Chemistry for Bio 9

1. Chemistry for Bio 9

2. Which of the following is/are properties of life? • Cellular structure   • the ability to take in energy and use it   • the ability to respond to stimuli from the environment   • the ability to reproduce   • All of the choices are correct.  

3. Lecture outline • Chemistry- definition, scope, and relevance to biology • Classification of matter • The atom and subatomic particles • Chemical bonding & reactions • Chemistry of water • Acids, Bases, and the pH scale

4. Chemistry is the study of matter • Matter is anything that has mass and takes up space

5. Chemistry is relevant to Biological Concepts • Chemistry is the study of matter and its interactions • All Living things are made of matter • Biolgists are interested in: • Complex biological molecules • Chemical energy • Biochemical reactions • The chemical environment

6. Complex biological molecules • All living things are made of complex macromolecules • Chemical principles rule their assembly

7. Chemical energy Photosynthesis creates molecules rich in energy: • 6CO2(g)+ 6H2O(l) + sunlight  C6H12O6(s) + 6O2(g) • Earth has been transformed by chemical reactions peformed by living things

8. Biochemical reactions • All living things are collections of a vast number of chemical reactions • Even the simplest living things contain impossibly complex pathways

10. The Chemical Environment • The physical properties of water determine the fate of life on earth • pH, salinity and other chemical factors influence • Living things are profoundly influenced by their chemical environment

11. Chemical reactions performed by living things have transformed earth over billions of years of its history

12. Classification of matter Mixtures can be homogeneous or heterogeneous Mixtures can vary in composition of their ingredients Compounds are defined substances with proportional amounts of ingredients: water, carbon dioxide, etc. Elements cannot be broken down into ingredients by chemical processes

13. Basic principles of chemistry

14. The periodic table is an organized display of all the elements in the universe

15. The Structure of the Atom Subatomic particles- protons, neutrons electrons Orbitals and the nucleus

16. All matter is ultimately comprised of atoms • Atoms are the smallest individual unit of matter • Atoms are comprised of protons, neutrons and electrons Proton: Charge= +1, Mass= 1 Neutron: Chg= 0, mass= 1 Electron: Chg = -1, mass= ~0 Mass= p + n Charge = p – e (for normal atoms = zero)

17. Electron cloud 6e– 2e– LE 2-4a Nucleus 2 Protons 6 Protons Mass number = 4 Mass number = 12 2 Neutrons 6 Neutrons 2 Electrons 6 Electrons Helium atom Carbon atom

18. Atomic structure • Protons and electrons in the nucleus • Electrons orbit around • Bohr atom- classic model featuring electrons in “planetary” orbitals • Each orbit holds a determined number of electrons (first holds two, 2nd and 3rd hold eight

19. Electron cloud model • Currently accepted model of atomic structure • 90% probability cloud • Mostly empty space • Unfilled orbitals found in unstable, reactive elements • Therefore, orbitals influence bonding

20. Reading the Periodic Table

21. Elements are defined by the number of their protons • There are 92 naturally occurring elements • Many others have been synthesized Atomic number = # protons Atomic mass (mass number) = protons + neutrons of an individual atom Atomic weight= Naturally occurring average of isotopes of a substance

22. The number of neutrons in atoms of a single element is variable • Isotopes are variants of an element, differentiated by numbers of neutrons • Some isotopes are stable, others are radioactive

23. Some isotopes are common, others rare

24. Many Isotopes for an element can exist; radioisotopes are radioactive

25. Radioisotopes can be used in medical diagnosis- Radioisotopes of iodine target the thyroid gland

26. How is atomic weight different from atomic mass?

27. The sodium atom contains 11 electrons, 11 protons, and 12 neutrons. What is the mass number (atomic mass) of sodium? A. 0 B. 11   C. 22   • 23   • 34  

28. 96% of human tissue is comprised of 6 elements • Carbon, Hydrogen, Nitrogen, Oxygen, Phosphorous, Sulfur (CHNOPS) • 25 elements serve known functions in the body, incl. Ca, K, Na, Cl, Mg, Fe • Trace elements are essential, but in small quantities

29. Electrons in the outermost shell of an atom are called valence electrons

30. The nucleus of an atom contains • protons and neutrons.   • protons and electrons.   • only neutrons.   • only protons.   • only electrons.

31. Intramolecular Chemical Bonds: Ionic, Covalent, and the formation of molecules

32. Atoms are stable when their valence shells are filled with electrons • What atoms are these? • How could they satisfy their valence shells?

33. Noble gases have a stable electron structure • Their outer orbitals have a full complement of electrons • Noble gases are very unreactive

34. Elements combine in chemical reactions to form compounds • Molecules- 2 or more atoms combined in specific ways • Compounds- different elements in a molecule, in exact, whole-number ratios, joined by a chemical bond • 2 major kinds of intramolecular chemical bonds: Covalent (incl. polar and nonpolar) and Ionic

35. LE 2-7 In ionic bonding, an atom takes an electron from another atom, forming ions Transfer of electron Na+ Sodium ion Cl- Chloride ion Sodium chloride (NaCl) Na Sodium atom Cl Chlorine atom

36. Ions • Ions- Charged atoms or molecules • Anion- negative ion • Cation- positive ion • Ionization- reaction producing ions • Salt- a neutral compound comprised of ions

37. LE 2-7a-2 Na+ Sodium ion Cl- Chloride ion Sodium chloride (NaCl)

38. LE 2-7b Na+ Cl-

39. A compound • A) is a pure element.   • B) is less common than a pure element.   • C) contains two or more elements in a fixed ratio.   • D) is exemplified by sodium.   • E) is a solution.  

40. Electronegativity and its effect on chemical bonds Ionic bonds, covalent bonds, and intermolecular forces

41. Electronegativity values can predict how atoms will bond

42. In covalent bonds, electrons do not always share time between bond partners equally Comparisions of electronegativity • Na: 0.9 • H: 2.1 • C: 2.5 • N: 3.0 • Cl: 3.0 • O: 3.5

43. Electronegativity = “electron greediness” • Large differences in polarity of atoms in a bond creates polar molecules • Relative electronegativity of Hydrogen and oxygen makes water a very polar molecule • Polar- regions of positivity and negativity • By Oxygen, water is (slightly) negative • By Hydrogens, water is (slightly) positive

44. Intermolecular forces and the chemistry of water Polarity and hydrophilicity, Nonpolarity and hydrophobicity, hydrogen bonding, and the chemistry of water

45. Water is a “universal solvent” and dissolves many polar and ionic compounds (“like dissolves like”)

46. The polarity of water allows hydrogen bonding • Polar regions of water molecules interact to form hydrogen bonds • Hydrogen bonds: weak/temporary intermolecular forces between positive and negative regions

47. Other molecules can engage in H-bonding, w/ water or other substances

48. Hydrogen bonds hold together the two strands of a DNA double helix

49. Hydrogen bonding in water determine many of water’s unique properties • H-bonds can form a lattice (ice) • H-bonds require much energy (usually heat) to break • H-bonds give water surface tension Hydrogen bond

50. 0 Hydrogen Bonds help to make water cohesive, allowing water surface tension and capillary action