Bonding: General Concepts – Study Notes

BONDING: GENERAL CONCEPTS

Covalent Bonds

Chemical bonds are the forces that hold atoms together in compounds. Covalent bonds are a type of chemical bond formed when atoms share electrons, resulting in the formation of molecules.

• Chemical bond: The force that holds two or more atoms together.

• Covalent bond: A bond formed by the sharing of electrons between atoms, typically between non-metals.

• Molecule: A collection of atoms held together by covalent bonds.

• Representations of molecules:

  • Chemical formula: Shows the types and numbers of atoms (e.g., CO2).

  • Structural formula: Shows how atoms are connected (e.g., H—O—H).

  • Condensed formula: A compact way to show structure (e.g., CH3CH2CHBrCH3).

Representation of Molecular Structures

Molecular structures can be visualized using different models to better understand their geometry and properties.

• Space-filling model: Indicates the relative sizes of atoms and their orientation in the molecule.

• Ball-and-stick model: Uses spheres (atoms) and rods (bonds) to show three-dimensional structure.

Attributes of Molecules

Molecules have various physical and chemical properties that can be measured and compared.

• Melting point

• Boiling point

• Electrical and thermal conductivity

• Solubility

• Electric charge

• Bond energy: The amount of energy required to break a bond between two atoms.

Types of Bonding

Ionic Bonding

Ionic bonding occurs between atoms that transfer electrons, typically between metals and non-metals.

• Ionic bond: Formed when one atom donates electrons to another, resulting in oppositely charged ions that attract each other.

• Ionic compounds: Formed when metals react with non-metals.

Covalent Bonding

Covalent bonding occurs when atoms share electrons, usually between non-metals.

• Non-metals tend to form covalent bonds with each other.

• Hydrogen can also form covalent bonds by sharing electrons.

Bonding in the H2 Molecule

The hydrogen molecule (H2) is a classic example of covalent bonding, where two hydrogen atoms share electrons.

• Electrons are shared in the space between the two nuclei.

• This sharing leads to increased stability compared to individual hydrogen atoms.

• The attractive forces between the nuclei and shared electrons lower the potential energy of the system.

Electronegativity

Definition and Effects

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond.

• Unequal sharing of electrons leads to charge separation (partial positive and negative charges) in a bond.

• This can result in polar molecules with distinct ends (dipoles).

Effect of Electric Field on Polar Molecules

Polar molecules, such as hydrogen fluoride (HF), align in an electric field due to their partial charges.

• In an electric field, the positive and negative ends of the molecules orient toward opposite charges.

Pauling's Method for Electronegativity

Linus Pauling developed a method to assign electronegativity values based on bond energies.

• For a molecule HX, compare the measured H—X bond energy to the expected value.

Pauling's equation:

$\text{Expected H—X bond energy} = \frac{\text{H—H bond energy} + \text{X—X bond energy}}{2}$

Trends in Electronegativity

• Electronegativity generally increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Fluorine is the most electronegative element.

Example: Bond Polarity

• Order the following bonds by increasing polarity: H—H, O—H, S—H, F—H.

• Answer: H—H (least polar) < S—H < O—H < F—H (most polar)

Electron Configuration and Compound Formation

Electron Sharing and Noble Gas Configurations

Atoms form bonds to achieve stable electron configurations, often resembling those of noble gases.

• Non-metals share electrons to complete their valence shells.

• Metals and non-metals form ionic compounds, resulting in ions with noble gas configurations.

Predicting Formulas of Ionic Compounds

Ionic compounds are formed to achieve electrical neutrality and maximum stability.

• Consider the valence electron configurations of the elements involved (e.g., oxygen and calcium).

• The empirical formula reflects the ratio of ions needed for neutrality (e.g., CaO for calcium oxide).

Naming Compounds

Binary Ionic Compounds (Type I)

Type I binary ionic compounds contain a metal (cation) and a non-metal (anion).

• The cation is named first, followed by the anion.

• Monatomic cations use the element name; monatomic anions use the root plus '-ide' (e.g., sodium chloride).

Common Monatomic Ions

Cation	Symbol	Anion	Symbol
Hydrogen	H+	Fluoride	F-
Lithium	Li+	Chloride	Cl-
Sodium	Na+	Bromide	Br-
Potassium	K+	Iodide	I-
Calcium	Ca2+	Oxide	O2-
Magnesium	Mg2+	Sulfide	S2-
Aluminum	Al3+	Nitride	N3-

Binary Ionic Compounds (Type II)

Type II compounds involve metals that can form more than one type of cation (usually transition metals).

• Use Roman numerals to indicate the charge (e.g., iron(III) chloride).

• Older system: '-ic' for higher charge, '-ous' for lower charge (e.g., ferric vs. ferrous).

