1. Ch. 6: The Periodic Table

2. Development of the Modern Periodic Table
Why is it called the periodic table?
Properties of the elements repeat in a periodic way
Periodic: predictable

3. History of the Table’s Development
Antoine Lavoisier
- 1700s
- Compiled a list of all the elements known at the time
- 33 elements in 4 categories
John Newlands
1864
Proposed an arrangement where elements were ordered by increasing atomic mass
Law of Octaves: when elements were arranged by increasing atomic mass, the properties repeat every 8th element

4. History of Table’s Development
Dmitri Mendeleev
1864
Created first periodic table
Arranged elements in atomic mass order
Major contribution: left blank spaces and predicted the properties of some undiscovered elements
Stated the 1st periodic law
Periodic Law: when the elements are arranged in increasing atomic number order, there is a periodic repetition of their chemical and physical properties

5. In this early version of Mendeleev’s periodic table, the rows contain elements with similar properties.

6. History of Table’s Development
Lothar Meyer
1800s
Also demonstrated a connection between atomic mass and elemental properties
Arranged elements in order of increasing atomic mass
Not as credited as Mendeleev because Mendeleev’s work was published first
Henry Moseley
1913
Realized arranging the elements by atomic mass was not the best way—some elements ended up in columns with elements of different properties
Established the atomic number
Rearranged the table by increasing atomic number which resulted in a clear periodic pattern

8. The Modern Periodic Table - Organization
Periods (series): horizontal rows
- The elements in each period have the same number of shells
Groups (families): vertical columns
- The elements in each group have the same number of valence electrons
Group numbers 1-18 (new system) or A/B system (older)
Representative Elements
- elements in the s and p blocks
- Groups 1,2 and 13-18
- Possesses a wide variety of chemical and physical properties
Transitional Elements
- elements in groups 3-12

9. Periodic Table Organization
Special Group Names
Alkali Metals: group 1 except for Hydrogen
Very reactive
Alkaline Earth Metals: group 2
Highly reactive
Transitional Metals: group 3-12
Inner Transitional Metals: 2 series below the periodic table
Lanthanide Series: 1st row in f-block
Actinide Series: 2nd row in f-block
Halogens: group 17
Highly reactive, “salt former”
Noble Gases: group 18
Extremely unreactive, stable

10. Inner Transitional Metals
Special Group Names
Inner Transitional Metals

11. Periodic Table Organization
Standard State
Solids- most elements
Liquids- Br and Hg
Gases- upper right corner, H, N, O, F, Cl, and nobles
Synthetic Elements- not found in nature, made in lab
- Tc (#43), Pm (#61) & all elements after uranium (#93 and higher)

12. Standard State

13. Periodic Table Organization
Classes of Elements
Metals
Shiny, solid at room temperature, good conductors, ductile, malleable, high melting points
Nonmetals
Generally gases or brittle, dull-looking solids, poor conductors, low melting points
Metalloids (B, Si, Ge, As, Sb, Te, Po, At) --staircase
Intermediate properties

14. Metalloids – along step line (except Aluminum)

15. A company plan to make an electronic device
A company plan to make an electronic device. They need to use an element that has a chemical behavior similar to that of Silicon (Si) and lead (Pb). The element must have a atomic mass greater than sulfur (S), but less than that of cadmium (Cd). Use the periodic table to determine which element the company can use.

16. Identify each of the following as metal, nonmetal, and metalloid
- Oxygen
- Barium
- Germanium
- Iron

17. If the periodic table were arranged by atomic mass, which of the first 55 elements would be ordered differently than they are in the existing table?

18. Periodic Trends
Many properties of the elements tend to change in a predictable way, known as a trend, as you move across a period or down a group.

19. Atomic Radius
One-half the distance between the nuclei of identical atoms that are bonded together
How closely an atom lies to a neighboring atom
Trends:
generally DECREASES across a period (LR)
generally INCREASES down a group

20. Atomic Radius
WHY the trends occur:
Trend across a period: As electrons are added to the s and p sublevels, they are pulled closer to the nucleus, decreasing the radii
Trend down a group: As electrons occupy sublevels in successively higher energy levels located farther from the nucleus, the sizes of the atoms increase

22. Ionic Radius
Ion: an atom or bonded group of atoms that has a positive or negative charge
When atoms lose electrons and become positive ions, they always become smaller (compared to the neutral atom)
Loss of valence electron can leave an empty outer orbital resulting in a small radius
When atoms gain electrons and become negative ions, they become larger

23. Octet Rule
Octet Rule: atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons
Used to predict what types of ions an element is likely to form
Exception: Hydrogen does not want 8 valence electrons, it will usually give up its one electron to make a positive ion

24. Ionic Radius
Trends:
generally positive ions DECREASE across a period (LR)
generally negative ions (beginning in group 15 or 16) DECREASE across a period (LR)
generally ionic size INCREASES down a group
WHY the trends occur:
Trend across a period: The loss of a valence electron can leave a completely empty outer orbital, which results in a smaller radius. The addition of an electron to an atom increases the repulsion between outer electrons, forcing them apart
Trend down a group: extra inner shell of electrons is added per period causing the radius to increase

25. Ionic Radius

27. Ionization Energy
The energy required to remove one electron from a neutral atom of an element
Energy required to remove the outermost electron from a gaseous atom
Measured in kJ/mol
First Ionization Energy: the energy required to move the first electron
Removing the second electron requires more energy
(second ionization energy)

28. Ionization Energy Trends: generally INCREASES across a period (LR)
generally DECREASES down a group
WHY the trend occurs:
Trend across a period: radii has been decreased, meaning a very strong attraction between the electrons and the nucleus. This means that it will be harder to remove an electron.
Trend down a group: the ionization energy decreases due to a larger radius. Attraction between the nucleus and the outer electrons is also decreased because there are more orbits in between blocking the way.

30. Electronegativity Ability to attract electrons in a chemical bond
Values range from 0.7 to 4.0 Pauling units (arbitrary units)
Most noble gases have no values
Trends:
generally INCREASES across a period (LR)
generally DECREASES down a group

32. Density The amount of mass per unit volume Trends:
generally it INCREASES from left to the middle, and then DECREASES (more rapidly) from the middle to the right across a period
generally INCREASES down a group
Why the trend occurs:
Increases as you go down because the atomic radius increases, volume increases, as well as mass at a larger rate, so it will be more dense.

33. Trends in Density Generally Increases Generally Increases