Module 5 - Periodic Table & Periodic Law

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%%Antoine Lavoisier (1743-1794)%% - compiled a list of all the elements that were known at the time.

%%John Newlands%% (1837-1898) - Proposed an organizational scheme for the elements.

• Law of Octaves - When elements are arranged by increasing atomic mass, their properties repeat every eighth element (this did not work for every element).

%%Lothar Meyer%% (1830-1895) - Demonstrated a connection between atomic mass and the properties of an element and arranged elements in order of increasing atomic mass.

%%Dmitri Mendeleev%% (1834-1907) - Demonstrated a connection between atomic mass and the properties of an element, arranged elements in order of increasing atomic mass, and predicted the existence and properties of undiscovered elements (not completely correct when new elements were discovered).

%%Henry Moseley%% (1887-1915) - Discovered that atoms contain a unique number of protons called the atomic number and arranged elements in order of increasing atomic number, which resulted in a periodic pattern of properties.

$**Periodic Law**$ - periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number.

$**Groups/Families**$ - Columns (vertical)

• Elements with similar physical and chemical properties

$**Periods**$ - Rows (horizontal)

$**Representative Elements/Main Group**$ - Groups 1, 2, and 13-18

• Wide range of physical and chemical properties

$**Transition Elements/Metals**$ - Groups 3-12

• Metals, nonmetals, and metalloids

$**Metals**$ - Luster- shiny, reflects light, good conductors of heat and electricity, malleable, ductile,

• Groups include- alkali metals, alkaline earth metals, transition metals, inner transition metals, and some elements in the lower left corner of the p Block on the periodic table.

$**Alkali Metals**$ - Group 1 (except for hydrogen)

• Extremely reactive metals

$**Alkaline Earth Metals**$ - Group 2

• Highly reactive

$**Inner Transition Metals**$ - Lanthanide and actinide series

$**Nonmetals**$ - Upper-right side

• Gases or brittle, dull solids, poor conductors

$**Halogens**$ - Group 17

• Extremely reactive nonmetals

$**Noble Gases**$ - Group 18

• extremely unreactive gases

$**Metalloids**$ - On the zigzag

• Possess properties of both metals and nonmetals

$**Valence electrons**$- electrons located in the outermost energy level of an atom

• Each of the group 1 electrons has similar chemical properties because they have the same number of valence electrons.

$**S-Block Elements**$ - Groups 1 and 2

• Partially filled orbitals

$**P-Block Elements**$ - Groups 13-18

• Filled or partially filled orbitals
• Group 18, the noble gases, are unique because they are so stable

$**D-Block Elements**$ - Transition metals (largest of the blocks)

• Usually, a filled outermost orbital of energy level n or partially filled orbitals of energy level n-1

$**F-Block Elements**$ - Inner transition elements,

• Filled or partially filled outermost s orbital, and a filled or partially filled 4f and 5f orbitals

$**Atomic Size**$ - how closely an atom lies to a neighboring atom

• Varies from substance to substance

$**Atomic Radius Metals**$- Half the distance between adjacent nuclei in a crystal

^^Atomic Radius^^ - Decrease left to right across a period and increase as you move down a group

^^Ionic Radius Positive & Negative Ions^^ - Decreases across a period and increases down a group

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$**Ion**$ - An atom or a bonded group of atoms that has a positive or negative charge.

$**Positive Ion**$ - An atom that has lost one or more electrons.

• Are smaller than their atoms

$**Negative Ion**$ - An atom that has gained one or more electrons

• Are larger than their atoms

When atoms lose electrons and form positively charged ions, they always become smaller.

• The electron lost will always be a valence electron
• The loss of a valence electron can leave a completely empty outer orbital (which results in a smaller radius)

$**Ionization Energy**$ - The energy required to remove an electron from a gaseous atom

• An indication of how strongly an atom’s nucleus holds onto its valence electrons (high ionization energy = a strong hold) (low ionization energy = less likely to form positive ions)

^^Ionization Energy^^ - Increases left to right across a period and decreases down a group

$**Octet Rule**$ - States that atoms tend to gain, lose, or share electrons in order to acquire a full set of eight valence electrons

$**Electronegativity**$ - Indicates the relative ability of its atoms to attract electrons in a chemical bond.

^^Electronegativity^^ - Generally increases left to right across a period and decreases down a group