In 1865 he organized the first 16 elements by atomic mass in two rows.
Law of Octaves
When John Newlands noticed similar properties in the elements.
In 1870 he organized the then known 63 elements by atomic mass, and grouped them by similar properties. He left open spaces and made predictions based on properties of elements around the open spaces.
In 1910 he "hit" known elements with x-rays and derived a relationship between x-ray frequency and the number of protons. Organized the elements by atomic number, decreased troubles with Mandeleev's table, and made the modern periodic table organization.
Across/Rows. Gives the # of energy levels.
Up and Down/Columns. Gives the number of valence electrons.
Effective Nuclear Charge
The net positive charge experienced by an electron in a multielectron atom (depends on # of protons, shielding effect of other electrons, distrance from nucleus)
Inner electrons block the attraction of the protons to the outer electrons.
An atom's size is determined by how far away the electrons are from the nucleus. The distance between the nuclei of two identical atoms, divided by two.
Atomic Radii Left to Right
Electrons are all in the same energy level. Adding more protons cause the electrons to move close to the nucleus.
Atomic Radii Top to Bottom
Adding another energy level from one period to the next. Inner electrons shield outer electrons from the full charge of the nucleus. (Shielding Effect) Noble Gases are an exception.
Radii of Ions (Cations)
Positive ions are smaller than their parent atoms.
Radii of Ions (Anions)
Negative ions are larger than their parent atoms.
A series with the same number of electrons and same electron configuration. Example: O-2, F-, and Na+ are an isoelectronic series because you changed the number of electrons causing them to have the same configuration.
The energy required to remove an electron from an atom. An atom is neutral, and to become a positive ion...lose an electron... there must be some energy put into it.
Ionization Energy Left to Right
As we go to the right, we need more energy to remove an electron because atoms want to gain electrons, not lose them.
Ionization Energy Top to Bottom
Less energy is required to remove an electron as we go down a column. Shielding effect: inner electrons block the attraction of the protons to the outer electrons. Therefore, if they are held less tightly, they are more easily extracted.
How strongly an atom attracts electrons of other atoms to itself. Basically the same as ionization.
Electronegativity Left to Right
Gets bigger. Same as Ionization Energy.
Electronegativity Top to Bottom
Gets smaller. Same as Ionization Energy.
Melting Points Metals
The melting point for metals generally decreases as you go down a group.
Melting Points Non-Metals
The melting point for non-metals generally increases as you go down a group.
These are the sub atomic particle chemical properties mostly depend on.
Oxides of Period 3 (Na, Mg, Al: Metals)
Oxides are ionic and conduct electricity in the liquid phase. Na and Mg oxides dissolve in water forming basic solutions. Al and Si oxides are acidic or basic.
Oxides of Period 3 (P, S, Cl: Nonmetals)
Oxides are covalent and they react with water forming. When you see H think acid, when you see OH think base.
Ions dissociate in solvent. (Aqueous)
Ions remain bonded insolvent.
2 compound exchange cations. General Form: AB + CD goes to AD + CB
Alkali Metals Reactivity
As we go down the group, size increases, ionization energy decreases and reactivity increases. React with water.
As we go down the group they are less reactive. Bigger atom and weaker attraction between nucleus and outer electrons so it is harder to gain electron.