BackBonding: General Concepts – Study Notes
Study Guide - Smart Notes
Tailored notes based on your materials, expanded with key definitions, examples, and context.
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. Molecules can be represented in several ways to convey their structure and composition.
Chemical bond: The force that holds atoms together in a compound.
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 for water).
Condensed formula: Groups atoms to show connectivity (e.g., CH3CH2CHBrCH3).
Representation of Molecular Structures
Understanding the three-dimensional arrangement of atoms in a molecule is essential for predicting its properties and reactivity. Two common models are used:
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 the three-dimensional structure.
Attributes of Molecules
Molecules exhibit various physical and chemical properties that depend on their structure and bonding.
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. The resulting ions are held together by electrostatic attraction.
Ionic bond: Formed when one atom donates electrons and another atom accepts them.
Ionic compounds: Formed from the reaction of metals with non-metals.
Covalent Bonding
Covalent bonding occurs when atoms share electrons to achieve stable electron configurations. Non-metals commonly form covalent bonds with each other.
Electron sharing: Leads to the formation of molecules.
Hydrogen: Can 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 to achieve stability.
Electrons are located in the space between the two nuclei.
Simultaneous attraction between nuclei and shared electrons increases stability.
Potential energy decreases as attractive forces increase, reaching a minimum at the bond length.
Electronegativity and Bond Polarity
What is Electronegativity?
Electronegativity is the ability of an atom in a molecule to attract shared electrons. Differences in electronegativity between atoms lead to unequal sharing of electrons, resulting in bond polarity.
Electronegativity: Tendency of an atom to attract electrons in a chemical bond.
Bond polarity: Unequal sharing of electrons creates partial positive (δ+) and partial negative (δ-) charges.
Example: In H—F, fluorine is more electronegative and attracts electrons more strongly, resulting in a polar bond.
Measuring Electronegativity
Pauling's method compares bond energies to determine relative electronegativities.
For a molecule HX, compare the measured H—X bond energy to the average of H—H and X—X bond energies.
Formula:
Effect of Electric Field on Polar Molecules
Polar molecules, such as hydrogen fluoride (HF), align in an electric field due to their partial charges, demonstrating the presence of bond polarity.
In an electric field, δ+ and δ- ends of molecules orient toward opposite charges.
Electron Configuration and Ionic Compounds
Electron Configuration in Compounds
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 by transferring electrons.
Resulting ions achieve noble gas configurations.
Predicting Formulas of Ionic Compounds
Ionic compounds are formed to achieve electrical neutrality and maximum stability.
Empirical formulas reflect the simplest ratio of ions that results in a neutral compound (e.g., CaO).
Valence electron configurations and electronegativity differences guide predictions.
Naming Compounds
Binary Ionic Compounds (Type I)
Type I binary ionic compounds contain a metal (cation) and a non-metal (anion). Naming follows specific rules:
The cation is named first, followed by the anion.
Monatomic cations use the element name (e.g., sodium).
Monatomic anions use the root of the element name plus the suffix -ide (e.g., chloride).
Common Monatomic Cations and Anions
Cation | Symbol | Anion | Symbol |
|---|---|---|---|
Sodium | Na+ | Chloride | Cl- |
Potassium | K+ | Bromide | Br- |
Calcium | Ca2+ | Oxide | O2- |
Magnesium | Mg2+ | Sulfide | S2- |
Aluminum | Al3+ | Nitride | N3- |
Barium | Ba2+ | Phosphide | P3- |
Binary Ionic Compounds (Type II)
Type II compounds involve metals (usually transition metals) that can form more than one type of cation. The charge is specified using Roman numerals.
Higher charge: -ic suffix (historical, less common now).
Lower charge: -ous suffix (historical, less common now).
Roman numeral system: e.g., iron(III) chloride for FeCl3.
Elements forming only one cation (e.g., 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. Prefixes indicate the number of each atom present.
Examples:
N2O: dinitrogen monoxide (nitrous oxide)
NO: nitrogen monoxide (nitric oxide)
NO2: nitrogen dioxide
N2O3: dinitrogen trioxide
N2O4: dinitrogen tetroxide
N2O5: dinitrogen pentoxide
Ionic Compounds with Polyatomic Ions
Polyatomic ions are charged groups of covalently bonded atoms. Oxyanions are polyatomic ions containing oxygen and another element.
