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The Periodic Table. Our current society takes for granted all of the hard work, research, chance, and luck that has gone into creating and discovering the materials that are used in the products we utilize every day.

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how d they come up with that
Our current society takes for granted all of the hard work, research, chance,

and luck that has gone into creating

and discovering the materials that are used in the products we utilize every day.

For example, who was the first person to set or find a random black rock (coal) on fire and discover that it provided a good, constant source of heat?

Who was the first person to discover that a substance found in some rocks was capable of being the ultimate explosive (uranium)?

How’d They Come Up With That?



organizing the elements
Organizing the elements

In nature and in the lab we have discovered over 100 different elements.

We’ve organized the elements into a table based on their PHYSICAL and CHEMICAL PROPERTIES

It took us almost 2000 years to figure out the properties of the elements currently in the Periodic Table of Elements and arrange them.

i developing the periodic table
I. Developing the Periodic Table
  • In the early 1800s, enough information was known about the elements that scientists wanted an easy way to categorize the Earth’s ingredients.
so why don t we just list the elements in a straight line why make a table
So why don’t we just list the elements in a straight line? Why make a table?
  • An analogy: When you go to the grocery store and are looking for a specific ingredient, you want things to be grouped in a logical way so that you can easily find what you need.
  • Scientists wanted the elements that they were using in their experiments to also be grouped so that they could easily find elements with the characteristics that they desired.
developing the periodic table
Developing the Periodic Table
  • Many methods of organization were tried before scientists found the most effective way of grouping the elements
dmitri mendeleev
Dmitri Mendeleev
  • In 1869 he published a table of the elements organized by increasing atomic mass.


developing the periodic table1
Developing the Periodic Table
  • Mendeleev stated that the properties of elements are a periodic function of their atomic masses.
  • If we lay out all of the elements in order by atomic mass, we would see the same properties showing up periodically
      • For example, as you would begin down the list of the first 18 elements, every 9th (or so) atom would be super reactive and want to gain 1 electron
lothar meyer
Lothar Meyer
  • At the same time, he published his own table of the elements organized by increasing atomic mass.


father of the modern periodic table

Both Mendeleev and Meyer arranged the elements in order of increasing atomic mass.

  • Both left vacant spaces where unknown elements should fit.
  • So why is Mendeleev called the “Father of the Modern Periodic Table” and not Meyer, or both?
Father of the Modern Periodic Table

Mendeleev …

  • used his periodic function theory and corrected the atomic masses of Be, In, and U.
  • was so confident in his table that he used it to predict the physical properties of three elements that were yet unknown (Sc, Ga & Ge).
  • After the discovery of these unknown elements between 1874 and 1885, and the fact that Mendeleev’s predictions were amazingly close to the actual values, his table was generally accepted.
mendeleev s periodic table1
Mendeleev’s Periodic Table
  • However, in spite of Mendeleev’s great achievement, problems arose when new elements were discovered and more accurate atomic weights determined.
  • By looking at our modern periodic table, we can identify some problems which might have caused chemists a headache.
  • Atomic Masses
  • Ar and K (40-39)
  • Co and Ni (59-58)
  • Te and I (128-127)
  • Th and Pa (232-231)
henry moseley
Henry Moseley
  • In 1913, through his work with X-rays, he determined the actual nuclear charge (atomic number) of the elements*. He rearranged the elements in order of increasing atomic number.

*There is in the atom a fundamental quantity which increases by regular steps as we pass from each element to the next. This quantity can only be the charge on the central positive nucleus.

His research was halted when the British government sent him to serve as a foot soldier in WWI. He was killed in the fighting in Gallipoli by a sniper’s bullet, at the age of 28. Because of this loss, the British government later restricted its scientists to noncombatant duties during WWII.


periodic law
Periodic Law
  • This arrangement allowed elements with similar properties to be lined up into the same column in the periodic table
    • Ex. All of the elements in group 1 are extremely reactive when by themselves and want to give away 1 electron
glenn t seaborg
Glenn T. Seaborg
  • After co-discovering 10 new elements, in 1944 he moved 14 elements out of the main body of the periodic table to their current location below the Lanthanide series.
  • These became known as the Actinide series.


glenn t seaborg1
Glenn T. Seaborg
  • He is the only person to have an element named after him while still alive.

