Lewis Dot Symbols and Ionic Bonds

• The periodic table and electron configuration provide a rationale for molecule and compound formation.
• Atoms combine to achieve a stable electron configuration, ideally isoelectronic with a noble gas.
• Chemical bonding primarily involves the valence electrons (outer region electrons).
• Lewis dot symbols track valence electrons during chemical reactions, ensuring electron conservation.
• A Lewis dot symbol includes the element symbol and dots representing valence electrons.
• Except for helium, the number of valence electrons equals the group number for representative elements.
• Elements in the same group possess similar outer electron configurations and Lewis dot symbols.
• Transition metals, lanthanides, and actinides have incomplete inner shells, making simple Lewis dot symbols difficult to write.

Lewis Dot Symbols

• Illustrate bonds between atoms
• Predict stability of products
• Predict the number of bonds

Ionic Bonds

• Elements with low ionization energies tend to form cations, while those with high electron affinities form anions.
• Alkali and alkaline earth metals form cations
• Halogens and oxygen form anions.
• 1A and 2A metals form bonds with oxygen and halogens
• An ionic bond is defined as an electrostatic force holding ions together in an ionic compound.
• In the reaction between lithium and fluorine, the 2s1 valence electron of lithium is transferred to fluorine.
• LiF is obtained industrially by purifying minerals containing the compound.
• Resulting calcium ion Ca2+ has the argon electron configuration
• The oxide ion (O²¯) is isoelectronic with neon
• The compound (CaO) is electrically neutral.

Balancing Charges

• In many cases, the cation and the anion in a compound do not carry the same charges to maintain balance.
• The oxygen atom receives two electrons (one from each of the two lithium atoms) to form the oxide ion in lithium oxide (Li2O)
• The Li⁺ ion is isoelectronic with helium.
• Magnesium reacts with nitrogen at elevated temperatures to form magnesium nitride (Mg3N2).
• The reaction transfers six electrons (two from each Mg atom) to two nitrogen atoms in magnesium nitride.
• The resulting magnesium ion (Mg2+) and the nitride ion (N³¯) are both isoelectronic with neon
• The charges are balanced, and the compound is electrically neutral due to three +2 ions and two -3 ions.

Lattice Energy

• Lattice energy indicates compound stability
• Is always a positive quantity because the separation of ions in a solid into ions in the gas phase is, by Coulomb's law, an endothermic process.
• Coulomb's law is used to calculate lattice energy (potential enery)
• Indicates that potential energy (E) between two ions is directly proportional to the product of their charges.
• Born Haber cycle calculates lattice energy indirectly
• Relates lattice energies of ionic compounds to ionization energies, electron affinities, and other atomic and molecular properties

Covalent Bonds

• A covalent bond is a bond in which two atoms share two electrons.
• Covalent compounds contain only covalent bonds.
• Shared electron pairs can be represented by a single line
• Covalent bonding between many-electron atoms involves only the valence electrons.
• Only two valence electrons participate in the formation of F2
• Other, non-bonding electrons, are called lone pairs—pairs of valence electrons that are not involved in covalent bond formation
• Lewis structures show shared electron pairs either as lines or as pairs of dots between two atoms, and lone pairs as pairs of dots on individual atoms
• Valence electrons are shown only

Octet Rule

• Atoms other than hydrogen tend to form bonds until surrounded by eight valence electrons.
• A covalent bond forms when not enough electrons exist for individual atoms to have a complete octet.
• Hydrogen requires the electron configuration of helium, or a total of two electrons.
• Works mainly for elements in the second period of the periodic table
• These elements only have 2s and 2p subshells

Multiple Bonds

• Single Bond = held together by one electron pair
• Multiple Bonds = Bonds formed when two atoms share two or more pairs of electrons

Double Bond

• When two atoms share two pairs of electrons, the covalent bond is called a double bond.
• Found in molecules of carbon dioxide (CO2) and ethylene (C2H4)

Triple Bond

• Arises when two atoms share three pairs of electrons, as in the nitrogen molecule (N2)

Bond Length

• Defined as the distance between the nuclei of two covalently bonded atoms in a molecule
• Multiple bonds are shorter than single bonds for the same pair of atoms

Electronegativity

• Defined as the ability of an atom to attract the electrons in a chemical bond toward itself
• Affects polarity
• Related to electron affinity and ionisation energy
• Elements with high electronegativity have a greater tendency to attract electrons than do elements with low electronegativity
• Electronegativity increases from left to right across a period and decreases down a group in the periodic table

Types of Bonds

• Ionic bonds form when there is an electronegativity differences above 2.0
• Polar covalent bonds form when the electronegativity difference is 0.3-2.0
• Covalent bonds have an electronegativity below 0.3 with little or no polarity

Electronegativity vs Electron Affinity

• Electronegativity signifies the ability of an atom in a chemical bond (with another atom) to attract the shared electrons
• Electron affinity refers to an isolated atom's attraction for an additional electron
• Electron affinity is experimentally measurable
• Electronegativity is an estimated number that cannot be measured

Oxidation Number

• It is the number of charges an atom would have if electrons were transferred completely to the more electronegative of the bonded atoms in a molecule
• Can be calculated based on electronegativity