1. The Periodic Table

2. Guiding Questions Why is the periodic table so important?
Why is the periodic table shaped the way it's shaped?
Why do elements combine? Why do elements react?
What other patterns are there in the world and how do
they help us?
Periodic Table Study Questions
1.         Why did chemists make the periodic table?
2.         Why was the table difficult to make?
3.         Why were Dobereiner’s triads of limited use at a periodic table?
4.         What did Newland discover about the elements?
5.         What did Meyer contribute to the development of the periodic table?
6.         What did Mendeleev use as the organizing property for the periodic table?
7.         What problem developed from the use of this property?
8.         What is common to elements in a column of the table?
9.         How did properties change in a row of the table?
10.       What was the significance of gaps in Mendeleev’s periodic table?
11.       What did Moseley use to order the elements in the periodic table?
12.       How did Moseley change the periodic law?
13.       What determines the identity of an element?
14.       Why do elements in a column of the periodic table have similar properties?
15.       With respect to the Periodic Table, what is the meaning of periodic?
16.       What does a row of the Periodic table represent?
17.       What happens to valence electrons as you move left to right in a row?
18.       When determines stability in an atom?
19.       List, from least to most, the stable configurations in an atom.
20.       What determines the column of the periodic table an element is in?
21.       What sublevels are in the outer level of an atom?
22.       What is the maximum number of electrons in the outer level of an atom?
23.       What determines the row and column of the periodic table an element is in?
24.       What are common properties of metals?
25.       What are common properties of non-metals?
26.       What three things can happen to electrons when atoms form compounds?
27.       The configuration of He is 1s2, but it is placed in column 18. Explain this discrepancy.
28.       Hydrogen is obviously not an alkali metal. Why is it in column 1 of the table?
29.       What is necessary for a metalloid to act as a semiconductor?

3. Mendeleev
“Father of Periodic Table”
organized elements based on increasing atomic mass. Found similarities in chemical properties and published his first periodic table in 1869.

4. Moseley
discovered while working with elements that they fit better into pattern when arranged by nuclear charge (number of protons-also known as atomic number).

5. Periodic Law states that physical and chemical properties of the elements are periodic functions of their atomic numbers.
This means that when elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals

6. Organization of the Periodic Table
Periods – rows across, horizontal
Indicates number of energy levels (Principle Quantum Number)
Groups/Families – down, columns 18
Indicate number of e- in outer most energy level (1-2, main group elements)
Share chemical properities

7. Three Categories of Elements
Metals
Non-Metals
Metalloids

8. Three Categories of Elements
Metals
- Groups 1-12 (except H) and under stair-step groups 13-15
- Form ionic and metallic bonds
Luster (shiny, reflects light)
Malleable - can be flattened into sheets
Ductile - can be drawn into thin wires
Good Conductors - heat and electricity can flow throughbecause the outer electrons are not held tightly to the nucleus and move freely
- Most are solid at room temperature except Mercury

9. The Periodic Table H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg
The Metals are represented in the Periodic Table in blue.

10. Three Categories of Elements
Non-Metals
Dull
Not Malleable/Ductile
Poor Conductors
- P block
- Form ionic and covalent bonds
- Most do not conduct heat or electricity
-All, except H, are found on right of periodic table
- At room temperature, most are solid or gaseous

11. The Periodic Table H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg
The Non-Metals are represented in the Periodic Table in yellow.

12. Three Categories of Elements
Non-Metals
Some Examples of Non-Metals
Oxygen (O)
Helium
(He)
Sulfur (S)
Chlorine
(Cl)
Neon (Ne)
Nitrogen
(N)

13. Three Categories of Elements
Metalloids
Have properties of both metals and non-metals.
Generally not shiny
Not malleable and ductile
- Form ionic and covalent bonds
They conduct heat and electricity better than nonmetals, but less than metals.

14. Three Categories of Elements
Metalloids
Here are the Metalloids
Boron (B)
Arsenic (As)
Tellurium (Te)
Silicon (Si)
- Polonium (Po)
Antimony (Sb)
Germanium (Ge)

15. The Periodic Table H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg
The Metalloids are represented in the Periodic Table in green.

