1. Chapter 6 The Periodic Table and Periodic Law

2. 6.1 Development
Antoine Lavoisier (late 1700s) organized the 33 known elements (at the time) into a list with four categories
John Newlands (1864) noticed that when elements were arranged by increasing atomic mass, properties repeated every 8 elements
periodic pattern – the law of octaves

3. Meyer, Mendeleev, and Moseley
1869-Meyer and Mendeleev both saw a connection between atomic mass and properties
Mendeleev got more credit because he was first to publish it
Left holes of undiscovered elements
Predicted properties of some undiscovered element- Sc, Ga, Ge
Still not completely correct
Mosley arranged his table by atomic number based on protons instead mass
fixed the problems with Mendeleev’s table

4. The Periodic Law
The statement that there is a periodic repetition of chemical and physical properties of the element when they are arranged by increasing atomic number.
horizontal rows are called periods
vertical columns are called groups or families

5. The Modern Periodic Table
Groups 1A through 8A = Representative Elements (a.k.a Main) because they contain a wide range of chemical and physical properties
Groups 1B through 8B = Transition Elements

6. Classifying the elements
Metals
shiny, smooth, clean, solid room temperature, good conductors of heat and electricity
Alkali Metals= 1A (excluding hydrogen)
highly reactive
Alkaline Earth Metals= 2A
highly reactive (not as much as 1A)‏

7. Classifying the Elements Continued…
Transition metals
Group B elements contained in the D block of the table
Inner transition metals
the lanthanide and actinide series
below the table

8. Classifying the elements continued
Nonmetals=
Generally a gas or a brittle, dull-looking solids
Poor conductors
Halogens=7A
REALLY REACTIVE
Noble Gases= 8A
unreactive and stable (all valence electrons are filled)‏

9. Classifying the Elements Continued…
Metalloids
contains the physical and chemical properties of both metals and nonmetal

10. 6.2 Classification of the Elements
The properties of each element in each group are similar because they have the same number of valence electrons
The number of the group that the element is in = the number of valence electron (excluding the transition metals)‏
The energy level of an element’s valence electrons indicates the period on the periodic table

11. S,P,D, and F Blocks 4 different energy levels: s, p, d, and f
S block= 1a and 2a
holds max of 2 electrons
P block= 3A through 8A
max holds 6 electrons
S block must fill before P block can fill
Noble gases are stable because of filled S and P blocks

12. S,P,D, and F Blocks Continued…
D block = transition metals
max of 10 electrons
F block= inner transition metals
unpredictable manner of filling
max of 14 electrons

13. 6.3 Periodic Trends Electron clouds are fuzzy.
So what does atomic size or radius mean?
Ideas?

14. Atomic Radius How it's normally done:
Take a crystal of the pure element
Find distance between adjacent nucleii

15. Atomic radius

16. Periodic Trends: Atomic Radius
DECREASES I N C R E A S

17. Atomic Radius Why do atomic radii increase as you move down a group?
Why does it decrease as you move across a period?

18. Atomic Radius
The increase from top to bottom is due to adding electron shells.
The decrease from left to right is due to increased nuclear charge as you move to the right, which draws electrons closer to the nucleus.

20. Ions Atoms can gain or lose one or more electrons to form an ion.
An ion is an atom or a bonded group of atoms with a positive or negative charge.

21. Ionic Radius an ion’s radius will be affected by its formation
Would a cation (positive charge) be smaller or larger than the corresponding neutral atom?
Would an anion (negative charge) be smaller or larger than the corresponding neutral atom?
Explain ionic radii trends within periods and groups.

22. Ionization Energy Remember excited states of electrons?
As energy level increases, distance from nucleus increases
Add enough energy, the electron is no longer bound to the nucleus.
Now, the atom has more protons than electrons.
What's its charge?

23. Ionization Energy

24. Periodic Trends: Ionization Energy
Ionization energy= energy required to remove an electron from a gaseous atom
Octet rule = atoms tend to gain lose or share electrons to acquire a full set of 8 valence electrons

25. Periodic Trends: Ionization Energy
INCREASES D E C R A S

26. Ionization Energy
Why does ionization energy tend to increase across a period?
Why does it tend to decrease down a group?

27. Periodic Trends: Electronegativity
Electronegativity= relative ability of an atom to attract electrons in a chemical bond.

28. Electronegativity
Arbitrary units called Paulings (after Linus Pauling) are used to express electronegativity.
Electronegativity increases from left to right across a period and from top to bottom down a group.
Why would these trends occur?

29. Periodic Trends: Electronegativity
INCREASES D E C R A S