1. The Periodic Table
Chapter 4

2. Unit Essential Question:
What information can be determined from the periodic table?

3. Lesson Essential Question:
How have elements been organized into the periodic table used today?

4. Section 1: How Are Elements Organized?
1865- when placed in order of atomic mass, scientists noticed that properties repeated every 8 elements.
Dmitri Mendeleev used this discovery and placed all 63 known elements into the first periodic table.
When chemical properties repeated he started another row.
Each column had elements with similar properties.
Able to predict missing elements using this repetition.
Problem with ordering elements by atomic mass:
Some did not match properties of other elements in the same column- needed to be switched around.
Instead of placing elements with similar properties in the same row, Mendeleev placed them into columns. He was the first to predict the existence of undiscovered elements and also to notice the issue with using atomic mass to order the elements. Thought that maybe the atomic masses were wrong, however more research indicated they were correct- unable to explain why some elements needed to be switched out of order of atomic mass to fit in the appropriate column. Made big contributions to the periodic table- ‘father of the periodic table’.

5. Here is a written version of the table Mendeleev worked on
Here is a written version of the table Mendeleev worked on. Notice the blank spaces as well as the ? marks.

6. Take a look at Mendeleev’s predictions
Take a look at Mendeleev’s predictions. He was extremely close to most of them without having any samples to test. These were all based on surrounding elements.

7. Periodic: happening or recurring at certain intervals.
About 40 years after Mendeleev’s table, Henry Moseley made an important change to the periodic table.
Organized elements by atomic number instead of atomic mass.
Elements that had not previously fit into the correct column when ordered by atomic mass were fixed.
Periodic: happening or recurring at certain intervals.
Ex: waves are periodic:
Periodic law: properties of elements are repeated at certain intervals when ordered by atomic number.
Moseley’s work allowed those “reversed pairs” to fit in with the pattern of properties that Mendeleev had seen.

8. More on the Periodic Law
Properties of elements change as electron configurations change.
Since the # e- = atomic #, you can also think of the periodic table being ordered by electron configurations.
Can easily determine the # of valence electrons , which are the outermost e- (comes from the column #).
Changing e- changes the elements’ properties/behavior!
Group = column (also known as families).
Recall that these elements all have similar properties!
This is because they have the same # of valence electrons!
Period = row (energy level).
Recall that each row begins when properties begin repeating again.
Since the periodic table is ordered by atomic #, it is also ordered by electron configurations (since atomic # in a neutral atom = # of electrons). Thus, we can see how many valence (outer) electrons exist in an element by what column (group) the element is in- each column/group has the same properties because they have the same valence electron configuration.

9. Lesson Essential Question:
Section 2: Tour of the Periodic Table
Lesson Essential Question:
How are elements grouped on the periodic table?

10. Yellow = nonmetals Blue = metalloids Green = metals
Here you can see a visual of how many elements are metals vs. nonmetals. The green elements are metals, the yellow are nonmetals. The blue elements are metalloids or semi-metals – they have properties of both metals and nonmetals. Notice Al and Po are not listed as metalloids. Most periodic tables show a “stair step” border between the metals and nonmetals.

11. Metals Most elements on the periodic table are metals.
Conduct electricity and heat.
Ductile
Can be drawn into a wire.
Malleable
Can be hammered or rolled into sheets.
Usually lustrous
Look shiny.
Dull in air or oxygen.
Solids at room temperature (except Hg).
About ¾ of the table consists of metals. Most are solids – the exception being Mercury.

12. Nonmetals Opposite characteristics from metals:
Do not conduct electricity and heat well.
Not very ductile.
Are not lustrous.
Can be solids, liquids, or gases at room temperature.
About ¾ of the table consists of metals. Most are solids – the exception being Mercury.

13. Transition Metals Groups 3-12.
d block elements.
Can lose a different number of valence electrons.
Less reactive than alkali and alkaline earth metals.
Some like Pd, Pt, and Au are very unreactive.
The middle of the periodic table is the transition metals. They vary in properties, but most form vivid colors when in compounds. Unfortunately many of these are toxic to the environment and we won’t get a chance to see them in lab.

14. Rare Earth Metals f block- 2 rows at the bottom of the table.
Fit into rows 6 & 7 (look for * or other symbol).
Lanthanide & Actinide series
Lanthanides = 4f
Reactive like alkaline earth metals.
Some used to produce color on TV screens.
Actinides = 5f
All of them are radioactive.
Nuclei are unstable and break down.
Most periodic tables list these two rows at the bottom of the table, but you may see a really ‘wide’ table that shows them placed in the table where they belong according to electron configuration. For simplicity, they are put at the bottom to make the table more visually appealing.

15. Other Properties of Metals
Varying melting points.
Example: W = 4322oC and Hg = -39oC
Used to make alloys.
Alloys: Homogeneous mixtures of metals.
New properties result from mixing metals.
Example: Brass = copper and zinc.
Harder than copper alone.
More resistant to corrosion.
Others include steel, stainless steel, sterling silver.
We talked about alloys in Chapter 1 – don’t forget they are homogeneous mixtures. New properties get rid of some disadvantages/undesired properties of the pure element.