"This is the greatest honor ever bestowed upon me - even better, I think, than winning the Nobel Prize."


video review
Video - Review
  • Periodic Table History (6:33)
periodic table geography

Periodic Table Geography

Where are elements found?


The elements in any group of the periodic table have similar physical and chemical properties!

The vertical columns of the periodic table are called GROUPS, or FAMILIES.

general properties of metals
General Properties of Metals
  • Metals are good conductors of heat and electricity
  • Metals are malleable
    • Easily bent and shaped
  • Metals are ductile
    • Able to be drawn out into a wire
  • Metals have a shiny luster
examples of metals
Examples of Metals

Copper, Cu, is a relatively soft metal, and a very good electrical conductor.

Zinc, Zn, is more stable than potassium

Mercury, Hg, is the only metal that exists as a liquid at room temperature

general properties of metalloids
General Properties of Metalloids
  • They have properties of both metals and nonmetals.
  • Metalloids are more brittle than metals, less brittle than most nonmetallic solids
  • Metalloids are semiconductors of electricity
  • Some metalloids possess metallic luster

Ex: Silicon possesses a metallic luster, yet it is an inefficient conductor and is brittle.

general properties of nonmetals
General Properties of Nonmetals
  • Nonmetals are poor conductors of heat and electricity
  • Good insulators
  • Nonmetals tend to be brittle when solids
  • Many nonmetals are gases at room temperature

Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element.

examples of nonmetals
Examples of Nonmetals

Microspheres of phosphorus, P, a reactive nonmetal

Sulfur, S, was once known as “brimstone”

Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure


The Periodic Table of Elements

Write the group names!




Noble Gases

Alkali Metals

Alkali Earth Metals

Transition Metals

Transition Metals

Transition Metals

Transition Metals



noble gases

Noble Gases

Last slide

iii periodic law
III. Periodic Law
  • Before we can discuss the characteristics of each group, we must introduce the Period Law.
periodic law1
Periodic Law
  • When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.
periodic table review
Periodic Table Review
  • Elements with similar e- configurations are placed in the same group. (same number of e- in last energy level)
  • The number of electrons in the outermost energy level determines how reactive an element will be
  • These electrons in the highest energy level are called “Valence Electrons”
    • Therefore elements in the same group have the same number of valence e-
group 1
Group 1
  • EX: H = 1s1

Li = 1s22s1

Na = 1s22s22p63s1

  • All elements in group 1 have 1 valence electron (1 electron in highest energy level)
group 17
Group 17
  • Ex: F = 1s22s22p5

Cl = 1s22s22p63s23p5

Br = 1s22s22p63s23p64s23d104p5

  • All elements in group 17 have 7 valence electrons (7 electrons in highest energy level)
valence electrons
Valence Electrons



  • Highlight the valence electrons on your periodic table







They differ: 2-3


octet rule
Octet Rule
  • Octet Rule = Eight electrons in an outer level render an atom unreactive (stable). Therefore all atoms are trying to get to this state of 8 valence electrons.
  • Atoms obtain an octet of electrons in their valence orbital by
    • taking electrons from another atom,
    • giving their valence electrons away,
    • or sharing electrons with another atom.
octet rule1
Octet Rule
  • Whether an atom gains, loses or shares e- to form an octet depends on how close the atom is to having the 8 e-.
    • If an atom has less than 4 valence e- then it will try to lose them so that the energy level below (which is full w/ an octet) will then be considered the outermost energy level.
      • Ex: Na = 1s22s22p63s1
        • If Sodium loses 1 electron in the 3s energy level, then energy level 2 will be considered the outermost energy level. Since it has an octet the atom will be stable.
octet rule2
Octet Rule
  • If an atom has more than 4 valence electrons it will try to steal electrons from other atoms to form an octet
    • Ex: Cl = 1s22s22p63s23p5
      • If Chlorine gains 1 electron it will have an octet and be stable.
octet rule3
Octet Rule
  • Atoms with 4 electrons in their valence shell will share4 electrons with other atoms to feel as though they have an octet.

