1. Dmitri Mendeleev( )
He develop the 1st
periodic table of the
elements.
Arranged elements in order of increasing
atomic mass and
created columns
with elements having similar properties.

2. Mendeleev’s
Table
Drawbacks
Tellurium and
Iodine
Potassium and Argon
Cobalt and Nickel

3. Henry Moseley (1887 – 1915)
Arranged elements in order of increasing atomic number thus reversing the order of the elements and correcting the drawbacks found in Mendeleev’s table.

4. Glenn Seaborg (1912 – 1999)
Rearranged the periodic table to pull out the lanthanide and actinide series during the Manhattan Project.

5. Periodic Law
Periodic law states the properties of the elements are periodic functions of their atomic number. In other words, when the elements are listed in order of atomic number, elements with similar properties appear periodically.
Therefore, elements in the same column have similar properties.
Periodic – to appear at regular intervals

6. Modern Periodic Table
Period – Row on the periodic table. Periods reflect the energy level of the electrons. .

7. A Group or Family is a column on the periodic table
A Group or Family is a column on the periodic table. Elements in the same column have similar chemical properties.

8. Metals are elements located to the left of the jagged stairs except hydrogen.

9. Metals are solids except mercury, which is a liquid.
Properties of Metals
Metals are solids except mercury, which is a liquid.
Metals have luster, are malleable, ductile, and have high tensile strength.
Metals are good conductors of
heat and electricity.

10. Nonmetals
Nonmetals are elements located to the right of the jagged stairs plus hydrogen.

11. Properties of Nonmetals
Nonmetals are solids or gases, except bromine, Br, which is a liquid. Nonmetals are dull, and lack other metallic properties. Nonmetals are generally poor conductors of heat and electricity.

12. Metalloids
Metalloids are elements bordering the stairs except aluminum. They have properties of metals and nonmetals.

13. Metalloids are generally semiconductors which means
that they conduct to varying degrees making them useful in
the computer industry.

14. Group A Elements
Group A elements (‘s’ and ‘p’ block)all have electrons in the outer s, or s and p orbitals. The Roman Numeral group number indicates the number of valence electrons except with helium which has 2 valence e-
Examples:
IIA - Ca (20) 1s22s22p63s23p64s2
VIA – S (16) 1s22s22p63s23p4

15. Group 1 (IA) Elements
Group 1(IA) elements are the alkali metals with one valence electron. Alkali metals are soft, silver in color, and are too reactive to be found in nature in their free form. Hydrogen is NOT an alkali metal.

16. Group 2 (IIA)
Group 2(IIA) elements are the alkaline earth metals with 2 valence electrons. Alkaline earth metals are harder, denser, and stronger than alkali metals. They have higher melting points and are less reactive than alkali metals but are also too reactive to be found in their free state in nature.

17. Group 17 (VIIA)
Group 17 (VIIA) elements are the halogens with 7 valence electrons. They are the most reactive nonmetal family (2nd most reactive family on the table), and not usually found pure in nature. They reactive vigorously with metals, to steal an electron and form negative ions.

18. Group 18 (VIIIA)
Group 18 (VIIIA) elements are the noble gases with 8 valence electrons, except helium which has 2. Noble gases are inert (nonreactive) in nature. They do not form ions.

19. Group B Elements
Group B elements or transition elements (d block) have electrons in their outer d orbitals. The have varying number of valence electrons but frequently have 2 (exceptions, Zinc is always +2 and Silver is always +1 as ions).

20. Transition elements form very colorful ions in solution and light emissions.

21. Lanthanoid and Actinoid Series or Inner Transition Elements
Lanthanoid and Actinoid elements
(f-block) have electrons in their outer
f orbitals. These elements have varying numbers of valence electrons. Inner transitions metals generally form +3 ions.
Example: Nd (60)
1s22s22p63s23p64s23d104p65s24d105p66s24f4

23. Stability of Electron Configurations
Octet Rule – Atoms having all their outer s and p orbitals filled are more stable (less reactive) than partially filled orbitals. Therefore, atoms will gain or lose electrons in order to achieve a stable configuration.
Stable configurations noble gases which have 8 valence electrons.

24. Example 1
Li (3)- Lithium will lose one valence electron to become 2 and resemble Helium.
O (8)- Oxygen will gain two valence electrons to become 10 and resemble Neon

25. Exceptions to Predicted Electron Configurations
According to the octet rule, filled and half-filled sublevels are more stable (less reactive). Therefore, in some cases, actual configuration varies from predicted configurations

26. Exceptions of the Octet Rule
Predicted Configuration
Chromium Cr 1s22s22p63s23p64s23d4
Actual Configuration
1s22s22p63s23p64s13d5

27. Exceptions to the Octet Rule
Predicted Configuration
Copper Cu 1s22s22p63s23p64s23d9
Actual Configuration
1s22s22p63s23p64s13d10

28. Ions
Ions- an atom that has gained or lost electrons.
Cations is an atom that has lost electrons and therefore has a positive charge.
Metals lose electrons to form positive ions, cations.
Metals lose all their valence electrons. Therefore, their ions are positive by the number they lose.

