1. The Periodic Table The most fun you can have with a bunch of squares.
Take a Moment to make a list of two or three elements you like, or are familiar with.
What do you know about that element? Why do you remember it?

2. Why do we organize the elements?
A quick history:
Many elements were familiar since ancient times. Au, Ag, Hg, Sn, C, Pb, Cu, Fe and S, had been discovered by various cultures.
The practice of alchemy led to the discovery of Phosphorus, by Hennig Brand around He hides his discovery.

3. What is an element? Early definition Earth, Air, Fire, Water, Ether….
Later definitions included the discovery that one element could not be transformed into another by ordinary means.
The discovery of radioactive decay alters this idea, but mainly, this is true.
Our definition? An element will always have the same atomic number!!!

4. Why do we organize the elements?
Between 1669 and 1869, the work of discovering and documenting new elements becomes one of the focuses of practitioners of science everywhere.
Practitioners of science everywhere are beginning to look for patterns.
Why?
1) They wanted to discover more.
2) Understanding relationships between elements would further their understanding of reactions, and compounds.

5. Attempts at Organization
Johann Dobereiner – begins to work with element properties and starts associating them in groups of three, called triads.
A.E. Beguyer de Chancourtois – makes a cylindrical table, first geometric representation of patterns found in elemental behavior.
John Newlands – proposes a musical analogy and organizes them into octets, groups of eight. Also shows the patterns, but is too limiting.

6. A.E. Beguyer de Chancourtois

8. Dmitri Mendeleev
Publishes a table in 1869 that gives rise to the modern table layout
He organizes the elements by mass, which is the information they have at the time
Some of the masses are incorrectly calculated so the table is not entirely in order
His table is considered the most influential because it leaves room for expansion and anticipates the existence of elements of other masses, leaving spaces for them to be placed there in the future.

9. Atomic Number
Later, work by Henry Moseley enables chemists and physicists to accurately find the number of protons in the nucleus of an atom.
Using this information the table is placed in the order we use today.
THE ELEMENTS ARE IN ORDER BY THEIR ATOMIC NUMBERS.
This allows patterns to emerge.

10. Elements are in groups and periods
Groups are also known as families, these are elements in columns, that share similar chemical behavior.
Periods are rows across the table.
While a regents chemistry table will show electron configurations, your “naked” tables will not, so you should know 2 things:
1) All elements in a group have the same # of valence electrons
2) All elements in a period have the same number of energy levels (shells).

12. Important Groups
Definition: Alkali – derived from the Arabic word, al-qali, which referred to ashes, which when mixed with water form a basic solution used throughout history to make soap.
Groups known as alkali will make salts, (you will recognize later as ionic compounds), that form basic solutions when placed in water.
Group 1 = Alkali metals

13. Alkaline Metals Group 1: Alkali metals:
- Form ions with 1+ charge by losing one electron.
- Do not occur in their metallic form in nature,
they are too reactive. They are always bound in a
compound.
- These are the most reactive metals on the table!!

14. Lithium Reacting with Water

15. Earth Alkaline Metals Group 2: Known as the Earth Alkaline Metals.
- Will acquire 2+ charges due to the loss of two valence electrons
- Also very reactive, these will also not be found in nature in their metallic forms.
- The elements Ra – Radium is a naturally occurring radioactive element, may not behave as others in the group.

16. Strontium and Barium are more reactive and are stored in oil.

17. Magnesium burns and is used for survival and camping type fire starting.

18. Transition Metals
Known as transition metals because they make the transition to using the d orbitals for bonding.
Transition metals behave differently than group 1 and two metals, and are often able to acquire a variety of charges, depending on which element they are reacting with.

19. Transition metals form colored compounds and solutions!

20. Other Metals

21. Metalloids

22. THESE ARE THE MOST REACTIVE
The Halogens
Group 17 are elements known as the Halogens
When they form compounds, they are often referred to as halide compounds
Halogens tend to gain one electron from other elements while bonding, so they acquire 1- charges
THESE ARE THE MOST REACTIVE
NON-METALS

23. The Noble Gases Group 18 on the table.
Often called INERT gases, under normal circumstances they will not react with anything.
Have a complete valence, with eight electrons, accounts for their reluctance to react.
As atoms of a gas they are always MONATOMIC, they are never diatomic. Their atoms won’t even bond with each other.