16. Groups Main group elements – s & p block elements
Groups 1,2 and 3-8
Group 1 = Alkali Metals
H is NOT included
Group 2 = Alkaline Earth Metals
Group 7 = Halogens
Group 8 = Noble Gases
Remember: the group/column number tells you how many valence electrons those elements have!
There are other groups that have names- other columns are simply named by the first element, such as the Carbon family (group 14). Students should label these names on their periodic tables!

17. Alkali Metals VERY REACTIVE !
React with water to make alkaline solutions.
Stored in oil to keep them from reacting with air and water.
Only 1 valence electron to lose- a filled valence shell is very stable.
Not found pure in nature, but combined with other elements (as compounds).
Soft – can be cut with a knife.
Usually lustrous but will dull in contact with air.
Form an oxide layer.
Group 1 is very reactive – you will see a video that demonstrates many of them reacting with water. I can show you Li, Na, and K. After that it is too dangerous for high school classrooms to store these elements (also too expensive).

18. Alkaline Earth Metals Also highly reactive.
Less reactive than alkali metals.
Have 2 valence electrons to lose.
Found as compounds, rather than pure substances.
Harder and higher melting points than group 1.
These elements do not react violently with water, although many of them still do react with water (just slowly). They are found in nature as compounds, the example on the slide is an emerald and aquamarine (Beryllium).

19. Halogens Most reactive nonmetals. 7 valence electrons.
Only need to gain one more electron to be stable.
Frequently react with alkali metals.
Recall that alkali metals have 1 valence electron to lose.
Ex: NaCl, KF, LiBr
Compounds formed from halogens typically are called salts.
Notice that in the examples an element that likes to lose one electron (alkali metals) is paired/bonded with an element that likes to gain one electron (halogens). One of the latest innovation in lights includes halogen lights. These contain at least one of the halogens from Group 17 as a gas instead of using one of the inert gases (Group 18). Picture- sodium fluoride tablets sold to help prevent cavities (fluoride does this).

20. Noble Gases Outermost energy level is completely filled with e-.
s2p6
Exception: He, which has s2. But 1st energy level does not have p sublevel, so it is filled.
Low chemical reactivity – very stable. They have no desire to gain or lose electrons!
Example – He used for blimps.
Typically inert – thought to be completely unreactive.
Exception: 1962, chemists were able to make some compounds with Xe.
The Noble gases are not really unreactive, just mostly – they do form some compounds, but are special situations since they do have full outer electron levels.

21. Hydrogen Most common element in the universe.
Group by itself – very unique.
Only 1 proton and 1 electron.
Can gain or lose an electron.
Hydrogen can gain or lose an electron – remember it only has one electron. This makes it unique. Even though it is found at the top of column 1, it is NOT an alkali metal. It actually is classified as a nonmetal.

22. Lesson Essential Question:
What trends can be found on the periodic table?

23. Section 3: Trends in the Periodic Table
Periodic trends exist since properties of elements repeat in the table.
We will look at the following trends:
ionization energy (IE)
atomic radius
electronegativity (e- neg)
ionic size
electron affinity
melting & boiling points
Since the periodic table is set up according to electron config, we can see a lot of recurring properties and trends. Trends allow us to predict properties when we cannot look up exact values.

24. The Basics.
Shielding Effect: inner electrons shield/block the valence electrons from the positive nucleus.
Results in less attraction.
Increases going down, stays the same going across.
(Effective) Nuclear Charge: how well valence electrons can feel the nucleus’ positive charge.
More shielding (inner electrons) = less nuclear charge felt by the valence electrons.
Decreases going down (due to shielding).
Increases going across.
Inner electrons stay the same, so does shielding.
Protons increase.
Since the periodic table is set up according to electron config, we can see a lot of recurring properties and trends. Trends allow us to predict properties when we cannot look up exact values.

25. Ionization Energy Increases Decreases
Energy needed to remove an e- (forms an ion- atom with a charge).
Decreases down a group.
More energy levels between nucleus and valence e-.
Shielding effect increases, effective nuclear charge decreases.
Increases across a period.
Energy levels between the nucleus and valence e- stay the same.
Increased nuclear charge (more p+ and e-, but e- are being added to the same energy level).
Noble gases have the highest values (very stable).
The way I put the arrows in this presentation will always point to where the trend increases. This is not the same as the book, but I feel its easier to remember the arrow head always points the highest values.
Remember the shielding effect only exists in a group – it has no effect in a period when energy levels stay the same.
Nuclear charge does increase in a group, so it can also explain trends.

