Quiz 1 cards

Pieces of Eight: Introduction to Lewis Dot Symbols and the Octet Rule

Lewis Dot Structure

➡Also known as Lewis dot diagrams or electron dot diagrams.

➡Represents valence electrons of an atom.

➡Visualizes valence electrons as lone pairs or bonds.

➡Number of dots = number of valence electrons in the atom.

➡Maximum of two dots per side.

• Electron Stucture of the atom

  

➡Start dot placement at 12 o’clock, move clockwise (3, 6, 9, back to 12)

➡Element symbol represents nucleus and core electrons.

➡Maximum Lewis dots:

• Noble gases (except He)

➡Highest ionization energy:

• Energy needed to remove one electron.

• Noble gases 

➡Why Noble Gases are stable:

• High ionization energy makes them non-reactive.

Octet Rule

➡Gilbert Lewis: Proponent of the Lewis Dot Structure and Octet Rule.

➡Applies to representative elements.

➡Atoms prefer to have eight electrons in their valence shell.

➡Fewer than eight electrons lead to reactions to form stable compounds.

➡A complete octet means all orbitals are full and stable.

➡Represents valence electrons as dots in Lewis diagrams.

➡Atoms seek their most stable state with a full valence shell of 8.

Exception to Octet Rule

➡Duet rule (Duplet)

• H and He are stable with two valence electrons.

Explore types of bonding and their properties

➡Bohr Model

• Center - Nucleus

• Orbits - energy levels

• Inner and Outer Electrons 

➡Valence Electrons

• Electrons in the outermost shell or highest energy level

• Group number = number of valence electrons 

Chemical Bonding

➡Elements combine through chemical bonding 

• Forces that hold atoms together

➡The process by which atoms combine to form compounds.

➡Bonds hold atoms together allowing them to achieve stability, often by fulfilling the octet rule (8 electrons in the valence shell).

➡Bonds form when valence electrons interact. 

Exceptions:

➡Noble Gases (He, Ne, Ar, Kr, Xe, Rn) occurs in air as atoms.

➡O, N, and S occurs in elemental form as molecules.

➡C occurs as a pure deposit (Ex. Coal).

➡Some metals like Cu, Ag, Au, and Pt may occur uncombined with other elements.

Why form bonds?

➡Atoms with fewer than eight electrons are unstable and tend to react with others.

➡They form bonds to reach the most stable (lowest energy) state possible.

Ionic Bonding

➡Occurs when electrons are transferred from one atom to another, forming charged ions (cations and anions). This bond typically forms between metals and nonmetals.

➡Ionic bonds usually form solid objects.

➡Formation of Ionic Bonds:

• A metal atom loses electrons, forming a cation (positively charged ion).

• A nonmetal atom gains electrons, forming an anion (negatively charged ion).

➡Example: Sodium Chloride (NaCl)

• Sodium (Na) loses one electron: Na ► Na⁺ + e⁻

• Chlorine (Cl) gains one electron: Cl + e⁻ ►Cl⁻

• The oppositely charged ions attract to form NaCi: Na⁺ + Cl⁻ ► NaCl

➡Characteristics:

• High Melting and Boiling Points: Strong ionic bonds require significant energy to break.

• Crystalline Structure: Form regular crystal lattice structures for stability.

• Solubility in Water: Many are soluble, as water stabilizes the ions.

• Electrical Conductivity:

  • Solid State: Non-conductive due to fixed ions.

  • Molten/Aqueous State: Conductive due to mobile ions.

➡The charges on the anion and cation correspond to the number of electrons donated or received.

➡Why is it called ionic:

• Generates two oppositely charged ions: cation (metal) and anion (nonmetal).

• Because they are opposite charges, opposite charges attract one another by electrostatic forces.

➡How do they achieve full octet:

• In ionic bonds, the net charge of the compound must be zero (the positive charges of the cations must equal the negative charges of the anions.)

Sodium Ion Charge: (1 atom of Na) x (1+ charge) = +1

Chlorine Ion Charge: (1 atom of Cl) x (1- charge) = -1

Algebraic Sum of the total charge: + 1 - 1 = 0

Steps on how to draw Lewis Structures of Ionic Compounds

➡Step 1: Draw the Lewis dot symbol of each atom

➡Step 2: Draw arrow/s to show transfer of electrons

➡Step 3: Show formation of ions

➡Step 4: Draw the Lewis structure for Ionic Compounds

Note: 

➡Charges of each ion indicates the number of electrons donated (+) or received (-). Anions are always enclosed in brackets with its charge outside the bracket.

➡For more than 1 atom used, write the numerical coefficient before the Lewis structure of the specific ion.

Covalent Bonding

➡Occurs when two nonmetal atoms share pairs of electrons to achieve stable electron configurations. This type of bond is typically formed between atoms with similar electronegativities.

➡Sharing of electrons between atoms

➡Formation of Ionic Bonds:

• Each atom contributes one or more valence electrons to form shared pairs.

• This allows each atom to achieve a full outer shell (often following the octet rule).

➡Types of Covalent Bonds:

• Single Bond:

  • One pair of shared electrons.

  • Example: H—H (hydrogen gas, H2)

• Double Bond:

  • Two pairs of shared electrons.

  • Example: O=O (oxygen gas, O2)

• Triple Bond:

  • Three pairs of shared electrons.

  • Example: N≡N (nitrogen gas, N2)

➡Characteristics:

• Lower Melting and Boiling Points: Compared to ionic compounds due to weaker intermolecular forces.

• Varied Physical States: Can be gases, liquids, or solids at room temperature.

• Solubility: Varies; polar covalent compounds tend to be soluble, while nonpolar compounds are generally insoluble in water.

• Electrical Conductivity: Do not conduct electricity in any state due to lack of charged particles.

➡Why share:

• Nonmetals have very high IE - too hard to remove an electron.

Electronegativity 

➡The tendency of an atom to attract shared electrons in a chemical bond.

➡It plays a key role in determining the bond type and bond polarity between atoms.

➡Bond Types:

• Nonpolar Covalent Bonds:

  • Occur when atoms have similar electronegativities.

  • Electrons are shared equally.

  • Example: H2, O2, Cl2

• Polar Covalent Bonds:

  • Occur when atoms have moderate electronegativity differences.

  • Electrons are shared unequally, creating partial positive and partial negative charges.

  • Example: H2O, HCl

• Ionic Bonds:

  • Occur when atoms have a large difference in electronegativity.

  • Electrons are transferred, not shared, forming ions.

  • Example: NaCl

➡Because both atoms have close electronegativities they have the same affinity (“love”) for electrons.

➡Neither atoms want to donate them completely.

➡Thus, they share electrons to achieve octet.

➡They will share their electrons with each other by an overlap of their outer energy levels.

Shared Pair

➡Known as a single (covalent) bond and is represented by a single line like a dash.

Lewis Structure of Covalent Compounds

➡Since covalent compounds are different from ionic ones, we have different steps in drawing the Lewis structures for Covalent compounds.

➡Let us consider CF4.

➡Step 1: Get the sum of the total number of valence electrons of all atoms.

C = 4 v.e. X 1 = 4 e-

F = 7 v.e. X 4 = 28 e-

              Total = 32 e-

➡Step 2: Place the atoms relative to each other.

• The less electronegative (except H) element is the center.

• Others are surrounding the central atom.

➡Step 3: Use a pair of electrons in the form of bond [-].

• Shared pair = 2 e-

Total Valence Electrons = 32 e-

               Bonding pairs = 8 e-

                    Remaining = 24 e-

➡Step 4: Distribute the remaining electrons in pairs so that each atom has eight electrons (or equal to two for H).