1. CH 6: The Periodic Table

2. 6.1: Organizing The Elements
By 1700: Only about 13 elements identified and known.
Chemists knew of and suspected other elements, but couldn’t separate them.
As Chemistry became more advanced, more elements were discovered at alarming rates!
Chemists decided they had to organize these elements in some way

3. Early Periodic Tables
Early Chemists used the Properties of elements to sort them into groups.
Ex: Bromine, Iodine, and Chlorine were grouped by their similar properties.

4. Early Periodic Tables
In 1869, a Russian Chemist by the name of Dmitri Mendeleev made his own Periodic Table of Elements.
Mendeleev organized his elements in groups of properties, and by Increasing Atomic Mass.
With this table, he was even able to predict properties of elements not discovered yet (Like Germanium!)

5. Modern Periodic Table
As we know, Atomic Mass can vary with different isotope abundances, breaking Mendeleev’s rule (he even broke this when atomic mass didn’t match with properties)
So, the Modern Periodic Table uses a different property that is more reliable.
The Modern Periodic Table is organized based on Increasing Atomic Number!

6. Modern Periodic Table Rows of the table: Called Periods
Columns of the table: Called Groups/Families
Each Period of the table corresponds to a Principal Energy level!
Ex: Row 1: 2 e-s, 2 elements, Row 2: 8 e-s, 8 elements… etc.
Same Group = Same outer e- configuration! (aka valence e-’s)

7. Periodic Properties
Properties of the elements change as I move leftRight across a Period, and repeat with every new Period of the table.
Called Periodic Law
Ends up putting the elements with similar Physical properties in the same Group.

8. Modern Periodic Table
Many ways to categorize elements on the Periodic Table
One way is in three Categories:
Metals
Nonmetals
Metalloids
These are organized by their properties AND Position on Per. Table

9. Metals The first category we will talk about are Metals.
Most elements of the periodic table are Metals
Some Examples of metals include: Iron (Fe), Cobalt (Co), Lithium (Li), and Zirconium (Zr). Any others? (There’s plenty!)

10. Properties of Metals
THINK PAIR SHARE: What are some properties of Metals that you know of?
Some Properties of Metals:
Luster (shiny), Malleable (bendy), Good Conductors, Low Electron Affinity, High Melting Points (usually solid), Hard.
Majority of elements on table are Metals!
Metals are all to the LEFT of the “Staircase” on the Per. Table!

11. Nonmetals
Nonmetals are generally considered the Opposite of the metals.
There are much less Nonmetals than Metals
The Nonmetals are located to the RIGHT of the “Stairstep”
Examples of Nonmetals include: Sulfur (S), Chlorine (Cl2), Phosphorous (P).

12. Properties of Nonmetals
If Nonmetals are the opposite of metals, can you guess some of their properties too?
Dull (opposite of shiny), Brittle (not malleable), High Electron Affinity, Lower Melting Points, All states at room temp, Bad Conductors (Insulators)

13. Metalloids
Metals are on the Left, Nonmetals are on the right, Metalloids are ON the “Stairstep”!
THINK PAIR SHARE: If Metalloids are between Metals and Nonmetals, what do you think their properties are like?
Examples of Metalloids: Arsenic (As), Germanium (Ge), Silicon (Si)
Aluminum is on the steps, but a metal, because of metallic props.

14. Periodic Table as a tool
The periodic table has lots of useful information!
See the picture below for a diagram of info you can get from a detailed Periodic Table
# on top: Atomic #
Ti: Element symbol
Element name (duh)
Bottom: Atomic Mass
Right: e- NRG level filling

15. Periodic Table as a Tool
Different colors/outlines= Different states at room temp!
Groups/Families: Similar Properties, due to outer e- configuration.
Some have own nicknames!
Others do not

16. Periodic Table ”A” Groups: Also known as Representative Elements
Any element that ISNT the transition metals
Transition Metals: in the d-block of the table
Inner Transition Metals: in F-block, actinide and lathanides.

17. Groups WHAT YOU NEED TO KNOW:
What elements are in what group (ex: what group is Sodium in?)
Find an element based on Group and Period (ex: element in period 3, group 7A is?) Chlorine
Nicknames for groups, and any outstanding/special properties
Outer e- configuration for each group

18. Alkali Metals These metals are on group 1A of the Table.
Highly Reactive, especially in water.
Examples: Lithium(Li), Sodium(Na), Potassium(K), Rubidium(Rb), Cesium(Cs).
E- conf: s1
Here’s a video showing what can happen with these metals in water:
Do we want to do our own???

