1. Development of the Periodic Table
Antoine Lavosier
In the late 1790’s, compiled a list of the 23 elements known at that time.
J. W. Dobereiner
In 1829, published a classification system in which he grouped elements in triads – a set of 3 elements with similar properties.
By this time, almost 70 elements were known.
Unfortunately, all known elements could not be grouped in triads.

2. Development of the Periodic Table
John Newlands
In 1864, proposed an organization scheme with elements arranged by increasing atomic mass.
He observed that properties repeated every eighth element, calling this relationship the “law of octaves” (an octave is a group of musical notes that repeats every eighth tone).
Acceptance of the law of octaves was very little because it did not work with all of the known elements and scientists associated the word octave with music, not science.
Dmitri Mendeleev
In 1869, developed a periodic table in which elements were arranged into groups based on a set of repeating properties.
Left spaces in his table for elements that he predicted would be discovered in the future based on gaps he noticed in the properties of the elements that had been discovered.

3. Development of the Periodic Table
When new elements were discovered, the actual properties of those elements were very close to the predicted properties.
Because of this table’s ability to predict properties of undiscovered elements, it was widely accepted by the scientific community.
This arrangement of the periodic table was based on increasing atomic mass since little was known about atomic structure at this time.
Henry Moseley
In 1913, determined an atomic number for each known element.
He re-arranged the elements by their atomic number – the order observed in the modern periodic table.
The patterns observed became known as periodic law: When elements are arranged in order of increasing atomic number, there is periodic repetition of their physical and chemical properties.

4. The Periodic Table

5. Metals Make up about 80% of the elements
Form cations (positively charged atoms)
Good conductors of heat and electricity
Malleable – can be hammered into thin sheets
Ductile – can be drawn into wire
High luster, or sheen

6. Nonmetals Most are gases at room temperature
Forms anions (negatively charged atoms)
Poor conductors of
heat and electricity
Solid nonmetals tend
to be brittle and dull-
looking

7. Metalloids Staircase arrangement of elements
Has properties similar to those of metals and nonmetals
Many are semi-conductors of heat and electricity

8. Alkali Metals – Group 1 (except H)
Silvery, soft solid
1 valence electron
Forms +1 charge
Very reactive with water
Very reactive, especially with Group 17 elements

9. Alkaline Earth Metals – Group 2
Soft, shiny solids (harder than alkali metals)
2 valence electrons
Forms +2 charge
Less reactive than alkali metals
Reactive with oxygen
Can be very reactive with water

10. Transition Elements - d and f blocks
Two types: transition metals (d-block) and inner transition metals (f-block)
Number of valence electrons varies

11. Halogens – Group 17
Called “salt formers” – they can form compounds with almost all metals
Exist as solids (I and At), liquids (Br), and gases (F and Cl) at room temperature
7 valence electrons
Forms -1 charge
Very reactive with metals

12. Noble Gases – Group 18 Colorless gases Unreactive, or inert
8 valence electrons – a full outer energy level, except He which has only 2 valence electrons
Very stable

13. Nuclear Charge and Electron Shielding
Nuclear Charge – directly proportional to the number of protons in the nucleus.
Increases across a period due to increases in the number of protons in the nucleus without adding energy levels.
Decreases down a group due to additional energy levels
Electron Shielding – refers to the repulsion between negative core electrons and valence electrons and will change the effective nuclear charge.
Increases down a group due to the increases in energy levels.

14. Atomic Radius Atomic Radius – the size of a neutral atom
Increases as you go down a group because of the addition of energy levels.
Decreases as you go across a period as a result of the increasing nuclear charge that occurs by the addition of one more proton and one more electron.
The electrons in the higher energy levels are pulled closer to the nucleus making the atom smaller.

15. Atomic Radius

16. Ionic Radius Ionic radius – the size of an ion, a charged atom
A cation (positively charged atom)
becomes smaller when it loses electrons.
An anion (negatively charged atom) becomes larger when it gains electrons.

17. Ionic Radius

18. Ionization Energy
Ionization energy (first ionization energy) – the energy required to remove the first valence electron from an atom.
First ionization energy decreases down a group because as the atom becomes larger the nuclear charge has a smaller effect on the outer electrons.
First ionization energy increases across a period because the nuclear charge increases.

19. Ionization Energy

20. Electronegativity
Electronegativity – the ability of an atom’s nucleus to attract electrons from another atom when the atom is in a compound.
Electronegativity decreases down a group because as the atom becomes larger the nuclear charge has a smaller effect on the outer electrons.
Electronegativity increases as you move across a period as a result of increasing nuclear charge.

21. Electronegativity
INCREASES
DECREASES

22. Reactivity
Reactivity– the chemical reactivity of an element depends upon the relative attraction between valence electrons of one element and the nucleus of another.
Metals
Nonmetals
increases down a group
decreases across a period
decreases down a group
Increases across a period