Last slide

iv oxidation numbers
IV. Oxidation Numbers
  • Our knowledge of e- configurations and the stability of noble gases allows us to predict oxidation numbers for elements.
  • Oxidation numbers represent the charge an ion obtains after losing or gaining valence electrons.

Ex. Ca = 1s22s22p63s23p64s2

2 valence e-, will want to get rid of them so that the full 3rd energy level (8 electrons) will then be it’s outermost shell


oxidation numbers
Oxidation Numbers

Ex. Ca = 1s22s22p63s23p64s2


*** Octet Rule: atoms want to have a full highest energy level (8 electrons)

** If an atom has less than 4 valence electrons – they will lose their valence electrons so that the full energy level below will then be considered the outermost shell, and become a positive ion (cation)

oxidation numbers1
Oxidation Numbers

Ex. F = 1s22s22p5

F- = 1s22s22p6

*** Octet Rule: atoms want to have a full highest energy level (8 electrons)

** If an atom has more than 4 valence electrons – they will gain (steal) electrons until they have 8 electrons (full energy level – octet), and become a negative ion (anion)

oxidation numbers2
Oxidation Numbers

*on your periodic table, write in the oxidation number of each group




Tend to have more than one oxidation number







3+ or 4+

  • What is the element with the electron configuration of: 1s22s22p63s2
  • How many valence electrons?
  • What would be the oxidation number?
  • What is the element with the electron configuration of:

[Ar] 4s23d104p3

  • How many valence electrons?
  • What would be the oxidation number?


2 valence e-


  • Arsenic
  • 5 valence e-
  • As-3
alkali metals1
Alkali Metals
  • Group 1 Metals
  • Soft silver metals
  • Less dense than other metals so melt and boil easily
  • All alkali metals have 1 valence electron
    • This makes them very reactive because they want to get rid of it to be stable (octet in energy level below)
  • NEVER found alone in nature – are always bonded to something else; they are too reactive

Potassium, K reacts with water and must be stored in kerosene

alkaline earth metals1
Alkaline Earth Metals
  • Group 2 Metals
  • Shiny silvery - white
  • All alkaline earth metals have 2 valence electrons
    • This makes them very reactive because they want to get rid of it to be stable (octet in energy level below)
  • Alkaline earth metals are less reactive than alkali metals
  • Alkaline earth metals are not found pure in nature; they are too reactive
  • Found in Earth’s crust in mineral form.
  • Group 17 Nonmetals
  • Halogens all have 7 valence electrons
    • This makes them very reactive because they are SO CLOSE to having an octet, just need 1 more e-
  • Halogens are never found unbonded in nature; they are too reactive
  • Halogens in their pure form are diatomic molecules (F2, Cl2, Br2, and I2)
  • Most important group to be used in industry

Chlorine is a yellow-green poisonous gas

  • Group 16 nonmetals
  • Have 6 valence electrons
    • Semi-reactive
  • Diverse group that includes nonmetals, metalloids, and metals
noble gases1
Noble Gases
  • Group 18 Nonmetals
  • Noble gases have 8 valence electrons
  • (except helium, which has only 2)
    • This makes them extremely stable and unreactive because they have filled valence shells. They don’t bond/react with other elements because they don’t need to gain or lose e-)
  • Noble gases are ONLY found alone in
  • nature – monatomic gasses
  • Colorless, odorless; they were among the last of the natural elements to be discovered
  • Which elements have the most similar chemical properties… Those in the same family or those in the same period
  • Which of the groups on the periodic table are too reactive to be found as monoatomic atoms in nature?