19. Alkaline Earth Metals Group 2A on the Periodic Table
Also reactive with water, more reactive with Chalcogen group like Oxygen (we will get to these)
Examples: Beryllium(Be), Magnesium(Mg), Calcium(Ca), Strontium(Sr), Barium (Ba).
E- conf: s2
Heres a video on what happens when you burn magnesium:

20. Group 3A No real ”Nickname” like Alkali Metals
Examples: Boron (B), Aluminum(Al), Gallium (Ga)
E- Conf: s2p1

21. Group 4A Group 4A on the Table, still no Nicknames
Examples: Carbon(c), Silicon(Si)
Special property of Carbon: Makes 4 bonds, really good for making big molecules
This is why much life is Carbon Based, but scientists think aliens out there could be Silicon-Based!
E- conf: s2p2

22. Group 5A No Nickname, again… just 5A Mildly Reactive.
Examples: Nitrogen(N), Phosphorus(P) E- conf: s2p3

23. Chalcogens Group 6A on the Periodic Table.
These elements are usually more reactive than 5A.
Examples: Oxygen(O), Sulfur (S), Selenium (Se).
Interesting Question: Is Oxygen Killing us? E-conf: s2p4

24. Halogens Group 7A of the Table!
Very Highly Reactive! These elements don’t mess around!
Examples: Fluorine(F),Chlorine(Cl),Bromine(Br),Iodine(I)
E-conf: s2p5
Here’s what I mean: Pointing Fluorine at stuff:
2:40-5:15

25. Noble Gases Group 8A of the Periodic Table
These gases are not reactive at all, hence the “Noble” gases.
Useful for keeping things from reacting!
E-conf: s2p6
Examples: helium (He), Neon(Ne), Argon(Ar), Krypton(Kr), Xenon(Xe)

26. Ions
These groups of our table, when reacted, can all form charged atoms
Recall: Charge practice earlier this section!
A Charged atom is also referred to as an Ion
Positive Ions: Cations
Negative Ions: Anions
Missing 1 electron: 1+
2 EXTRA electrons: 2-
Groups usually make specific ions (Group 1= 1+)
Group2=2+…... Group 5= Group 7=1-

27. Ions Cations Attract anions in what we call ionic bonds
Ionic Bonds make special compounds: called Ionic Compounds
Happens when a metal readily gives up an electron to a nonmetal
What happens when you lose a proton/gain an electron?
Result: Opposite charged atoms stick together. Often form crystalline structures

28. Metal Reactivity
Think Pair Share: Why do you think some metals corrode (rust) faster than others?
Answer: Not all Metals are created Equal!
As you move left/down on the table: Metallic properties Increase
As you move up/right: Metallic Properties decrease
What this means: Cesium is more metallic than Fluorine, and will corrode much faster than less metallic elements.(amongst other metallic properties).
This is called a Periodic Trend

29. Trends
As we move Across the Periodic Table, and Down the Periodic Table different changes occur in atoms.
Changes in the following:
Electronegativity
Atomic Radius
Ionic Radius
Ionization Energy

30. Atomic Radius
Atomic Radius is also known as how Large a neutral atom is
Left to Right across a Period: Atomic Radius generally Decreases
This is due to atoms holding their electrons less tightly
Top to Bottom down a group/family: Atomic Radius Generally Increases
This is due to atoms with many more protons and neutrons in their nucleus, and electrons to orbit.

31. Practice:
Which of the following elements have the largest atomic radii?
Cl or Br: Br, further down a group
Na or S: Na, further left on a period
O or P: P, both further down AND left
N or S: inconclusive, further down, but to the right, won’t give you these.

32. Ionic Radius When atoms make Ions, their Radius changes.
Cations get smaller (lose e-s, they hold the rest tighter)
Anions get bigger (extra e-s cant be held as tightly)
The more “extreme”/charged the ions, the larger/smaller they become
In General:
As you move down a group, Ionic Radius INCREASES

33. Practice Which of the following has a larger Ionic Radius?
Zn or Zn+: Zn: Zn+ shrinks
Na+ or K+: K+, lower on a group
P or P3-: P3-, Anion becomes larger
Al3+ or P3-: P3-, anion is larger and cation is smaller

34. Ionization Energy
Ionization Energy: The amount of energy required to remove an electron from that atom.
IN GENERAL: FIRST IONIZATION ENERGY
INCREASES Left==> Right across a Period
DECREASES Top to bottom of a Group

35. Ionization Energy Cont.
If I can remove one electron, who’s to say I can’t remove two? Or three?
As I remove electrons, Ie (Ionization Energy) goes UP steadily.
RECALL: e-s on outermost energy levels from CH5, called Valence electrons.
These are relatively easy, or need little energy to remove.
Once I try to take an e- from a full NRG level, or Core Electron, Ie goes WAAAAAAY up!
That “Jump” can tell you valence e-s, and therefore? The GROUP the element is in!
Can look these up!

36. Practice Which has higher 1st Ie?
Cs or Na? Na, it’s higher on a group, holds e-s closer
F or O? F, further right, holds e-s tighter
Ie’s in order (kJ/mol): 300, 400, 500, 1900, 2250,3900… which group?
Group 3A!
Ie’s (kJ/mol): 250, 300, 450, 600, 750, 900, 22500, 30000
Group 6A!

37. Electronegativity
Electronegativity is the ability for an atom to attract electrons
We will use this later as we go over Bonding in later chapters!
IN GENERAL:
Left to right across a period: E-negativity INCREASES
Top to bottom of a group: E-negativity DECREASES
Fluorine is most E-neg! (4.0, top of scale)
Cesium is least E-neg! (0.7, bottom of scale)

38. That’s all!