Ionic and Covalent Chemical Bonds

Chemical Bonds

• A chemical bond involves an attractive force that holds atoms together in a more complex unit.
• Chemical bonds result from interactions between electrons in combining atoms.
• Two primary types of chemical bonds exist: Ionic and Covalent.

Ionic Bonds

• Ionic bonds form through the transfer of one or more electrons.
• The transfer happens from one atom or group of atoms to another.
• An ionic compound has ionic bonds.

Covalent Bonds

• Covalent bonds form through the sharing of one or more pairs of electrons between two atoms.
• Molecular or covalent compounds consist of atoms joined through covalent bonds.

General Bonding Principles

• Most bonds are not entirely ionic or covalent; they exhibit a degree of both ionic and covalent characteristics.
• Valence electrons are the electrons participate in bonding.
• Certain arrangements of electrons are more stable than others based on the octet rule.

Valence Electrons and Lewis Symbols

• Valence electrons are electrons found in the outermost electron shell of representative or noble-gas elements.
• Valence electrons are located in either s or p subshells.
• A Lewis Symbol is a chemical symbol surrounded by dots equal to the number of valence electrons.
• Elements in the same group have the same number of valence electrons.
• The number of valence electrons for representative elements corresponds to the Roman numeral periodic-table group number.
• The maximum number of valence electrons for any element is eight.
• Calcium (Ca) has 2 valence electrons (4s²).
• Selenium (Se) has 6 valence electrons (4s²4p⁴).
• Carbon (C) has 4 valence electrons (2s²2p²).

The Octet Rule

• Certain arrangements of valence electrons lead to greater stability.
• Noble gases' valence electron configurations have the most stable valence electron configurations.
• Atoms lose, gain, or share electrons to achieve a noble-gas electron configuration when forming compounds.

The Ionic Bond Model

• An ion is an atom, or group of atoms, that has an electrical charge due to the loss or gain of electrons.
• An atom that gains electrons becomes a negatively charged ion.
• An atom that loses electrons becomes a positively charged ion.
• For instance: when a potassium atom loses one electron, it forms K⁺; when a sulfur atom gains two electrons, it forms S²⁻.

The Sign and Magnitude of Ionic Charge

• Atoms gain or lose electrons to achieve an electron configuration like a noble gas.
• Example: K⁺ (1s²2s²2p⁶3s²3p⁶) loses one electron to achieve the electron configuration of Argon (Ar), which is 1s²2s²2p⁶3s²3p⁶.
• Metal atoms with one, two, or three valence electrons tend to lose electrons to achieve a noble-gas configuration.
  • Group IA forms ions with a 1+ charge.
  • Group IIA forms ions with a 2+ charge.
  • Group IIIA forms ions with a 3+ charge.
• Nonmetal atoms with five, six, or seven valence electrons tend to gain electrons to achieve a noble-gas configuration.
  • Group VIIA forms ions with a 1- charge.
  • Group VIA forms ions with a 2- charge.
  • Group VA forms ions with a 3- charge.
• Group IVA elements occupy a unique position as they could either gain or lose four electrons.
• Isoelectronic species are a series of ions or atoms with the same number and configuration of electrons, such as O²⁻, F⁻, Ne, Na⁺, Mg²⁺, and Al³⁺, all having the configuration 1s²2s²2p⁶.

Lewis Structures for Ionic Compounds

• Ion formation requires two elements, a metal to donate electrons and a nonmetal to accept them.
• The electrons lost by the metal are the same ones gained by the nonmetal.
• Simultaneously formed positive and negative ions attract one another.
• A Lewis structure is using Lewis symbols to represent the transfer or sharing of electrons in chemical bonds.

Chemical Formulas for Ionic Compounds

• Ionic compounds must be neutral, with no net charge.
• Positive and negative ions combine in a ratio that achieves charge neutrality.
• The symbol for the positive ion, cation, is written first.
• Charges on the ions are not included in the formula.
• Subscripts indicate the combining ratio of the ions.
• For example, the compound formed between Li⁺ and O²⁻ where two Li⁺ ions are needed to balance the 2⁻ charge on oxygen, yields a formula of Li₂O.
• Chemical formulas for compounds formed by interacting ion pairs:
  • Ba²⁺ and Cl⁻ form BaCl₂
  • Fe³⁺ and O²⁻ form Fe₂O₃
  • Pb⁴⁺ and O²⁻ form PbO₂

The Structure of Ionic Compounds

• Solid ionic compounds consist of positive and negative ions closely arranged, with each ion surrounded by neighbors of the opposite charge.
• Ions are bonded by electrostatic attractions to all surrounding ions of opposite charge.
• Formula Unit is the smallest whole-number repeating ratio of ions in an ionic compound, ensuring charge neutrality.
• Chemical formulas for ionic compounds represent the simplest ratio of ions present.

Recognizing and Naming Binary Ionic Compounds

• Binary compounds have two elements and include ionic and covalent compounds.

• Binary ionic compounds contain a metal and a nonmetal.

• The metal is always present as the positive ion, and the nonmetal as the negative ion.

• The full name of the metallic element is given first, followed by a separate word with the stem of the nonmetallic element name and the suffix -ide.

• Metallic elements with variable charges can form more than one type of positive charge, requiring specification of the charge on the metal ion using Roman numerals for transition metal cations.

• Examples of naming binary ionic compounds:

  • KCl is Potassium chloride
  • MgBr₂ is Magnesium bromide
  • CaO is Calcium oxide
  • CuBr is Copper(I) bromide
  • FeS is Iron(II) sulfide
  • PbO₂ is Lead(IV) oxide

Polyatomic Ions

• A polyatomic ion is an ion formed from a group of atoms (held together by covalent bonds) through the loss or gain of electrons.
• Common compounds containing polyatomic ions:
  • NaOH is Sodium hydroxide
  • Mg(NO₃)₂ is Magnesium nitrate
  • (NH₄)₂SO₄ is Ammonium sulfate.
• Most polyatomic ions have a negative charge.
• Two negatively charged polyatomic ions, OH⁻ and CN⁻, have names ending in -ide
• The rest of them typically end in -ate or -ite.
• For -ate,-ite pairs, the -ate ion has one more oxygen atom than the -ite ion, both carrying the same charge.
• In pairs of ions differing by a hydrogen atom, the charge on the ion containing hydrogen is always one less than that on the other ion.

Chemical Formulas and Names for Ionic Compounds Containing Polyatomic Ions

• The positive and negative charges present must add to zero.
• Examples of chemical formulas using polyatomic ions:
  • Na⁺ and OH⁻ form NaOH
  • Mg²⁺ and NO₃⁻ form Mg(NO₃)₂
  • NH₄⁺ and SO₄²⁻ form (NH₄)₂SO₄.