Chemistry Chapter 4 Flashcards

Ion Formation

• Ions form when an atom loses a valence electron, resulting in a positive charge, or gains an electron, giving it a negative charge.

Ionic Compounds

• Ionic compound formulas consist of the positive ion followed by the negative ion, with subscripts to balance charges.
• Naming conventions for ionic compounds place the name of the positive ion first, followed by the negative ion.

Properties of Ionic Compounds

• Characterized by hard, brittle crystals with high melting points.
• Conduct electricity when dissolved in water or melted.

Noble Gases

• Noble gases are nonreactive due to having a complete valence shell of electrons.

Valence Electrons

• Valence electrons are the highest energy electrons that determine an element's chemical properties.

Electron Dot Diagram

• Depicts a chemical symbol surrounded by dots that represent the atom's valence electrons.

Chemical Bonds

• Chemical bonds are the attractive forces that hold atoms together, resulting from electron rearrangement.

Atom Bonding

• Atoms bond to achieve stability by filling their valence shells with electrons.

Element Chemistry

• An element's chemistry is primarily determined by its number of valence electrons.

Group Properties

• Elements within a group share the same number of valence electrons, leading to similar properties.

Behavior of Metals

• Metals typically react by losing their valence electrons.
• Metal reactivity decreases from left to right across the periodic table based on the ease of losing valence electrons.

Nonmetals and Stability

• Nonmetals achieve stability by gaining or sharing electrons to complete their valence shell.

Halogens

• Halogens (Group 17) are highly reactive, needing only one more electron to fill their valence shell.

Metalloids

• Metalloids can both lose and share electrons during bonding with other elements.

Hydrogen

• Hydrogen is unique on the periodic table as a nonmetal with distinct properties compared to alkali metals, often forming bonds by sharing its electron.

Valence Electron Characteristics

• Eight valence electrons confer stability, rendering an element nonreactive.

Ions

• Ions are atoms or groups of atoms with an electric charge, which can be positive or negative.
• Polyatomic ions consist of multiple atoms and can have overall positive or negative charges.

Ionic Bonds and Compounds

• Ionic bonds are formed through the attraction between oppositely charged ions, resulting in ionic compounds that maintain an equal ratio of positive and negative ions.

Chemical Formulas

• Chemical formulas utilize symbols to indicate the ratios of elements in compounds, with subscripts showing the quantity of each element.

Crystal Structure

• Ionic compounds arrange in an orderly, three-dimensional crystal structure.

Covalent Bonds

• Covalent bonds form when two atoms share electrons; the bond strength is due to attractions between shared electrons and nuclei.

Molecular Compounds

• Composed of molecules, molecular compounds typically do not conduct electricity like ionic compounds and generally have lower melting and boiling points.

Polar and Nonpolar Bonds

• Polar bonds result from unequal electron sharing; nonpolar bonds result from equal sharing of electrons.

Metallic Bonds

• Metallic bonds arise from the attraction between positively charged metal ions and delocalized electrons, allowing metals to conduct heat and electricity effectively.

Alloys

• Alloys are mixtures containing two or more elements, at least one being a metal.

Metal Properties

• Metals exhibit a shiny luster, high malleability, ductility, and excellent electrical and thermal conductivity.

Covalent Bonding Details

• Single, double, and triple bonds refer to the number of electron pairs shared between atoms.

Conductivity in Metallic Bonds

• Metallic bonds facilitate electrical and thermal conductivity due to the free movement of valence electrons.

Examples

• Nonpolar bond example: H₂ molecule, because of equal electron sharing.
• Polar bond example: Water (H₂O), where oxygen attracts electrons more than hydrogen.