1. UNIT 2 Atomic Structure and The Periodic Table (Chapters 4,5,6)

2. The smallest particle of an element that retains the properties of that element.

3. Atoms consist of electrons, protons and neutrons. These are called subatomic particles.

4. Sub Atomic Particles Electrons (e-):   negatively charged   revolve around the nucleus in electron clouds   mass much less than protons or neutrons (mass = 0 a.m.u)   Up to 2 electrons on 1 st energy level, up to 8 on each level after 1 st.

5. Protons (p + ):   positive charge   inside the nucleus   Mass equals 1 a.m.u   Mass is 1840 times more than e- Neutrons (n 0 ):  Neutral (no) charge   inside the nucleus   mass equals 1 a.m.u.

6.  Central core of atom  Composed of  Composed of protons and neutrons   Tiny (compared to atom as a whole)   Very Dense   Contains most of the atoms mass Nucleus

8. ATOMIC NUMBER Number of protons (p+) in the nucleus  p+ = e- (in a neutral atom)  Each element contains a different # of protons (what makes each element different)  “the element’s address on the periodic table”

9. MASS NUMBER  Total number of protons & neutrons in an atom  n 0 + p + = Mass #  Remember (atomic # = # of p+)  Mass # - atomic # = n 0

10. ISOTOPES Atoms of the same element that have: same number of protons different number of neutrons different masses

11. Atomic # Mass # Chemical Symbol How many neutrons does this atom have? NOTE: Some periodic tables list the mass # on top and atomic # on the bottom while other periodic tables do the reverse. Remember, the smaller of the two numbers is always the atomic number!!

12. For about 50 years past the time of John Dalton, the atom was considered a solid indivisible mass. The later discovery of subatomic particles shattered this theory.

13. Dalton’s Atomic Theory 1. atoms 1. Elements are composed of tiny particles called atoms. 2. 2. Atoms of the same element are identical. Different elements have different atoms. 3. 3. Atoms of different elements can physically or chemically combine. 4. 4. Chemical Reactions (RXNs) occur when atoms are separated, joined, or rearranged. 5. 5. Atoms of one element never change into atoms of another element as a result of a chemical reaction.

14. He discovered the electron and proposed the plum-pudding model of the atom. This model said nothing about the number of protons and neutrons.

15. He discovered the nucleus and proposed a nuclear atom in which electrons surround a dense nucleus. He thought of the rest of the atom as empty space.

16. He had nothing to do with discovering a model of the atom but really likes teaching chemistry. (Just checking to see if you are paying attention!)

17. A student of Rutherford’s, he proposed that electrons are arranged in orbits around the nucleus. He explained that the electrons do not fall into the nucleus because they have fixed energy levels.

18. Development of Periodic Table  70 elements known by 1800s  Dmitri Mendeleev discovered a way to systematically & logically group elements  What are some ways elements are grouped?

19. Mendeleev’s Work  Arranged Elements in columns, so elements w/ similar properties were side by side  Arranged in columns of increasing Atomic Mass  Left blank spaces in table because there were no known elements w/ the correct properties and mass for that space.

20. Moseley’s Periodic Table Corrections  Put table according to atomic number instead of atomic mass

21. PERIODIC LAW 1) Elements arranged in order of increasing atomic number 2) Periodic Repetition of physical & chemical properties 3) Elements with same properties in columns

22. Modern Periodic Table  Periods – horizontal rows going across the periodic table  7 periods (each has its own energy level)  Properties of elements change as you move across  Pattern of properties repeats when you move from one period to the next.

23.  Groups/Families– elements in the periodic table going down in columns.  Similar chemical & physical properties  Classified by number and/or letter Ex: 1A, 2A, 3A… Ex: 1A, 2A, 3A… Representative Elements = Group A Transition Metals = Group B

24. METALS a) Left side of periodic table (except H) b) High electrical conductivity c) High Luster (very shiny when clean) d) Ductile (drawn into wires) e) Solid @ room temperature (except Hg)

25. 1) Alkali Metals: Group 1A (not H) Very Reactive Very Reactive 2) Alkaline Earth Metals: Group 2A 3) Transition Metals: Group B Middle of the periodic table Middle of the periodic table