Oxyanions with fewer oxygens: -ite suffix (e.g., nitrite, NO2-).
Oxyanions with more oxygens: -ate suffix (e.g., nitrate, NO3-).
Series with more than two oxyanions:
Fewest oxygens: hypo- prefix (e.g., hypochlorite, ClO-).
Most oxygens: per- prefix (e.g., perchlorate, ClO4-).
Examples of Oxyanions
Prefix/Suffix | Chlorine | Bromine | Iodine |
|---|---|---|---|
Hypo-...-ite | ClO- (hypochlorite) | BrO- (hypobromite) | IO- (hypoiodite) |
-ite | ClO2- (chlorite) | BrO2- (bromite) | IO2- (iodite) |
-ate | ClO3- (chlorate) | BrO3- (bromate) | IO3- (iodate) |
Per-...-ate | ClO4- (perchlorate) | BrO4- (perbromate) | IO4- (periodate) |
Hydrates
Hydrates are ionic compounds that contain water molecules within their crystal structure. The number of water molecules is indicated by Greek prefixes.
When water is removed, the compound is called anhydrous.
Examples:
Ba(OH)2·8H2O: barium hydroxide octahydrate
CuSO4·5H2O: copper(II) sulfate pentahydrate
Greek Prefixes for Numbers
Number | Prefix |
|---|---|
1 | mono- |
2 | di- |
3 | tri- |
4 | tetra- |
5 | penta- |
6 | hexa- |
7 | hepta- |
8 | octa- |
9 | nona- |
10 | deca- |
Acids
Acids are molecules in which one or more hydrogen ions (H+) are attached to an anion. The naming depends on the type of anion present.
If the anion ends in -ide: use the prefix hydro- and the suffix -ic (e.g., hydrochloric acid for HCl).
If the anion is an oxyanion ending in -ate: use the suffix -ic (e.g., HNO3 is nitric acid).
If the anion is an oxyanion ending in -ite: use the suffix -ous (e.g., HNO2 is nitrous acid).
Examples of Acid Names
Anion Name | Acid Name |
|---|---|
Perchlorate | Perchloric acid |
Chlorate | Chloric acid |
Chlorite | Chlorous acid |
Hypochlorite | Hypochlorous acid |
Cl- | Hydrochloric acid |
NO3- | Nitric acid |
NO2- | Nitrous acid |
SO42- | Sulfuric acid |
SO32- | Sulfurous acid |
PO43- | Phosphoric acid |
CH3COO- | Acetic acid |
Naming Hydrocarbons
Alkanes
Hydrocarbons are compounds containing only carbon and hydrogen. Alkanes are saturated hydrocarbons with the general formula CnH2n+2.
Names are based on the number of carbon atoms, using a prefix and the suffix -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. Ionic size is determined by the number of electrons and the position of the element in the periodic table.
Isoelectronic ions: Ions with the same number of electrons but different nuclear charges.
As nuclear charge increases for isoelectronic ions, ionic radius decreases.
Localized Electron (LE) Model and Lewis Structures
LE Model
The Localized Electron (LE) Model describes molecules as collections of atoms bound by shared pairs of electrons (bonding pairs) and lone pairs (non-bonding pairs).
Bonding pairs: Electrons shared between atoms.
Lone pairs: Electrons localized on a single atom.
Lewis Structures
Lewis structures represent the arrangement of valence electrons among atoms in a molecule. Atoms tend to achieve noble gas electron configurations (octet rule).
Only valence electrons are shown, represented as dots.
Duet rule: Hydrogen achieves stability with two electrons.
Octet rule: Most main-group elements achieve stability with eight electrons.
Exceptions: Boron can have fewer than eight electrons; elements in Period 3 and beyond can exceed the octet.
Formal Charge
Formal charge is used to estimate the distribution of charge in a molecule and to identify the most stable Lewis structure.
Formula:
The sum of formal charges in a molecule or ion equals the overall charge.
Example: In the nitrate ion (NO3-), formal charges help determine which oxygen atom carries the negative charge.
Additional info: These notes expand on the original slides by providing definitions, examples, and context for key concepts in chemical bonding, suitable for introductory college-level chemistry or general biology students.
