Chemistry

1. Chemistry General Biology 1114

2. Living Things Are Composed of Matter Matter has mass and takes up space

3. Properties of Matter • Physical Properties • Observed and measured without changing the identity of the matter. (mass, weight, volume) • 2. Chemical Properties • Describe a substances ability to change into a new substance as the result of a chemical change. (rusting, radioactivity)

4. Phases of Matter • Solid • Liquid • Gas

5. Democritus (Greek Philosopher) • World is Made of Empty Space and Particles (Atoms) • Different Type of Particles for Different Things • Discontinuous Theory • Theory Said the Atom is the Basic Unit of Matter

6. Aristotle • Said Matter Was Continuous and Not Made Up of Smaller Particles • Called the Continuous Substance Hyle • Proposed the Continuous Theory

7. Atoms • Smallest Unit of Matter Retaining its Properties • Atoms 0.1 nm • Organelles 0.1 micron • Cells 0.1 mm

8. Parts of the Atom • Protons • Neutrons • Electrons Subatomic Particles were discovered during the 1800’s.

9. Protons • Nucleus • +1 Charge • 1 amu = 1 Dalton • Defines the Atomic Number

10. Neutron • Nucleus • No Charge • 1 amu = 1 Dalton • Numbers Vary • Atomic Mass – Atomic Number = # Neutrons

11. Isotopes Isotopes result when an element can have a varying number of neutrons. Each isotope has the same chemical properties but different physical properties

12. Natural Very Rare and Magnetic Even Rarer and Radioactive

13. Electrons • Orbit Nucleus • -1 Charge • Mass is Negligible • Speed of Light • Electrically Neutral Atoms Have Equal Numbers of Protons and Electrons

15. If an atom gains or loses an electron it becomes an ion. When Sodium loses an electron it becomes a positive ion – a cation If an element gains an electron it becomes a negative ion – an anion.

16. Atomic Number • The number of protons an atom possesses. • It is unique and characteristic for each element.

17. Atomic Mass • Sometimes called the atomic weight • Protons + Neutrons = Atomic Mass

18. Practice • How many protons are in Zinc (Zn)?

19. Practice • How many protons are in Zinc (Zn)? 30

20. Practice • How many electrons are in Lithium (Li)?

21. Practice • How many electrons are in Lithium (Li)? 3

22. Practice • How many neutrons are in Gold (Au)?

23. Practice • How many neutrons are in Gold (Au)? 118

24. Practice • How many protons are in Lead (Pb)?

25. Practice • How many protons are in Lead (Pb)? 82

26. Practice • How many electrons are in Neon (Ne)?

27. Practice • How many electrons are in Neon (Ne)?

28. Practice • How many neutrons are in Chlorine (Cl)?

29. Practice • How many neutrons are in Chlorine (Cl)? 18

30. Practice • What is the atomic number of Hydrogen (H)?

31. Practice • What is the atomic number of Hydrogen (H)? 1

32. Practice • What is the atomic mass of Copper (Cu)?

33. Practice • What is the atomic mass of Copper (Cu)? 64 amu

34. Dimitri Mendeleev Russian Chemist First to be partially successful in arranging the known elements into a periodic table Arranged elements by their atomic masses Later rearranged by Mosley according to increasing atomic numbers

35. Periodic Table of the Elements A Chemistry Reference Tool

36. What is the Periodic Table? • Chemistry Reference • Array of Known Elements • Left to Right • Top to Bottom • Increasing Atomic Number • Contains Elements (Substances which cannot be broken down into simpler substances by a chemical reaction.)

37. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 Grouping Within the Periodic Table 1 H He Hydrogen Helium 2 Li Be B C N O F Ne Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon 3 Na Mg Al Si P S Cl Ar Sodium Magnesium Aluminum Silicon Phosphorus Sulfur Chlorine Argon 4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton 5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon 6 Cs Ba * Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Cesium Barium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury Thallium Lead Bismuth Polonium Astatine Radon 7 Fr Ra ** Rf Db Sg Bh Hs Mt Uun Uuu Uub Francium Radium Unnilquadium Unnilpentium Unnilhexium Unnilseptium Unniloctium Unnilennium Ununnilium Unununium Ununbium * La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium ** Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Element Groups (Families) Alkali Earth Alkaline Earth Transition Metals Rare Earth Other Metals Metalloids Non-Metals Halogens Noble Gases Rows are Periods Columns are Groups or Families

38. 105-109 Elements 92 Naturally Occurring 20 Used by Living Organisms 5 Account for 99% of all Body Mass

39. Elements Found in the Human Body 65% Oxygen “O” 19% Carbon “C” 10% Hydrogen “H” 3% Nitrogen “N” 2% Calcium “Ca” 1% Trace Elements (Cu, Fe, Zn, Mg)

40. Group 1 • Alkali Metal Group • Strong Metals • Unusually Soft • Very Reactive Toward Oxygen and Water Vapor • Must be Stored Under an Inert Liquid

41. Group 2 • Alkaline Earth Metals • Combine with Oxygen to Form Bases • Not as Soft as Group 1 • Not as Reactive as Group 1

42. Group 3-12 • Transition Metals • Can Combine in a Variety of Ways • Not Very Predictable

43. Group 17 • Halogens • Combine with Elements to From Salts

44. Group 18 • Noble Gases • Inert Gases • Do Not Tend to Combine

45. Chemical BondingMolecules Join Together to Form Compounds • Octet Rule • Atoms are more stable when the outer shell is full. • Valency • The number of electrons to lose or gain to obtain a full outer shell. • Monovalent 1 Electron • Divalent 2 Electrons • Trivalent 3 Electrons

46. Bond Formation Requires Energy • Types of energy (Potential, Kinetic, etc.) • Energy is Measured in Joules • Law of Conservation of Energy Energy is neither created nor destroyed. • Law of Conservation of Mass/Energy Mass and Energy can be interconverted, but the total amount in the universe is constant • E = mc2 • Types of Energy Reactions 1. Exothermic = Heat Produced 2. Endothermic = Heat Absorbed

47. Bond Types

48. Ionic Bonds are formed when atoms become ions by gaining or losing electrons.

49. The sodium atom on the left has only one electron (blue) in its outermost orbit so it is unstable. To achieve a stable electron shell, sodium atoms like to GIVE AWAY or DONATE its outermost electron to other atoms. Chlorine has seven electrons in its outermost orbit and only needs ONE to achieve stability. When sodium donates its electron to chlorine, both achieve orbital bliss and stability. However, since electrons carry a negative charge, the sodium atom becomes POSITIVELY charged when it donates an electron and chlorine becomes NEGATIVELY charged upon gaining an electron. The two charged atoms are now called IONS and they strongly attract each other forming the compound SALT or NaCl.

50. Covalent bonds in water. In covalent bonds one electron spends part of its time orbiting around one atom and the rest of its time orbiting around the other atom in the compound. Biological molecules like proteins, nucleic acids and lipids are held together mostly by covalent bonds. In the water molecule shown above, note that the two types of atoms SHARE their electrons (yellow oval) so that both fill their outermost orbits to the maximum numbers of 2 & 8 respectively.