Science 10 Chemistry: The Periodic Table, Energy Level Diagrams, and Ions

The Periodic Table: History and Organization
Mendeleev's Contribution
- Dmitri Mendeleev organized elements according to recurring or "periodic" properties.
- He utilized a system that left gaps for elements that were "missing" or yet to be discovered at the time.
- Prior versions of the periodic table were organized based on atomic weights.
- Mendeleev arranged elements according to their atomic number ( $\text{atomic \#}$ ).
- Instruction noted in the text: "Do Not Copy".
Early Attempts and Specialized Groupings
- Early versions of the Periodic Table featured different layouts, including identified groups for:
  - Superactinides: Elements beyond the standard actinide series.
  - Transition Metals: Located in the central block of the table.
  - Lanthanides & Actinides: Positioned at the bottom of the table to maintain organizational flow.

Definition: Periods are the horizontal rows that run across the periodic table.
Labeling: There are $7$ primary periods labeled from top to bottom.
Function: The period number represents the total number of electron shells (energy levels) present in the atoms of those elements.
Specific Elements Mentioned in Rows:
- Period 1: Hydrogen ( $H$ , $1.01$ ) and Helium ( $He$ , $4.00$ ).
- Period 2: Lithium ( $Li$ , $6.94$ ), Beryllium ( $Be$ , $9.01$ ), Boron ( $B$ , $10.81$ ), Carbon ( $C$ , $12.01$ ), Nitrogen ( $N$ , $14.01$ ), Oxygen ( $O$ , $16.00$ ), Fluorine ( $F$ , $19.00$ ), and Neon ( $Ne$ , $20.18$ ).

Definition: Groups or Families are the vertical columns that run down the periodic table.
Labeling: These are labeled from $1$ to $18$ .
Significance:
- A family may consist of one single column or several columns grouped together.
- Elements within the same group share similar chemical properties.
- The group number determines the number of valence electrons (electrons in the outer shell).
Named Groups and Specific Properties:
- Alkali Metals: Group $1$ . These are soft metals that react very easily with $H_2O$ to form a base (alkali).
- Alkaline Earth Metals: Group $2$ . These are light, reactive metals that form oxide coatings.
- Halogens: Group $17$ . These elements react with metals to form "salts."
- Noble Gases: Group $18$ . These are unreactive or "inert" gases.
- Transition Elements: Groups $3$ through $12$ . These exhibit a wide range of physical and chemical properties.
- Lanthanides: Atomic numbers $58$ to $71$ .
- Actinides: Atomic numbers $90$ to $103$ . These are mostly radioactive or synthetically produced.
- Rare Earth Elements: A collective term for the Lanthanides and Actinides.
- Transuranic Elements: These are the synthetic elements found beyond Uranium ( $U$ ) on the periodic table.

The "Staircase" Line:
- This line acts as a boundary separating metals from non-metals.
- Metalloids: These are elements located along the staircase line that exhibit properties of both metals and non-metals.
Properties Comparison:
- Metals:
  - State: Solids at room temperature (with the exception of Mercury, $Hg$ ).
  - Appearance: Shiny (lustrous).
  - Conductivity: Good conductors of both heat and electricity.
  - Malleability: Malleable (can be hammered into sheets).
  - Ductility: Ductile (can be drawn into wires).
- Non-Metals:
  - State: Can be gas, liquid, or solid at room temperature.
  - Appearance: Not very shiny (dull).
  - Conductivity: Do not conduct heat or electricity well.
  - Malleability: Brittle (tend to break or shatter when stressed).
  - Ductility: Not ductile.

Bohr Model of e-:
- Electrons ( $e^-$ ) orbit the nucleus in fixed energy levels.
- The period number equals the number of energy levels in its atoms.
Energy Level Capacity:
- Each energy level has a maximum capacity for electrons:
  - $1^{st}$ Energy Level: Holds up to $2$ electrons.
  - $2^{nd}$ Energy Level: Holds up to $8$ electrons.
  - $3^{rd}$ Energy Level: Holds up to $8$ electrons.
- Once an energy level is full, electrons begin filling the next available level.
Valence Electrons and Stability:
- Valence electrons ( $e^-$ ) are the electrons located in the outermost shell of an atom.
- Elements are most stable when their outer level is full.
- This stability is the reason why atoms of different elements undergo chemical reactions to combine.

Energy Level Diagrams:
- A model representing the atom with the nucleus at the center and electrons drawn in orbits around it.
- Elements in the same group have the same number of valence electrons.
- Specific elements targeted for practice: $H$ , $He$ , $Li$ , $Be$ , $P$ , and $Ar$ .
Lewis Dot Diagrams:
- These focus exclusively on the valence electrons.
- The diagram consists of the element's chemical symbol surrounded by dots representing the valence electrons.
- Example elements to draw: $Li$ , $Be$ , $P$ , $Ar$ , and $He$ .

Neutral Atoms:
- In a neutral atom, the number of protons ( $p^+$ ) is equal to the number of electrons ( $e^-$ ).
Ionization Process:
- Elements want to lose or gain electrons in their valence shell to achieve a full shell, known as a "stable octet."
- By reaching a full shell, they achieve an electron configuration similar to the nearest noble gas.
- If the number of electrons does not equal the number of protons, the atom carries a charge and becomes an ION.
- Representation: Ions are represented by the element symbol and the ion charge (e.g., $Na^+$ , $N^{3-}$ , $Mg^{2+}$ , $Br^-$ ).
Metal Ions (Cations):
- Metals tend to lose electrons to become positive.
- These positive ions are called CATIONS.
- The name of a metal ion is the same as the name of the element.
Non-Metal Ions (Anions):
- Non-metals tend to gain electrons to become negative.
- These negative ions are called ANIONS.
- The names of non-metal ions are modified to end in the suffix "-ide" (e.g., Chlorine becomes chloride).

Example C3: Magnesium Ion ( $Mg^{2+}$ ):
- Atom vs. Ion: A neutral Magnesium atom has $12$ protons. The magnesium ion ( $Mg^{2+}$ ) loses its $2$ valence electrons to achieve stability.
- Diagram for $Mg^{2+ nuestros}$ :
  - Nucleus: $12 p^+$ , $12 n^0$ .
  - Energy Levels: $2e^-$ in the first level, $8e^-$ in the second.
- Nearest Noble Gas: The magnesium ion has the same number of electrons ( $10$ ) as Neon ( $Ne$ ).
- Neon ( $Ne$ ) Diagram: $10 p^+$ , $10 n^0$ ; $2e^-$ , $8e^-$ .
Example C4: Sulfide Ion ( $S^{2-}$ ):
- The sulfide ion is the ionic form of Sulfur.
- Diagram for Sulfide ( $S^{2-}$ ):
  - Nucleus: $16 p^+$ , $16 n^0$ .
  - Energy Levels: $2e^-$ (first level), $8e^-$ (second level), $8e^-$ (third level).
- Nearest Noble Gas: The sulfide ion has the same number of electrons ( $18$ ) as Argon ( $Ar$ ).
- Argon ( $Ar$ ) Diagram: $18 p^+$ , $22 n^0$ ; $2e^-$ , 8e^-$, 8e^-$$.