chem

Chemical Reactions and Compounds

• Example of Chemical Reaction
  • Interaction of methane ($CH4$) with oxygen ($O2$)
  • Produces carbon dioxide ($CO2$) and water ($H2O$)
  • Observes the movement of atoms to form new compounds
  • Definition: Chemical reactions yield new compounds.

Types of Chemical Bonds

Ionic vs. Covalent Bonds

• Ionic Bonds:

  • Occur between metals and nonmetals
  • Involves the transfer of electrons

• Covalent Bonds:

  • Occur between nonmetals
  • Involves the sharing of electrons

Overview of the Periodic Table

• Metals:

  • Includes alkali, alkaline, and transition metals

• Metalloids:

  • Located along the zig-zag line of the periodic table
  • Can behave like metals or nonmetals depending on the reaction

• Nonmetals:

  • Found on the right side of the periodic table (in yellow, including hydrogen)
  • Bond with each other covalently

Examples of Bonding Scenarios

• Covalent Bonding Examples:

  • Polar bears sharing an ice cream cone: Both share equally in nonpolar covalent bonding
  • Polar bears and penguins sharing an ice cream cone: Represents polar covalent bonding, where one (polar bear) has more dominance in sharing the cone

• Ionic Bonding Example:

  • The penguin does not receive any of the ice cream cone, representing complete transfer of electrons

• Valence Electrons:

  • Sodium ($Na$) has one valence electron
  • Loses that electron easily to achieve stability, which refers to ionic bonding

Electronegativity

• Definition:
  • Electronegativity indicates how strongly an atom attracts electrons towards itself
• Electronegativity Trends:
  • Top right elements (like Fluorine) have the highest electronegativity
  • As we move away from Fluorine, electronegativity decreases

Bond Classification by Electronegativity Difference

• Nonpolar Covalent Bonds:

  • Difference of $0 - 0.4$
  • Example: Carbon ($C$, 2.5) and Hydrogen ($H$, 2.1), difference of 0.4: electrons are shared equally

• Polar Covalent Bonds:

  • Difference of $0.4 - 1.8$
  • Example: Carbon ($C$, 2.5) and Oxygen ($O$, 3.5), difference of 1: electrons are shared unequally, with oxygen pulling more

• Ionic Bonds:

  • Difference greater than $1.8$
  • Example: Sodium ($Na$, 0.9) and Chlorine ($Cl$, 3.0), difference of 2.1: complete electron transfer

Ionic Charges and Compound Formation

• Cations:

  • Carry a positive charge (fewer electrons than protons)
  • Example: Sodium ($Na^+$) loses one electron and carries a +1 charge

• Anions:

  • Carry a negative charge (more electrons than protons)
  • Example: Chlorine ($Cl^-$) gains one electron and carries a -1 charge

Predicting Ion Charges

• Sodium ($Na$)

  • 1 valence electron (alkali metal)
  • Loses it easily, resulting in a +1 charge

• Magnesium ($Mg$)

  • 2 valence electrons
  • Loses two to become +2 charged

• Aluminum ($Al$)

  • 3 valence electrons
  • Loses three to become +3

• Carbon ($C$)

  • 4 valence electrons
  • Can either gain or lose, prefers covalent bonding

• Nitrogen ($N$)

  • 5 valence electrons
  • Gains three to become -3 charged

• Oxygen ($O$)

  • 6 valence electrons
  • Gains two to become -2 charged

• Fluorine ($F$)

  • 7 valence electrons
  • Gains one to become -1 charged

Charge Balancing in Ionic Compounds

• When forming ionic compounds, charges must balance out

• Example: Sodium Chloride ($NaCl$)

  • Sodium (+1) and Chlorine (-1) balance each other out: $NaCl$

• Example: Magnesium Bromide ($MgBr_2$)

  • Magnesium (+2) and Bromine (-1): need 2 bromine for every 1 magnesium

Polyatomic Ions and Naming Conventions

• Polyatomic Ions:
  • Groups of atoms that act like a single ion
  • Example: $SO_4^{2-}$ (sulfate ion)
• **Naming Conventions:
  • Compounds can end in -ate or -ite reflecting different oxygen numbers
  • More oxygens = -ate, fewer = -ite

Covalent Bonding

• Covalent bonds involve sharing electrons to achieve stable electron configurations (octet)

Covalent Bonding Examples

• Carbon and Hydrogen:

  • Carbon has 4 valence electrons.
  • Forms covalent bonds with 4 hydrogens ($H$) to achieve 8.

• Ammonium ($NH_4^+$):

  • Nitrogen has 5 valence electrons, needs 3 more electrons which are provided by 3 hydrogens.
  • Results in a positive charge due to one less electron for ammonia.

• Formal Charge:

  • A theoretical charge based on the number of valence electrons an atom brings versus how many it has in its bonded state.
  • Example: Nitrogen brings 5; if bonded with 4 electrons, it is a +1 charge due to missing one electron.