Alkaline metals & alkaline earth metals

Same Family

using the periodic table to predict properties of elements
Using the Periodic Table to Predict Properties of Elements
  • The basis of the periodic table is the atomic structure (number of protons, neutrons and electrons) of the elements.
  • Position on the table and properties of these elements arise from the e- configurations of the atoms.
  • Properties such as density, atomic radius, oxidation numbers, ionization energy, and e- affinity can be predicted.
a atomic radius size of atoms
A. Atomic Radius (size of atoms)
  • As principal quantum number increases (you go down a group), the size of the electron cloud increases (more orbitals).
  • Size of atoms increase moving down Per. Table.
atomic radius
Atomic Radius
  • Atoms in the same period have the same quantum number; however, positive charge on the nucleus increases by one proton for each element in a period. This pulls the e- cloud in tighter, decreasing atomic radius.
predicting atomic radius
Predicting Atomic Radius
  • General rule: atomic size increases as you move diagonally from top right corner to bottom left corner.
b ionization energy
B. Ionization Energy
  • The energy required to remove an e- from an atom.
  • Remove the most loosely held e- is first ionization energy.
  • Measured in kilojoules per mole (kJ/mol)
ionization energy
Ionization Energy
  • As you move down the group, the ionization goes from high to low ionization energy. The larger the atom, the less energy is required because the e- are farther from the positive center.
ionization energy1
Ionization Energy
  • As you move from left to right across a period, the ionization goes from low to high ionization energy.
classification based on first ionization energy
Classification based on First Ionization Energy



  • 1. Low 1st ionization energy.

2. Located on left side of Periodic Table.

3. Form positive ions

4. Will give away their electrons.

Ca = [Ar] 4s2

2 valence e-

positive ion: Ca+2

  • 1. High 1st ionization energy.

2. Located on the right side of Periodic Table.

3. Form negative ions

4. Will NOT give away their electrons, they want to gain electrons.

F = [He] 2s2 2p5

7 valence e-

negative ion: F-

multiple ionization energies
Multiple Ionization Energies
  • Additional e- can be lost from an atom and the ionization energies can be measured. Notice that the energy increases greatly to remove more e-

IONIZATION ENERGIES (kilojoules per mole)

Element 1st 2nd3rd 4th 5th

H 1312.0

He 2372.3 5220

Li 520.2 7300 11750

Be 899.5 1760 14850 20900

B 800.6 2420 3660 25020 32660

c ionization energies
C. Ionization Energies
  • It increases from left to right, bottom to top and bottom left to top right corners.
  • Ionization Energy (1:52)
c electronegativity
C. Electronegativity
  • Electronegativity is the ability of an atom to capture an electron.


  • As you move from left to right across a period, the electronegativity increases because elements on the left side want to lose electrons to be like the nearest noble gas (octet rule) so do notwant to grab electrons(low electronegativity)
  • Elements on the right side want to gain electrons to be like the nearest noble gas, so the will grab electrons a lot. (high electronegativity)
  • They have no electronegativity at all because they are noble gases, and don’t need any more electrons. (No electronegativity)
  • As you move from higher to lower in a group, the electronegativity decreases due to the shielding effect (which states that electrons in outer energy levels are held less tightly than those in lower energy levels). 
  • Generally, if an atom doesn’t hold the electrons it has very much, it won’t grab electrons from other atoms much, either.
c electronegativity1
C. Electronegativity
  • It increases from left to right, bottom to top and bottom left to top right corners.
  • Based on our trends:

The most reactive metal element would be


The most reactive nonmetal element would be


Last slide

in summary
In Summary
  • Periodic table is a chart of elements in which the elements are arranged based on their e- configurations which dictates their properties.
  • Moving down a group in the periodic table, atomic radii becomes larger because more energy levels are needed for more e-.
in summary1
In Summary
  • As the size becomes larger, the e- are located farther away from the positive center.
  • This decreases the affinity of that atom to hold on to these outer e-, thus decreasing e- affinity.
  • Ionization energy is low because it is easy for the atom to lose these outer e-.
in summary2
In Summary
  • Moving across a period in the periodic table, atomic radii becomes smaller because the energy levels of periods are the same but the positive centers of atoms increase. This pulls the e- cloud closer to the nucleus, making the atom smaller.
  • Ionization energy increases for these smaller atoms.