Chemical Bonding: Atoms, Molecules, and Compounds

• Chemical bonding is being analyzed in this lesson.
• Mrs. Sheryl M. Mallari RPh. prepared the lesson.

Atom

• It is the smallest unit of an element that retains the element's properties.
• Atoms combine to form molecules and compounds, their interactions govern chemical properties.

Molecules

• This is a group of two or more atoms chemically bonded.
• Examples include Oxygen ($O_2$), Water ($H_2O$), and Methane ($CH_4$).
• Diatomic molecules consist of two atoms of the same element bonded together ($O_2$, $N_2$, $H_2$).
• Polyatomic molecules consist of three or more atoms bonded together ($H_2O$, $CO_2$, $NH_3$).

Compound

• A substance composed of two or more different elements chemically bonded together in fixed proportions.
• All compounds are molecules, but not all molecules are compounds.
• H2 is a molecule, but not a compound because only H is present.
• H2O is both a molecule and a compound because H and O are present.

Chemical Bonding

• Interactions between atoms hold them together to form molecules or compounds.
• Compounds are divided into two broad categories: ionic and molecular.
• Ionic compounds form when electrons are transferred between atoms, resulting in ion formation.
• Molecular compounds form through sharing one or more pairs of electrons between two atoms.

Valence Electron

• Valence electrons reside in the outermost shell of an atom.
• There are 2 fundamental concepts to understanding ionic and covalent bonding models.
• Not all electrons in an atom participate in bonding.
• Certain electron arrangements are more stable than others, explained by the octet rule.
• Oxygen has an atomic number of 8, 2 valency, with 6 valence electrons.

Determination of Valence Electron

• To determine the number of valence electrons in atoms of Mg, Si, As.

Covalent Bonding

• A chemical bond results from two nuclei attracting the same shared electrons.
• Considers two hydrogen atoms, each with a single electron brought together.

Covalent Bond Lewis Structure

• It is an electron dot structure.
• Valence shell electrons of an atom are represented as dots, known as Lewis dot symbols or electron dot symbols.
• Representative elements of the same group in the periodic table has the same number of valence electrons.
• The number of valence electrons for typical elements corresponds to the Roman numeral periodic-table group number.
• The maximum number of valence electrons for any element is eight.

Covalent Bond Kekule Structure

• Line bond structure
• Each shared electron is represented by a line between the atom symbols.
• H has one bond, C has four bonds, N has three bonds and one unshared pair of electrons.
• O has two bonds and two unshared pair of electrons, F, Cl, Br, and I have one bond and three unshared pairs of electrons.

Lone-pair electrons or non-bonding electrons

• Pair of valence electrons that are not used for bonding.
• Multiple Bonds: When two atoms share more than one pair of electrons, it forms a double or triple bond.

Identifying Formal Charges

• Formal Charge is associated with any atom that does not exhibit the appropriate number of valence electrons.
• Determine the number of valence electrons, then determine whether the atom exhibits the appropriate number of electrons.
• Formal Charge = V - [N + (B/2)], with V = # of valence electrons in the free atom, N = # of non-bonding (lone pair) electrons, B = # of bonding (shared) electrons
• Includes formulas to determine the formal charge of:
  • Oxygen in Water ($H_2O$)
  • Nitrogen in Ammonia ($NH_3$)
  • Carbon in Methane ($CH_4$)
  • Oxygen in Carbon Dioxide ($CO_20$)
  • Carbon in Carbon Monoxide (CO)

Induction and Polar Covalent Bonds

• INDUCTION refers to the phenomenon where electron density within a molecule is redistributed due to varying electronegativities.
• POLAR COVALENT BONDS occur where electrons are shared unequally between two atoms due to electronegativity differences.

Electronegativity Values

• If the difference is less than 0.5, it's a non-polar covalent bond with equally shared electrons.
• A difference of 0.5-1.7 indicates a polar covalent bond, not equally shared, and induction occurs.
• If greater than 1.7, it's an ionic bond, and electrons are not shared (e.g., NaOH).

Hybridization

• Process where one s orbital and three 'p' orbitals from the same atomic shell mix to form four new equivalent orbitals.
• Methane ($CH_4$) is a prominent example of $sp^3$ hybridization.
• Each carbon atom creates four sigma bonds with four hydrogen atoms in methane.

$sp^3$ Hybridization

• Hybridization occurs when an atom mixes one s orbital and three p orbitals to form four equivalent $sp^3$ hybrid orbitals.
• The sum of attached atoms + No. Lone Pairs= 4
• Solid line represents bonds in the plane of the page.
• Heavy Wedged Line represents bonds coming out of the page, toward the viewer.
• Dashed Line represents bonds preceding back behind the page, away from the viewer
• The indicated atoms can be $sp^3$ hybridized

$sp^2$ Hybridization

• $sp^2$ Hybridization: Sum of attached atoms + No. Lone Pairs = 3.

sp Hybridization

• Sum of attached atoms + No. Lone Pairs= 2.

Shortcut for Determining Hybridization

• To determine the hybridization of any atom: count the number of attached atoms then consult the table.
• If atoms + lone pairs = 4, hybridization is $sp^3$
• If atoms + lone pairs = 3, hybridization is $sp^2$
• If atoms + lone pairs = 2, hybridization is sp
• When counting, don't count the bonds - do count the # of attached atoms

Exceptions noted:

• Atoms with lone pairs next to pi bonds
• For Oxygen: 2 atom + 2 lone pair
• For Nitrogen: 3 atom + 1 lone pair
• For sp³ hybridizations you need hybridization from sp³ to sp²

VSEPR Theory: Predicting Geometry

• Steric Number: total number of sigma bonds and lone pairs.
• Indicates the number of electron pairs (bonding and non-bonding) that are repelling each other.

Drawing Chemical Structures

• It is a way of representing molecules where atoms are listed in a sequence that shows their bonding, without explicitly drawing all the bonds.
• Grouping of Atoms: Atoms directly bonded to each other are written together.
  • Ethane: $CH_3CH_3$
  • Propane: $CH_3CH_2CH_3$
• Parentheses for Branches or Groups: Substituents or repeating groups are enclosed in parentheses.
  • Isopropanol: $(CH_3)_2CHOH$
  • Hydrogens Implied: Hydrogen atoms are often implied based on carbon's valency.
  • For example, $CH_4$ means one carbon is bonded to four hydrogens.
• Double and Triple Bonds: Double or triple bonds are indicated using “=” or “≡”.
  • Ethene: $CH_2=CH_2$
  • Ethyne: $HC≡CH$