29. Anions
Anion- is an atom that has gained electrons and therefore, has a negative charge. Anion named end in –ide. S-2 is called the sulfide ion.
Nonmetals gain electrons to from negative ion, anions. They gain electrons to have a stable octet (8) of electrons. Therefore, nonmetal ions are negative by the number of electrons they gain.

31. Ion Formulas
Ion formulas consist of the element’s symbol followed by its charge or oxidation state.
Rubidium - Rb Iron – Fe+2
Aluminum – Al Lead – Pb+4
Sulfur – S Iodine – I-1
Nitrogen – N-3
Hydrogen – H+1 or H-1

32. Periodic Properties
As you have seen, elements in the
same column are similar in their outer
electron configurations.
This results in these elements having relatively the same physical properties such as density, melting point, and boiling.
These physical properties are then said to be periodic properties.

33. Periodic Trends
A periodic trend is a general tendency that occurs across periods or down groups on the periodic table. Exceptions are always present in trends.

34. Atomic radii is the radius of an atom.
Down a Group – radius increases.
Reasons – * Addition of energy levels
* Shielding of outer electrons
from the nucleus by inner
electrons in larger atoms.
* Electron – electron
repulsion in outer energy
levels

35. Across a Period – Radius Decreases Reasons – 1. no addition of energy
levels
2. increased nuclear
charge causes
electrons to be pulled
closer

37. Isoelectric Particles
Isoelectric particles- are particles containing the same number of electrons.
This would include a noble gas and all of the ions that resemble it exactly in electron configuration.
Across a period, electron affinity increases

38. Ionic Radii- is the radius of an ion.
Positive Ions (cations)– are smaller than their neutral atom.
Reasons – 1. loss of an energy
level
2. The nucleus is attracting fewer electrons.

39. Ionic Radii- is the radius of an ion.
Negative Ions (anions)– are larger than their neutral atom.
Reasons –
1. The nucleus is attracting more electrons.
2. Those attracted electrons continue to repel each other (e-/e- repulsions)

40. Electric Force, Fields, Current
04/18/2006
Coulomb’s Law
Like charges repel
Unlike charges attract + d – + –
F = Force of Attraction
k = coulomb constant
q1= charge of 1st particle
q2= charge of 2nd particle
d= distance between the two particles midpoints
F = k q1 q2 d2
Lecture 6

41. Coulomb Force Law, Qualitatively
Electric Force, Fields, Current
04/18/2006
Coulomb Force Law, Qualitatively
Double one of the charges
force doubles
** Charge and Force are directly proportional
Double the distance between charges
force four times weaker (otherwards ¼ the force of attraction)
Distance and Force are exponentially Inversely Proportional
Double both charges
force four times stronger
Lecture 6

42. First Ionization Energy
1st Ionization Energy is the energy required to remove an electron from an atom.
Down a Group – Ionization Energy
Decreases
Reason – Outer electrons in larger
atoms are held more loosely
by the nucleus.

43. 1st Ionization Energy
Across a Period – Ionization Energy
Increases
Reasons – 1. Outer electrons in
smaller atoms are
held more tightly by
the nucleus.
2. An octet of electrons
is approached.

44. Periodic Trend for 1st Ionization Energy

45. Electronegativity is the relative attraction of an atom for electrons when forming chemical bonds.

46. Down a Group
Electronegativity decreases.
Reasons
Larger atoms- the nucleus is farther away from the outer energy level and cannot attract more electrons
Larger Atoms – the nucleus is shielded from the outer energy level and cannot attract more electrons

47. Electronegativity increases across a period.
Reasons:
Smaller atoms – the nucleus is closer
to the outer energy level and can attract more electrons
Smaller atoms – the nucleus is not shielded from the outer energy level and attracts more electrons
An octet of electrons is approached.

48. Periodic Trend for Electronegativity

49. More electron affinity
The attraction of an atom for additional electrons.
Because they already have a stable number of electrons, NOBLE CASES HAVE NO ELECTRON AFFINITY.
As you go down a group, electron affinity decreases.
This is because the nucleus is less attracted to additional electrons in larger atoms because:
Additional energy levels
Shielding of outer electrons from the nucleus by inner electrons in larger atoms
Electron-electron repulsion in outer energy levels