26. NON-METALS a) Upper right corner of periodic table b) Non-lustrous (not shiny) c) Poor conductors of electricity d) Some are gases at room temp.(O,N) --others brittle solids (S)

27. 1) Halogens – group 7A 1) Very Reactive 2) 7 e- on outer energy level 2) Noble Gases -- group 8A 1) Unreactive – Few to no chemical RXNs 2) Full outer energy level

28. Semi-Metals (metalloids)  Divides metals from nonmetals  properties intermediate between metals and nonmetals  Ex: Si, Ge,...  “stair case” on right Side of periodic table

29. The Modern Atom  Quantum Mechanical Model:   The Energy level of an electron is the region around the nucleus where the electron is likely to be moving.

30.   Each period has its own principle energy level   principle energy levels are assigned values in order of increasing energy: n = 1, 2, 3, 4, and so forth

31. One Important Rule  Aufbau Principle:  Aufbau Principle: Electrons enter orbitals of lowest energy first.

32. Energy Sublevels   each principal energy level has energy sublevels.   sublevels are called atomic orbitals and are represented by the letters: s, p, d, f.

33. n =1 n = 2 n = 3 n = 4 1 2 3 4 s s, p s, p, d s, p, d, f Principle NRG Level Number of Sublevels Orbital Types

34. An s orbital can contain 2 electrons. A p orbital can contain 6 electrons A d orbital can contain 10 electrons. A f orbital can contain 14 electrons. Each sub-orbital can hold a different number of electrons

35. s p d f

36. 3 QUESTIONS TO ASK  What Row? (principle energy level) (principle energy level)  What section? (type of sub-orbital) (type of sub-orbital)  What seat? (how many electrons in that sub-orbital) (how many electrons in that sub-orbital)

37. Example 1: Write the electron configuration for nitrogen. 7 N 1s 2 2s 2 2p 3 Example 2: Write the electron configuration for Fe. 26 Fe 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6

38. s p d f

39. S orbitals  Starts in 1 st NRG level  1 orbital per energy level  Each can contain up to 2 electrons  Alkali and Alkaline Earth metals  Spherical in nature

40. P orbitals  Starts in 2 nd NRG level  3 orbitals per NRG level  Each orbital can contain up to 2 e- (6 e- total per sublevel)  Groups 3A – 8A  Dumbbell shaped

41. D orbitals  Starts at 3 rd NRG level  5 orbitals per NRG level  2 e- per orbital (10 total e- per sublevel)  Clover shape  Transition Metals

42. F orbitals  Starts in 4 th NRG level  7 orbitals per NRG level  2 e- per orbital (14 e- total per sublevel)  Complicated, unexplainable shape  Lanthanide & Actinide series

43. 2 Important Rules   Pauli Exclusion Principle: An atomic orbital may contain a maximum of two electrons. The electrons must have opposite spins.   Hund’s Rule: When electrons occupy orbitals of equal energy, one electron enters each orbital before they pair.

44. For Example: 2s 2p After the s sublevel gets two electrons, three electrons enter the p orbitals before they pair.

45. What do we know about the table so far?

46. Trends in Atomic Size   radius of an atom cannot be measured directly   atomic radius is estimated as half of the distance between the nuclei of two like atoms in a diatomic molecule.

47. Atomic size generally increases as you move down a group of the periodic table. More principal energy levels are added.

48. Atomic size generally decreases as you move from left to right across a period. Increasing Nuclear charge pulls electrons in closer to nucleus

49. Trends in Ionization Energy   When an atom gains or loses an electron, it becomes an ion.   Ionization Energy is the amount of energy required to remove an electron from a neutral atom.   The smaller the atom, the more ionization energy required.

50. Increases Decreases

51. Trends in Ionic Size   Nonmetallic elements: gain electrons to form negative ions. Negative ions (anions) are always larger than the neutral atoms from which they form.   Metallic elements: lose electrons to form positive ions. Positive ions (cations) are always smaller than the neutral atoms from which they form.

52. Trends in Electronegativity   tendency for an atom to attract electrons when it is chemically combined with another atom.   decreases as you move down a group   increases as you go across a period from left to right.