24. The Noble Gases
Adventures in literacy: during science classes you will learn many Latin and Greek prefixes that can help with word decoding. Because the noble gases are so non-reactive, they were difficult to discover.
Noble Gas Greek Meaning Vocabulary/SAT Word
Helium Helios – Sun Heliocentric
Neon Neos – New Neophyte, Neoclassical
Krypton Hidden Cryptography, Cryptic
Xenon Stranger Xenophobic
Argon Lazy, inactive

25. The Diatomic Gases Sometimes called the magic seven
Br, I, N, Cl, H, O and F
Form the shape of a seven on the table, except for Hydrogen
These gases all form diatomic molecules in order to become more stable. Several are very reactive, Nitrogen, N2, is the least reactive and most stable.
In most cases you can assume that elemental diatomic gases should be written as Cl2, etc. Particularly if you have been told in words that it is its gaseous form.

26. Understanding Patterns in Compound Formation
We noted earlier that elements in groups 1, 2, 13, 16 and 17 tend to pick up certain charges.
These elements will pair with one another when reacting so that their charges balance/cancel out to zero.
You will sometimes be asked to identify an element based on it’s charge.
You will be given main charges, label them on your table.

27. Balancing Charge in Compound Formation
What element could form a compound with the formula MgX2 ? How about X2O?
What element could form a compound with the formula
XCl?

28. Strength Does not lie in what you have. It lies in what you can give
Strength Does not lie in what you have. It lies in what you can give. Quote I found on the tag of a herbal tea bag. Of course it reminded me of the metals. Remember Metals react by giving up electrons!!!!

29. Elemental Properties Metals Non-metals No luster, may be glassy
Gain electrons to acquire negative charges, (exception – Hydrogen).
DO NOT conduct, are insulators because they prevent heat and electricity from passing through
Tend to have low melting and boiling points
Form molecules and are found in molecular substances
Shiny..have metallic luster
Lose electrons to acquire positive charges
Conduct heat and electricity
Are malleable and ductile, (can be made into sheets and wires)
Have high melting and boiling points
React with acids to form Hydrogen gas

30. Properties and Reactivity
Francium represents the most reactive metallic element, so it has the greatest metallic character in its group
Francium is the most reactive metal in this group for the same reason, it is the element that will most readily lose its valence electron
Fluorine represents the most reactive non-metallic element, so it has the least metallic character of the elements in its group.
Fluorine is the most reactive non-metal for the same reason. Fluorine will literally tear an electron off of pretty much any element.
So metals are reacting by giving up an electron and non-metals are reacting by taking an electron.

31. States of Matter
You will need to know the following about what elements are solids, liquids, and gases at room temperature.
Most metals are solids, except Mercury, (Hg), which is a liquid.
Many non-metals are gases, however NOT ALL OF THEM. C, P, S, Se, Te, Po and, At are all solids.
At standard conditions the Halogens are not all gases. Bromine is a liquid, and Iodine is a purple crystalline solid.
The Noble gases are of course, gases.

32. Fluorine and Chlorine are both pale yellow/green gases
Bromine is a liquid, however it will begin to evaporate and form dangerous vapor immediately, see the orange cloud?
Fluorine and Chlorine are both pale yellow/green gases
Iodine forms an almost metallic crystal, which is also very volatile, and sublimes easily into a purple vapor.

33. What are Allotropes?
Remember: the state and the physical properties of a substance are dependent on the arrangement of its atoms or molecules.
Allotropes are forms of elements which have adopted different molecular/crystalline structure. Despite still being pure samples of the element, the differences in structure result in differences in their PHYSICAL properties.

34. Carbon
3 Forms of Phosphorus

35. What is Metallic Character?
We sometimes refer to the elements as having more or less metallic character as you move around the table.
Essentially the closer you are getting to the metals, the more metallic character the elements will have. Therefore they will give up electrons more easily, etc.

36. Characteristics of Metalloids
As you cross from the metals to the non-metals you will encounter the metalloids, also referred to as the semi-metals.
Metalloids have hybrid characteristics. In other words, a metalloid may be brittle, but might conduct electricity. They are literally in between.
The properties of metalloids allows elements like silicon, to be used in electrical applications where metals could not function, for example, in microchips. This is why the birth place of many computer companies in California is referred to as Silicon Valley. They are also used to produce solar cells, solid state data storage, transistors etc.etc.

