1. The Periodic Law
[CHEM1.PS1:Matter & Its Interactions] – Explain the origin & organization of the Periodic Table of Elements. Predict chemical & physical properties of main group elements (reactivity, ion charge, ionization energy, atomic radius & electronegativity).

3. History of the Periodic Table
Who is the father of the periodic table?
Dmitri Mendeleev ~1869
He initially arranged ~60 elements in order of increasing atomic mass. 2 questions remained…
1. Why could most of the elements be arranged in the order of increasing atomic mass but a few could not?
2. What was the reason for chemical periodicity?
Give an example of something that is periodic.
Henry Mosely, who worked with Ernest Rutherford discovered by viewing spectral data, that the table fit into patterns better when arranged by increasing nuclear charge – or increasing _________ __________.

4. History of the Periodic Table
Periodicity or the periodic law states that the physical and chemical properties of the elements are functions of their atomic numbers – when elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals.

5. Electron Configuration and the Periodic Table
Based on electronic configurations of the elements, the periodic table can be divided into four blocks  s, p, d, and f.

6. All atoms ”seek” the lowest possible energy state naturally available to them.
The lowest possible energy state naturally available is an outer electron shell that is completely filled.
Atoms that are isoelectronic to noble gases have the lowest possible potential energy state.
Valence (outer) electrons govern chemical properties IE what they react with & how they react.
Allotropes are elements that have two or more distinct sets of chemical and physical properties.
Oxygen O2 & Ozone O3 are allotropes.
Graphite, diamond and buckminsterfullerene C60 are three allotropes of carbon.
Sulfur & phosphorous also have allotropic forms
Hydrogen does not.

7. Electron Configuration and the Periodic Table
Alkali metals are group 1 on periodic table.
silvery
soft enough to be cut with a knife
very reactive
cannot be found in nature as free elements
react strongly with water
Alkaline earth metals are group 2 on the periodic table.
harder, denser and stronger than group 1
higher melting points than group 1
less reactive than group 1
Groups 1 & 2 are also known as the s-block elements.

8. Electron Configuration and the Periodic Table
d-block elements are called transition metals.
good conductors
high luster
typically less reactive than groups 1 & 2
exist in nature as free elements
Three very unreactive elements listed by your text…
palladium, Pd
platinum, Pt
gold, Au

9. Electron Configuration and the Periodic Table
p-block elements are called the main-group elements
all metalloids
all nonmetals (except hydrogen and helium)
the halogens (group 17)
the halogens are the most reactive nonmetals
why?
post transition metals are in the p-block and include gallium, indium, tin, thallium, lead, and bismuth.
f-block elements are the lanthanides and actinides
lanthanides are shiny and similar in reactivity to group 2
actinides are mostly laboratory-made elements

10. Electron Configuration and Periodic Properties
_________ ___________ may be defined as one-half the distance between the nuclei of identical atoms that are bonded together.
From left to right atomic radii trend smaller. Why?
As atomic numbers get bigger (one per element), this means that each subsequent element has one more ___________. Because of this increasing nuclear charge (or positive charge), the electrons orbiting the nucleus are pulled in closer and closer.
From top to bottom atomic radii trend larger. Why?
As elements trend down a group, more s-sublevels are added and so the 1s is small, but each s sublevel is a bit larger than the previous.

11. Electron Configuration and Periodic Properties

12. Ionic Radii
Cations have lost one or more electrons and have a (+) charge.
Anions have gained one or more electrons and have a (-) charge.
Cations are always smaller than neutral atoms of the same element. (more protons than electrons so nuclear charge pulls in electrons and make electron cloud smaller)
Anions are always larger than the neutral atoms. (extra repulsive forces between added electrons)

13. Metallic character of elements increases from top to bottom of a group.
Which has more metallic character: Nitrogen or Bismuth?
Beryllium or Barium?
This is because electrons become easier to lose as the atomic radius increases. The increase in atomic radius decreases attraction between the positive nucleus and the negative electrons, causing the electrons to be held more loosely.

15. Melting and boiling points of metals decrease from top to bottom.
Melting and boiling points of nonmetals increase from top to bottom.

16. __________ __________ are available to be lost, gained, or shared in the formation of chemical compounds. They are often located in incompletely filled outer orbitals. They are the electrons that participate in chemical bonding.
Often, the charge of valence electrons is concentrated closer to one atom than to another. This uneven concentration of charge has significant effects on the chemical properties of a compound.
______________ is a measure of the ability of an atom in a compound to attract electrons from another atom in a compound. The most electronegative element is fluorine.
Electronegativity generally increases left to right and decreases top to bottom …Why?
…because, although there is a dramatic increase in nuclear charge, the bonding electrons are in much higher energy levels so are much farther from the nucleus. There is also more shielding of the attractive force (protons in the nucleus attracting bonding electrons) by electrons in the lower energy levels; this decreases the effect of the nuclear charge.

17. Electron Configuration and Periodic Properties

18. Electron affinity is defined as the energy change that accompanies the addition of an electron to an atom.
Some atoms readily attract electrons and EA has a negative value, meaning energy is released.
Most atoms do not accept additional electrons readily and the EA is a positive value, indicating that energy must be USED to add the electron.
Fluorine has the highest EA & francium has the lowest. Trend much like electronegativity.

19. Electron Configuration and Periodic Properties

20. An ______ is an atom or group of bonded atoms that has a positive or negative charge.
Any process that results in the formation of an ion is referred to as _____________.
Ionization energy (IE) is the energy required to remove an electron from a neutral atom of an element. This energy is always endothermic.
IE increases left to right (across a period) and decrease top to bottom (down a group). Why?
As elements become more and more like the noble gases, more and more energy is required to remove an electron and create an ion.
Lots of energy is required to disrupt a completely filled octet.
Predicting IE of any atom is made easier when an orbital notation is drawn out for the atom to visualize electrons in their valence orbitals.

21. Electron Configuration and Periodic Properties

22. The Periodic Law Slide Quiz
In the modern periodic table, elements are arranged according to what?
Group 17 (halogens) are the most reactive of the nonmetals because?
What does the periodic law state?
Across period 3 (left to right), the energy needed to remove an electron from an atom (increases/decreases)?
Which element has the highest electronegativity?
How does ionization trend?
How does atomic radius trend?
How does electronegativity trend?

23. Funny Chemistry