16. The Periodic Table H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg
What do you notice about the way the element groups are arranged in the Periodic Table?

17. Hydrogen In a class by itself
Most common element in the universe (3/4)
Behaves unlike any other element because it consists of one proton and one electron.

18. Main Group Elements S and p blocks Groups 1,2, and 13-18
Four groups have special names:
Alkali metals (Group 1)
Alkaline-earth metals (Group 2)
Halogens (Group 17)
Noble gases (Group 18)

19. Group 1 – Alkali Metals Characteristics Soft, silver metals
Low melting points and densities
Highly reactive (esp. w/ water)
Do not occur in nature in elemental form
Stored in kerosene or mineral oil
Have one Valence electron

21. Group 2 – Alkaline-earth Metals
Characteristics
Gray, metallic solids
Reactive; but, less than alkali
Not found as free elements in nature
Bright fireworks, aircraft
Harder, denser, and stronger than alkali metals
Higher melting points than alkali
Have 2 valence electrons

23. Transition Elements Groups 3-12 (elements in transition) D block
Vary in reactivity and can be found as free elements
Metal characteristics
Form colored compounds
Often occur in nature as uncombined elements
Hg – mercury – liquid metal

24. Valence Electrons for Transition Elements
Group 3: 3 valence electrons
Group 4: 2 to 4 valence electrons
Group 5: 2 to 5 valence electrons
Group 6: 2 to 6 valence electrons
Group 7: 2 to 7 valence electrons
Group 8: 2 or 3 valence electrons
Group 9: 2 or 3 valence electrons
Group 10: 2 or 3 valence electrons
Group 11: 1 or 2 valence electrons
Group 12: 2 valence electrons

26. Group 13 – Boron Group or Icosagens
Have 3 valence electrons
Boron is the only metalloid in the family
The rest are poor metals.

28. Group 14 – Carbon group or Crystallogens
Have 4 valence electrons
Unique feature is that the elements can form different anions and cations
Ex: C: 4-
Si and Ge: 4+
Sn and Pb: 2+

30. Group 15 – Nitrogen Group or Pnictogens
Have 5 valence electrons
Able to form double and triple bonds

32. Group 16 – Oxygen group or Chalcogens
Have 6 valence electrons
From -2 ions
Physical properities vary dramatically
Ex: oxygen – colorless gas
Sulfur – yellow solid
Tellurium – silver metalloid
Selenium - black

34. Group 17 - Halogens Most reactive non-metal Irritating odor
Interact with alkali metals to form salts
7 electrons in outer energy level
Easily pickup one electron
Bromine – only liquid nonmetal
Flourine and chlorine – gases at room temp
Iodine and astatine – solids at room temp

36. Group 18 – Noble Gases No color or odor
Exist as individual gas atoms (monatomic)
Full outer energy level
Relatively un-reactive (inert gases)

38. Inner Transitional Metals
f-block, n-2
ALL are radioactive and unstable
Includes lanthanides (atomic number 58-71)
Shiny metals, similar in reactivity to alkaline earth metals
and actinides (atomic number )
First 4 are found on earth, the remaining are synthetic

42. Diatomic Molecules
Elements That Exist as Diatomic Molecules in Their Elemental Forms
Element Present Elemental State at 25 oC Molecule
hydrogen colorless gas H2
nitrogen colorless gas N2
oxygen pale blue gas O2
fluorine pale yellow gas F2
chlorine pale green gas Cl2
bromine reddish-brown liquid Br2
iodine lustrous, dark purple solid I2

43. How do we use an elements location on the periodic table to determine its ionic charge?
Atomic Radius (Angstrom)
½ the distance from the nuclei to another
From the nucleus to edge of e- cloud
Going down a group, Atomic radius increases because of the increasing number of energy levels
Going across a period (left to right), atomic radius decreases because of the increase in positive charge in the nucleus

44. Atomic Radii

45. Coulombic attraction Attraction of + and – charges
Two factors determine the strength:
Amount of charge
Distance between charges

