chemical symbol of the element: • Oxygen (6 valence electrons) shares electrons with two hydrogen atoms to complete its valence shell of 8 – each hydrogen atom completes a valence shell of 2

47. Carbon Bonds • Carbon has 4 electrons in its valence shell, all of which are available for covalent bonding (e.g. methane, CH4) • Each orbital holds a maximum of 2 electrons

48. Single, Double, and Triple Covalent Bonds • When one pair of electrons is shared between two atoms, the covalent bond is called a single covalent bond • A double covalent bondis formed when two pairs of electrons are shared(represented by two parallel solid lines) • A triple covalent bond is formed when three pairs of electrons are shared (represented by three parallel solid lines)

49. Electron Sharing in Covalent Compounds

50. Hydrogen (H) Hydrogen (H) Molecular hydrogen (H2) or H H (a) Single covalent bond formation. Two hydrogen atoms achieve stability by sharing a pair of electrons, thereby forming a molecule of hydrogen. In the structural formula on the right, the straight line between the hydrogen atoms represents a single covalent bond. Fig. 2-5a, p. 33