• Group 1, Group 2, Al, Ag do not use Roman numerals.

Binary Covalent Compounds (Type III)

Type III compounds are formed between two non-metals.

• Use prefixes to indicate the number of each atom (e.g., dinitrogen monoxide for N2O).

• Common prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-.

Ionic Compounds with Polyatomic Ions

Polyatomic ions are charged groups of covalently bonded atoms. Oxyanions are polyatomic ions containing oxygen.

• Oxyanions with fewer oxygens end in '-ite'; with more, '-ate' (e.g., nitrite NO2-, nitrate NO3-).

• With more than two oxyanions: 'hypo-' (fewest), 'per-' (most) as prefixes (e.g., hypochlorite ClO-, perchlorate ClO4-).

Oxyanion Naming Table

Prefix/Suffix	Example (Cl)
hypo-...-ite	hypochlorite (ClO-)
...-ite	chlorite (ClO2-)
...-ate	chlorate (ClO3-)
per-...-ate	perchlorate (ClO4-)

Hydrates

Hydrates are ionic compounds that include water molecules in their crystal structure.

• Water molecules are called 'waters of hydration.'

• When water is removed, the compound is 'anhydrous.'

• Name includes Greek prefixes for the number of water molecules (e.g., barium hydroxide octahydrate: Ba(OH)2·8H2O).

Acids

Acids are molecules in which one or more H+ ions are attached to an anion.

• If the anion ends in '-ide', the acid name uses 'hydro-' and '-ic' (e.g., hydrochloric acid for Cl-).

• If the anion ends in '-ate', the acid name ends in '-ic' (e.g., nitrate → nitric acid).

• If the anion ends in '-ite', the acid name ends in '-ous' (e.g., nitrite → nitrous acid).

Acid Naming Table

Anion Name	Acid Name
Perchlorate	Perchloric acid
Chlorate	Chloric acid
Chlorite	Chlorous acid
Hypochlorite	Hypochlorous acid
Chloride	Hydrochloric acid

Naming Hydrocarbons

Alkanes

Hydrocarbons are compounds containing only carbon and hydrogen. Alkanes are saturated hydrocarbons with only single bonds.

• General formula: $C_nH_{2n+2}$

• Name is based on the number of carbon atoms and ends with '-ane'.

Number of Carbons	Prefix	Formula	Name
1	meth-	CH4	methane
2	eth-	C2H6	ethane
3	prop-	C3H8	propane
4	but-	C4H10	butane
5	pent-	C5H12	pentane
6	hex-	C6H14	hexane
7	hept-	C7H16	heptane
8	oct-	C8H18	octane
9	non-	C9H20	nonane
10	dec-	C10H22	decane

Sizes of Ions

Factors Affecting Ionic Size

The size of ions influences the structure and stability of ionic solids.

• Determined by measuring the distance between ion centers in a crystal.

• Depends on the charge and position in the periodic table.

• Isoelectronic ions: Ions with the same number of electrons.

Trends in Ionic Size

• Cations are smaller than their parent atoms; anions are larger.

• Within an isoelectronic series, higher nuclear charge means smaller ion size.

Localized Electron (LE) Model

Overview

The LE model describes molecules as collections of atoms bound by shared electron pairs (bonds) and lone pairs (non-bonding electrons).

• Lone pairs: Electron pairs localized on a single atom.

• Bonding pairs: Electron pairs shared between atoms.

Parts of the LE Model

• Lewis structures: Show valence electron arrangement.

• VSEPR model: Predicts molecular geometry based on electron pair repulsion.

• Atomic orbitals: Describe the regions where electrons are likely to be found.

Lewis Structures

Lewis structures use dots to represent valence electrons and lines for bonds.

• Only valence electrons are shown.

• Atoms achieve noble gas configurations (octet rule).

The Duet and Octet Rules

• Duet rule: Hydrogen forms stable molecules with two electrons in its valence shell.

• Octet rule: Main group elements form stable molecules when surrounded by eight valence electrons.

Exceptions to the Octet Rule

• Boron can have fewer than eight electrons (e.g., BH3).

• Elements in Period 3 and beyond can have expanded octets (e.g., SF6).

Formal Charge

Formal charge helps determine the most likely Lewis structure for a molecule or ion.

• Calculated as:

$\text{Formal charge} = \text{(valence electrons on free atom)} - \text{(non-bonding electrons)} - \frac{1}{2}\text{(bonding electrons)}$

• The sum of formal charges in a molecule or ion must equal the overall charge.

Examples

• Lewis structures for molecules like CH4, H2O, F2, BH3, SF6, XeF4, PF5.

• Assigning formal charges to atoms in ions such as NO3- and CN-.