26. Atomic Radius (Atomic Size)
Decreases
Inc reases
Atomic Radius (Atomic Size)
Half the distance between 2 bonded atoms’ nuclei.
Hard to measure exact size due to e- cloud.
Bond radius is easier to measure- then cut in half.
Increases down a group.
Energy levels added farther and farther from the nucleus.
Shielding effect increases.
Decreases across a period.
Shielding remains the same, effective nuclear charge increases.
Electrons feel the attraction more and are pulled closer.
I often think of atomic radius as atomic size – how big are the atoms? This is usually easier to remember and visualize.

27. Atomic Radius Diagram Measured in picometers (pm) or Angstroms (Å).
Where should we consider the outside of the atom to be?
Measured in picometers (pm) or Angstroms (Å).
Can see why it’s easier to use the distance between two bonded atoms’ nuclei and divide it by two, rather than trying to measure the radius with just one atom. Because electron location depends upon probability, we have electron clouds (as discussed in chapter 3). Therefore, it becomes difficult to say exactly where the outside of the atom is.
distance between two bonded atoms’ nuclei 2

28. Atomic Radius Cont.
Noble gases are an exception- because they don’t bond with other atoms, we cannot measure their atomic radius the same way as other atoms- must use Van der Waals forces, which don’t ‘smoosh’ the atoms as much, and therefore results in a larger measured radius.

29. Electronegativity
Ability of an atom to attract e- when bonded with another atom.
Electrons from each atom are involved when atoms bond.
Each atom’s ability to attract e- is different.
Linus Pauling invented a scale to indicate how well an atom can attract an e- in a bond.
No units, just numbers.
Ranges from 0 – 4.0.
F assigned 4.0 (highest value- has the highest ability to attract e- when bonded).
Noble gases don’t have a value (don’t need to form bonds- they are stable).
We will use the table of electronegativity in the next chapter when we look at bonding. Depending on the values, it will determine whether a compound is ionic or covalent.
Linus Pauling is an American chemist. He is know for promoting the use of vitamin C to keep healthy.

30. Electronegativity Increases Decreases Decreases down a group.
Shielding effect increases.
Effective nuclear charge decreases.
Thus it’s tougher to attract electrons in a bond.
Increases across a period.
Shielding stays the same.
Effective nuclear charge increases.
Thus it’s easier to attract electrons in a bond.
Another way to look at it: as more e- are added, elements get closer to noble gas configuration which is very stable!
The closer to stability the better.

31. Other Trends- Honors Ionic size Electron affinity
Ions: atoms that have lost or gained e- (have a charge).
Same trend as atomic size.
Same for both positive and negative ions.
Electron affinity
Ability to gain e- (when not bonded).
Same trend as electronegativity.
Melting & boiling points (see pg. 140)
Very different trend from ones previously discussed.
Peaks appear in the p and d blocks across a period.
Tend to increase as atoms approach half filled p and d sublevels, and then decrease.
More opportunities for atoms to bond together.
These trends are not as important but worthy of discussion. We will do a lab that will look at other trends. It is especially good to note that a trend is just that – it is not true 100% of the time. There are exceptions to many of the trends we looked at. If you look at the figures in the book you will see several of these exceptions. We are not concerned with the exceptions, but we should know that some exist.

32. Lesson Essential Question:
How have the elements been created?

33. Section 4: Where Did the Elements Come From?
Only 93 of the elements are found in nature.
3 of these are not found on Earth.
Technetium, Promethium, Neptunium
Found in stars.
Most living things contain C,H,O,N,P, & S.
Compounds that are carbon-based are called organic compounds.
Found in living things.
Big Bang Theory: these elements were created when universe was formed in a violent explosion.
Some elements exist here on earth, but in such small quantities, that is more advantageous to synthetically prepare them. Francium is a good example of this.

34. Big Bang Theory Cont.
VERY high temperatures existed after the big bang. This form of energy cooled and formed matter (e-, p+, n).
Further cooling allowed subatomic particles to join together to form H.
Gravity pulled H clouds together and formed stars.
Stars worked as nuclear reactors to form He (under high temperature and pressure).
4 H  1 He + energy (gamma radiation)
Other elements were formed as He and H combined (fusion) to form even heavier elements.
Supernovas formed all elements heavier than iron.
Star collapses and blows up, releasing heavier elements into space.
Remember that the big bang theory is just that – a theory. It is however believed by most scientists, since the evidence we have gathered points to this being true.
During a supernova, the star can emit more energy than the sun will emit over its entire life span!

35. Supernovas

36. Synthetic & Superheavy Elements
Transmutations: type of nuclear reactions that change one element into another element
All elements greater than number 93 (except 61) are not naturally occurring– synthetic elements.
Particle accelerators can be used to create them. Different types exist.
Nuclei collide and fuse together.
Superheavy elements are those that have an atomic number greater than 100.
Only exist for fractions of a second.
In the Middle Ages, alchemists tried to change materials into gold – they were not successful. They did not understand nuclear reactions. We know today that elements can be created from other elements. It does not really pay to make gold this way!!