2. Atoms and Atomic Theory Atom - the smallest particle of a given type of matter. Atomic Theory – the idea that matter is made up of fundamental particles called atoms.

3. What discoveries led to the development of atomic structure? Democritus – Father of the Atom 400 B.C. atomos – “indivisible”

4. John Dalton – Father of the Modern Atom Dalton’s Atomic Theory. 1.) All matter is made of atoms. 2.) All atoms are indestructible and cannot be divided into smaller particles. 3.) All atoms of one element are exactly alike, but they are different from all other elements.

5. Subatomic Particles Electrons Protons Neutrons Nucleus

6. Electrons Symbol: e - Mass = 1/1837 amu 1 amu ~ the mass of a hydrogen atom Charge = (-1) J.J. Thomson – Discovered the electron using a cathode ray tube J.J. Thomson

7. Indirect Evidence Evidence providing only a basis for inference about the fact in dispute. Indirect evidence provided information about electrons. Diagram this experiment.

8. Protons Symbol: + p Mass = 1 amu Charge = (+1)

9. Neutrons Symbol: n 0 Mass = 1 amu and there is no charge

10. Thomson’s Model Thomson put both the electron and proton together and proposed the “Plum Pudding” model of the atom.

11. Nucleus Ernest Rutherford Using the “Gold Foil” experiment, Rutherford determines:Gold Foil 1. An atom is mostly empty space. 2. There is a nucleus in the middle of the atom containing the protons. 3. The electrons orbit a large distance away from the nucleus. Proposes the nuclear model of the atom.

12. Let’s Review http://www.brainpop.com/science/matterandchemist ry/atomicmodel/

13. How Big is an Atom? Scale of the Universe website: http://www.youtube.com/watch?v=fUAFqkS7y9M&fe ature=related Current Atomic Model: http://www.youtube.com/watch?v=xqNSQ3OQMGI Unit 1 Investigation III

14. Atomic Number Designates the number of protons in a nucleus of an atom. Each element has a characteristic atomic number. The number of electrons equals the atomic number in a neutral atom.

15. Mass Number Designates the total number of protons and neutrons in an atom. Number of neutrons = mass number - atomic number. Atoms of the same element can have different mass numbers.

16. Objective 2 Explain why isotopes differ.

17. Isotopes Atoms with the same number of protons, but different numbers of neutrons. Isotopes of an element have the same atomic number, but different mass numbers. Nuclear Symbol or isotopic Nuclear Symbol symbol – shows number of protons, neutrons and electrons in an atom. Isotopes may also be described by labeling the name following by the mass number (Example: carbon-12).

18. Atomic Mass Weighted average mass of the atoms in naturally occurring sample of an element. Masses are based off of the atomic mass unit (amu) defined as one twelfth the mass of a carbon-12 atom.

19. Atomic Mass Example In nature carbon is composed of 98.89% 12 C atoms and 1.11% 13 C atoms. 12 C has a mass of 12 amu and 13 C has a mass of 13.0034 amu. What is the average atomic mass of carbon?

20. Review John Dalton – Dalton’s Atomic Theory stating atoms are solid, indivisible spheres J. J. Thomson – Thomson Model or “Plum Pudding Model” stating electrons are dispersed through a sphere of positive charge Ernest Rutherford – Rutherford Model stating atoms have a small, dense, positively charged nucleus surrounded by electrons

21. Objective 3 Diagram the Bohr Model of the atom.

22. Bohr Model Niels Bohr (1913) – Bohr Model stating electrons travel around nucleus in energy levels. Uses rings to show the energy levels. The number of rings should match the row the element is in on the periodic table.

23. Bohr Model Diagram The protons in the nucleus are found by looking at the atomic number. The neutrons in the nucleus are found by subtracting the atomic number from the mass number. Electrons fill the shells from low to high until the correct number of electrons are added (number of electrons = number of protons). This model works for the first 20 elements.

24. Diagram Lithium-7

25. Diagram Nitrogen-14

26. Diagram Aluminum-27

27. Diagram Calcium-40

28. Objective 4 Identify the position of groups, periods, and different chemical families on the periodic table.

29. Periodic Table Dmitri Mendeleev – mid 1800’s Proposed a table for 70 elements based on mass and properties Henry Moseley – 1913 Determined the atomic number of elements and arranged the table in order of atomic number

30. How is the periodic table of elements arranged? Periods: Rows of the periodic table Groups: Columns of the periodic table Periodic Law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their chemical and physical properties.

31. Groupings to know on the Periodic Table Periodic Table Representative Elements Metals Non-metals Metalloids Transition Metals Inner-transition Metals Alkali Metals Alkaline Earth Metals Halogens Noble (inert) gases

32. Properties of Metals Metals: Bright metallic luster Solids are easily deformed Good conductors of electricity and heat

33. Properties of Nonmetals Nonmetals: Non-lustrous, various colors Solids may be hard or soft, usually brittle Poor conductors of electricity

34. Properties of Metalloids Metalloids have properties intermediate between metals and nonmetals.

35. Objective 5 Identify forces between atoms.

36. Lewis Dot Structures Show valence electrons. Look at the group number. This is how many dots you draw for each atom. Ex: PotassiumFluorineSulfurXenon PhosporousSiliconAluminumBeryllium

37. Bonding & Reactions The main goal of chemical boding and reacting is to fill the outermost energy level. This is called the octet rule. (8 electrons in the outermost shell, EXCEPT H, He, Li, Be, which need 2.)

38. Two Types of Bonds 1.) Ionic 2.) Covalent

39. Ionic Bonds When a metal ion with a positive charge sticks to a nonmetal ion with a negative charge. Also called salts.

40. Ions: an atom that has extra or missing electrons Metals lose their outer electrons, becoming positive. Ex: NaBe Al Nonmetals gain electrons in their outer shell, becoming negative. Ex: NO Cl

41. Where do these gained or lost electrons go? Ex: Sodium chloride Metal = sodiumnonmetal = chlorine Ex: Aluminum oxide Metal = aluminumnonmetal = oxygen

42. Use the valence electrons! To find the charge of an ion, look at the valence electrons and count how many they need to give away or add to reach a full energy level (8 electrons for most). Ex: What charge will the following ions have? sodiummagnesium phosphorouscarbon Brominesulfur

43. Forming the compound The compound is then balanced by crossing charges, reducing if possible. MetalChargeNonmetalCharge Balanced Compound Formula Li+C= Mg+F= K+O= Al+P= Brain Pop: Ions

44. Covalent Bonds Non-metals share electrons to fill outer shells. (1 bond = 2 electrons) Remember, nonmetals like to gain electrons. Since they both want to gain, they must share. Follow these rules: 1.) Look at the formula. Then add up total valence electrons needed for your drawing. 2.) Single bond all of the atoms, picking a center atom when possible. 3.) Fill every atoms energy level. Remember, most need 8 electrons, but H only needs 2. 4.) Count the electrons & erase some if you have too many. Move the electrons into double or triple bonds if needed.

45. Practice Cl 2 H 2 OH 2 O 2 CCl 4 CH 4 CO 2

46. Ionic or Covalent? (Metal + Nonmetal = Ionic Nonmetal + Nonmetal = Covalent) CompoundMetal + Nonmetal or Nonmetal + Nonmetal? Ionic or Covalent? NaCl H2O2 C6H12O6 MgF2 Al2O3 CO2 O2 Li3P

47. Lewis Dot Structures Practice N 2 F 2 OCl 2 O 2 CI 4 CBr 4 CO 2