GOB Chemistry Study Notes: Covalent Bonding, Molecular Structure, and Chemical Reactions

Covalent Bonding and Molecular Structure

Difference Between Ionic and Covalent Bonds

Chemical bonds are the forces that hold atoms together in compounds. The two main types are ionic bonds and covalent bonds.

• Ionic bond: Transfer of electrons from a metal to a nonmetal, resulting in the formation of oppositely charged ions.

• Covalent bond: Sharing of electrons between two nonmetals, forming a molecule.

Example: NaCl (ionic), H2O (covalent)

The 7 Diatomic Elements

Certain elements exist naturally as diatomic molecules, meaning two atoms of the same element are bonded together.

• Diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2

Example: O2 in the atmosphere

Forces in Covalent Bonds

Covalent bonds involve both attractive and repulsive forces between atoms.

• Attractive force: Between the nuclei and shared electrons.

• Repulsive force: Between the two nuclei and between electrons.

Example: The balance of these forces determines bond length and strength.

Predicting Number of Covalent Bonds

The number of covalent bonds an atom forms depends on its valence electrons and the need to complete an octet.

• Atoms share electrons to achieve a stable octet (8 valence electrons).

• Use the number of valence electrons to predict how many bonds are needed.

Example: Oxygen (6 valence electrons) forms 2 covalent bonds in H2O.

Single, Double, and Triple Bonds

Covalent bonds can involve sharing one, two, or three pairs of electrons.

• Single bond: 2 electrons shared ($A-B$)

• Double bond: 4 electrons shared ($A=B$)

• Triple bond: 6 electrons shared ($A\equiv B$)

Example: O2 has a double bond; N2 has a triple bond.

Physical Properties of Covalent Compounds

Covalent compounds have distinct physical properties due to their molecular structure.

• Low melting and boiling points

• Do not conduct electricity in water

• Often soft or brittle solids

Example: Sugar (C12H22O11) is a covalent compound.

Drawing Lewis Dot Structures

Lewis dot structures represent the arrangement of valence electrons in molecules.

• Use valence electrons to draw structures with shared (bonding) and lone pairs.

• Helps predict molecular shape and reactivity.

Example: Lewis structure of H2O shows two lone pairs on oxygen.

Predicting Molecular Shape (VSEPR Model)

The Valence Shell Electron Pair Repulsion (VSEPR) model predicts molecular geometry based on the number of charge clouds (bonds and lone pairs) around the central atom.

• 2 charge clouds: linear (e.g., CO2)

• 3 charge clouds: trigonal planar or bent (e.g., SO2)

• 4 charge clouds: tetrahedral, trigonal pyramidal, or bent (e.g., CH4, NH3, H2O)

Example: H2O is bent due to two lone pairs on oxygen.

Electronegativity Trends

Electronegativity is the ability of an atom to attract electrons in a bond.

• Increases across a period (left to right) and up a group (bottom to top) in the periodic table.

• Fluorine is the most electronegative element.

Example: In HCl, Cl is more electronegative than H.

Partial Charges in Polar Bonds

When atoms in a bond have different electronegativities, partial charges develop.

• More electronegative atom: $\delta^-$ (partial negative)

• Less electronegative atom: $\delta^+$ (partial positive)

Example: In H2O, O is $\delta^-$, H is $\delta^+$.

Bond Type by Electronegativity Difference

The difference in electronegativity between atoms determines bond type.

Electronegativity Difference	Bond Type
< 0.4	Nonpolar covalent
0.4 – 1.7	Polar covalent
> 1.7	Ionic

Example: H–Cl (polar covalent), Na–Cl (ionic)

Polarity of Molecules

A molecule is polar if it has polar bonds and an uneven charge distribution due to its shape.

• CO2: Nonpolar (linear, charges cancel)

• H2O: Polar (bent, charges do not cancel)

• CH4: Nonpolar (tetrahedral, charges cancel)

Example: Water is a polar molecule.

Naming Binary Molecular Compounds

Binary molecular compounds are named using prefixes to indicate the number of atoms. The second element ends in "-ide"; never use "mono-" for the first element.

Prefix	Number
mono-	1
di-	2
tri-	3
tetra-	4
penta-	5

Example: CO2 is carbon dioxide; N2O is dinitrogen monoxide.

Classification and Balancing of Chemical Reactions

Reactants and Products; Law of Conservation of Mass

Chemical reactions involve the transformation of reactants into products. The Law of Conservation of Mass states that mass is conserved; atoms are neither created nor destroyed.

• Reactants: Substances before the reaction

• Products: Substances after the reaction

Example: $\mathrm{CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O}$

Balancing Chemical Equations

Balancing equations ensures the same number of each type of atom on both sides of the reaction.

• Adjust coefficients, not subscripts.

• Check each element for balance.

Example: $\mathrm{2H_2 + O_2 \rightarrow 2H_2O}$

Types of Chemical Reactions

Chemical reactions are classified by the changes that occur.

• Precipitation: Formation of an insoluble solid (precipitate) from solution.

• Acid-base: Reaction of acid and base to form salt and water.

• Redox: Transfer of electrons; oxidation and reduction occur.

Examples:

• Precipitation: $\mathrm{NaCl + AgNO_3 \rightarrow AgCl(s) + NaNO_3}$

• Acid-base: $\mathrm{HCl + NaOH \rightarrow NaCl + H_2O}$

• Redox: $\mathrm{Zn + CuSO_4 \rightarrow ZnSO_4 + Cu}$

Solubility Rules & Precipitation Products

Solubility rules help predict whether a compound will dissolve in water or form a precipitate.

• Use rules to determine if a solid forms in a reaction.

Example: $\mathrm{BaCl_2 + Na_2SO_4 \rightarrow BaSO_4(s) + 2NaCl}$

Acid-Base Reactions

Acid-base reactions produce a salt and water.

• General equation: $\mathrm{\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}}$

Example: $\mathrm{HNO_3 + KOH \rightarrow KNO_3 + H_2O}$

Identifying Acids and Bases

Acids and bases are defined by the ions they produce in water.

• Acids: Produce H+ ions

• Bases: Produce OH- ions

Example: HCl (acid), NaOH (base)

Assigning Oxidation Numbers

Oxidation numbers are used to keep track of electron transfer in redox reactions.

• Set of rules assigns charges to atoms in compounds.

Example: In H2O, H is +1, O is -2.

Redox Reactions: Oxidation & Reduction

Redox reactions involve the transfer of electrons between species.

• Oxidation: Loss of electrons

• Reduction: Gain of electrons

• Oxidizing agent: Gets reduced

• Reducing agent: Gets oxidized

Example: $\mathrm{Zn \rightarrow Zn^{2+} + 2e^-}$ (oxidation), $\mathrm{Cu^{2+} + 2e^- \rightarrow Cu}$ (reduction)

Molecular, Ionic, and Net Ionic Equations

Chemical reactions in solution can be represented in three ways:

• Molecular equation: Shows all reactants and products as compounds.

• Ionic equation: Shows all soluble substances as ions.

• Net ionic equation: Shows only the species that change during the reaction.

Example:

• Molecular: $\mathrm{NaCl(aq) + AgNO_3(aq) \rightarrow AgCl(s) + NaNO_3(aq)}$

• Net ionic: $\mathrm{Ag^+(aq) + Cl^-(aq) \rightarrow AgCl(s)}$

Additional info: Academic context and examples have been expanded for clarity and completeness. Tables have been recreated for bond type and naming prefixes. All equations are provided in LaTeX format as required.