46. Shielding Effect
Kernel electrons “shield” valence electrons from attractive force of the nucleus
Caused by kernel and valence electrons repelling each other
The more electron shells there are, the greater the shielding effect.
Explains why valence electrons are more easily removed

47. + Shielding Effect - - - - nucleus Electron
Valence + -
nucleus - -
Electrons -
Some schools use the term “screening” as I use the term “shielding”.
Electron
Shield
“kernel”
electrons
Kernel electrons block the attractive force of the nucleus from the valence electrons

48. Why is cesium bigger than sodium?
Sodium has the electron configuration 1s22s22p63s1. There is one valence electron.
The attraction between this lone valence electron and the nucleus with 11 protons is shielded by the other 10 core electrons.
The electron configuration for cesium is 1s22s22p63s23p64s23d104p65s24d105p66s1.
While there are more protons in a cesium atom, there are also more electrons shielding the outer electron from the nucleus. The outermost electron, 6s1, therefore, is held very loosely. Because of shielding, the nucleus has less control over this 6s1 electron than it does over a 3s1 electron.

49. Ionization Energy
Ionization is a process that results in the formation of an ion
An ion is an atom or group of atoms that have a “+” or “-” charge
Change is created by gain or loss of e-
Losing an e creates a “+” charge (cation)

50. (also known as 1st ionization energy)
Losing the e- requires energy. Energy required to remove one e- from a neutral atom is called
IONIZATION ENERGY
(also known as 1st ionization energy)

51. Factors that affect ionization
Nuclear charge – the larger the nuclear charge , the greater IE
Shielding effect – the greater the shielding effect, the less IE
Radius – greater the distance between the nucleus and the outer electrons of an atom, the less IE.
Sublevel – an electron from a sublevel that is more than half-full requires additional energy to be removed

52. Ionization energy decreases as you go down a group because valence electrons are farther from the nucleus
And increases as you go across a period because of the greater positive charge leads to greater attraction to electron

53. Ionization Energy

54. ELECTRON AFFINITY Electron Affinity
Gaining an electron results in an ion with a “-” charge (anion)
When an atom gains an e- it causes an energy change. The energy change when a neutral atom gains an e- is the
ELECTRON AFFINITY

55. Electron Affinity Electron affinity decreases as you move down a group
Increases as you move across a period
Halogens have the highest electron affinities

56. Electron Affinity

57. Electronegativity (Pauling)
Ability of an atom to attract (or remove) an e- from another atom
Fluorine is most electronegative (4)
Metals have Electronegativity of less than 2

58. Electronegativity
Electronegativity decreases as you go down a group because the electrons are farther from the nucleus
Electronegativity increases as you go to the right because atoms are more inclined to gain electrons in order to gain a full shell

59. Electronegativity

61. Summary of Periodic Trends
Shielding is constant
Atomic radius decreases
Ionization energy increases
Electronegativity increases
Nuclear charge increases 1A
Ionization energy decreases
Electronegativity decreases
Nuclear charge increases
Atomic radius increases
Shielding increases
Ionic size increases 2A 3A 4A 5A 6A 7A
• Elements with the highest ionization energies are those with the most negative electron affinities, which are located in the upper-right corner of the periodic table.
• Elements with the lowest ionization energies are those with the least negative electron affinities and are located in the lower-left corner of the periodic table.
• The tendency of an element to gain or lose electrons is important in determining its chemistry.
• Various methods have been developed to describe this tendency quantitatively.
• The most important method is called electronegativity (), defined as the relative ability of an atom to attract electrons to itself in a chemical compound.
Rules for assigning oxidation states are based on the relative electronegativities of the elements — the more-electronegative element in a binary compound is assigned a negative oxidation state
Electronegativity values used to predict bond energies, bond polarities, and the kinds of reactions that compounds undergo
Trends in periodic properties:
1. Atomic radii decrease from lower left to upper right in the periodic table.
2. Ionization energies become more positive, electron affinities become more negative, and electronegativities increase from the lower left to the upper right.
Ionic size (cations) Ionic size (anions)
decreases decreases