Chemistry Valence Electrons and Bonding Quiz

Valence Electrons and Energy Levels

• The maximum number of electrons that can be accommodated in the valence energy level is eight.
• Valence electrons are the electrons in the outermost energy level of an atom and are involved in chemical bonding.

Formation of Ions

• A positive ion (cation) is formed when an atom loses electrons, resulting in a net positive charge.

Metalloids

• Metalloids are a group of elements that exhibit properties of both metals and nonmetals.

Ionic Bonding

• In ionic bonding, a cation (positively charged ion) and an anion (negatively charged ion) are held together by electrostatic attraction.
• The cation loses electrons and the anion gains electrons.

Covalent Bonding

• A covalent bond is formed when two atoms share electrons.
• Triple bonds are formed when atoms share three pairs of electrons.

Nonmetals

• Nonmetals are elements that are typically poor conductors of heat and electricity.
• Nonmetals tend to gain electrons to form negative ions.

Types of Bonds

• Ionic bonds are formed when electrons are transferred from one atom to another.
• Covalent bonds are formed when atoms share electrons.
• If two atoms have identical electronegativities, they form a nonpolar covalent bond.
• Polar covalent bonds form when electrons are shared unequally between atoms due to differences in electronegativity.
• An electronegativity difference of 1.7 or greater indicates an ionic bond.

Electron Dot Diagrams

• In an electron dot diagram for an ionic compound, the unpaired electrons represent the electrons that are transferred to form the ions.

Ionic Compounds

• Ionic compounds are formed by the electrostatic attraction between oppositely charged ions.
• The resulting charge of a negative ion (anion) after gaining electrons is negative.
• Ionic compounds typically have a crystal lattice structure, characterized by a regular, repeating arrangement of ions.
• Ionic compounds have high melting and boiling points due to the strong electrostatic forces holding the ions together.

Metallic Bonding

• Metallic bonding is characterized by a "sea" of delocalized electrons that move freely throughout the crystal lattice.

Types of Crystals

• Molecular crystals are formed by weak intermolecular forces holding molecules together.
• Covalent network crystals are characterized by a continuous network of covalently bonded atoms and have very high melting points.
• Molecular crystals have weak intermolecular forces compared to ionic and covalent bonds.

Metallic Bonds

• Valence electrons in metallic bonds behave like a sea of delocalized electrons. This is what allows metals to conduct electricity and heat well.

Examples

• Diamond is an example of a covalent network crystal.

Oxygen

• Oxygen has six valence electrons.

Valence Electrons in Bonding

• Valence electrons are the electrons that participate in chemical bonding.
• The atom with the highest number of valence electrons is oxygen.

Lone Pairs

• In the context of valence orbitals, "lone pairs" refer to pairs of electrons that are not involved in bonding.
• Chlorine has three lone pairs in its electron configuration.

Energy Levels

• The first energy level of an atom can hold a maximum of two electrons.

Lewis Structures

• In an electron dot diagram, chlorine's valence electrons are represented as four dots surrounding the chemical symbol Cl.
• During electron transfer, the metal atom becomes positively charged and the nonmetal atom becomes negatively charged.
• In a Lewis dot structure, the atom with the lowest electronegativity typically occupies the center position.
• A water molecule (H2O) has a total of eight valence electrons.
• In a structural formula for a covalent bond, a line represents a single bond.
• Carbon has a bonding capacity of four based on its valence electrons.
• In a Lewis structure example where hydrogen atoms are bonded on either side of a carbon atom, the carbon atom serves as the central atom.
• Drawing Lewis structures involves arranging atoms, placing valence electrons, and connecting atoms with lines to represent bonds.
• A single bond in structural formulas represents two bonding electrons.

Polarity

• Whether a molecule is polar or nonpolar is determined by the arrangement of its atoms and the types of bonds present.
• Dipole-dipole forces are specifically associated with polar molecules.
• London dispersion forces occur between molecules with no permanent dipoles.
• The atoms in a water molecule are arranged in a bent shape.
• A molecular dipole is formed due to an uneven distribution of electron density.

Intermolecular Forces

• Intermolecular forces are weaker than the forces holding atoms together in molecules.
• In a nonpolar molecule like O=N=O, a zero total dipole results from the symmetrical arrangement of the atoms and the equal sharing of electrons.
• London dispersion forces are considered the weakest type of intermolecular force.