37. Navigating By Landmarks

38. Effective Nuclear Charge
What type of charge is found on the nucleus of an atom?
How does this effect its electrons?
How does the number of electrons, or the number of energy levels affect the atoms structure?
How does it affect the way the atom will behave?

39. Nuclear Charge VS Energy Levels
The amount of nuclear charge, which is basically the number of protons, affects how many electrons an atom can hold onto.
More importantly, it affects HOW WELL the atom is able to hold on and the AMOUNT OF ENERGY it takes to do so.
You may recall that energy levels that lie closer to the nucleus tend to be lower in energy, while electrons located further away from the nucleus are in higher energy levels.

40. Nuclear Charge VS Energy Levels
The larger an atom becomes, the more energy levels it will have to have to house all of its electrons.
The more electrons there are the more difficult it becomes to pull them close to the nucleus.
HOWEVER… as we said earlier, all the elements located in a group will have the SAME number of energy levels. SO…
Despite having more electrons, they are not moving into higher energy levels and are no more difficult to hold onto.

41. Nuclear Charge VS Energy Levels
But… as you move down a group, you will see that atoms begin to pick up more energy levels. This makes it increasingly difficult to hang onto electrons, or to pull them in close to the nucleus.
Additionally, the effect of more electrons is that they are interfering with the electromagnetic attraction of the nucleus reaching the outermost electrons. This effect is known as SHEILDING. Think of it as the electromagnetic force being partially used up before it can reach the outer electrons at full force. This means the outermost electrons of bigger atoms will be more vulnerable, and will bond, or be lost, more readily to achieve stability.

42. HOMEWORK
1) Re-label your blank periodic tables again!!! Take new ones if needed. You must be able to label all the important information: charge, group names, increasing and decreasing metallic character.
2) Make a table that can be used as a study guide. Do it on the computer or design by hand. The table should go over the properties of the metals, non-metals and metalloids, including information about major groups found in those areas!!! Be ready for tomorrow I will be checking it and collecting it!!!

43. What is Electronegativity?
Have a working definition ready…
What’s a working definition?
- A working definition is an informal way of considering a concept that is more useful for you to use to both navigate questions and apply the concept while working.

44. What is Electronegativity?
For Electronegativity you need to have in your working definition:
- It is a tendency to attract electrons.
- It varies across the table according to a pattern or trend.
- It is impacted by shielding and effective nuclear charge.

45. Can you predict the electronegativity trend across the periods of the table?

46. Navigating By Landmarks

47. Remember your landmarks!!
Fluorine Just happens to be the
MOST ELECTRONEGATIVE
element on the table
Whenever you are headed towards fluorine, the elements are increasing in electronegativity.

48. Ionization Energy
What is your working definition of Ionization Energy?
- It should include energy and the
loss of electrons
- It should include the pattern you find moving from metals to non-metals on the table.

49. Can you predict the trend for Ionization Energy across a period?

50. Remember your landmarks!!
If Fluorine is the MOST ELECTRONEGATIVE element,
it will take the MOST ENERGY to remove an electron from a fluorine atom.
So as you are moving towards Fluorine, Ionization energy is also increasing.

51. Size VS Radius
Before we talk about radius, it’s important to remember that we are talking about size –
IN TERMS OF VOLUME…
In other words, how much space is this atom taking up? How far out are its electrons spreading?

52. Size VS Radius
During a test, when you are nervous, if you forget this trend, you may be tempted to look at the MASS of the elements as you go across a period, and assume that more mass means the atom is getting bigger.
THIS IS NOT TRUE, remember most of an atom is empty space, (thanks Rutherford ). We are not talking about MASS we are talking about HOW FAR THE ATOM’S ELECTRONS ARE GOING TO SPREAD OUT!!!!

53. Radius

54. Radius

55. Ionic Radius Recall that ions are formed by loss or gain of electrons.
Losing electrons: Cation formation, Acquiring a positive charge. Reduces the radius of an atom.
Gaining electrons : Anion formation, Acquiring a negative charge. Increases the radius of an atom.

56. Regents Reference Tables
Take out your Regents Reference Tables, Let’s review the helpful information found here!!!

57. Label the Table
Now that you have graphed, and know the major trends found in the groups and periods, you should add arrows that indicate these trends to your periodic table.
Do this tonight!!! Include Electronegativity, ionization energy, and atomic and ionic radius. By tomorrow you should have produced a table showing ALL the important things you need to know about the periodic table.