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Learn about valence electrons, the formation of positive and negative ions, and the nature of ionic bonds. Explore the properties and energy changes in ionic compounds, as well as the naming and formulas for these compounds.
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Ionic Compounds • Chapter 8
Section 8.1 - Forming Chemical Bonds 1. Valence Electrons & Chemical Bonds • A chemical bondis the force that holds two atoms together. • Chemical bonds form by the attraction between the positive nucleus of one atom and the negative electrons of another atom. • Atom’s try to form the octet—the stable arrangement of eight valence electrons in the outer energy level—by gaining or losing valence electrons.
2. Positive Ion Formation • A positively charged ion is called a cation. • This figure illustrates how sodium loses one valence electron to become a sodium cation. • Metals are reactive because they lose valence electrons easily. • Transition metals commonly form 2+ or 3+ ions, but can form greater than 3+ ions. • Other relatively stable electron arrangements are referred to as pseudo-noble gas configurations.
3. Negative Ion Formation • An anionis a negatively charged ion. • The figure shown here illustrates chlorine gaining an electron to become a chlorine ion. • Nonmetal ions gain the number of electrons required to fill an octet. • Some nonmetals can gain or lose electrons to complete an octet.
Section 8.2 - The Formation and Nature of Ionic Bonds 4. Formation of an Ionic Bond • The electrostatic force that holds oppositely charged particles together in an ionic compound is called an ionic bond. • Compounds that contain ionic bonds are called ionic compounds. • Binary ionic compounds contain only two different elements—a metallic cation and a nonmetallic anion.
5. Properties of Ionic Compounds • Positive and negative ions exist in a ratio determined by the number of electrons transferred from the metal atom to the non-metal atom. • The repeating pattern of particle packing in an ionic compound is called an ionic crystal. • The strong attractions among the positive and negative ions result in the formation of the crystal lattice. • A crystal latticeis the three-dimensional geometric arrangement of particles, and is responsible for the structure of many minerals.
Properties of Ionic Compounds (cont.) • Melting point, boiling point, and hardness depend on the strength of the attraction. • In a solid, ions are locked into position and electrons cannot flow freely—solid ions are poor conductors of electricity. • Liquid ions or ions in aqueous solution have electrons that are free to move, so they conduct electricity easily. • An ion in aqueous solution that conducts electricity is an electrolyte.
6. Energy and the Ionic Bond • Reactions that absorb energy are endothermic. • Reactions that release energy are exothermic. • Energy changes occur during the formation of ionic bonds from ions during a chemical reaction. • The formation of ionic compounds is always exothermic. • A more stable system is formed that is lower in energy than the individual atoms. • If the amount of energy released during bond formation is added to an ionic compound, the bonds that hold the ions together will break.
Energy and the Ionic Bond (cont.) • The energy required to separate 1 mol of ions in an ionic compound is referred to as the lattice energy. • Strength of the forces holding ions in place is reflected by the lattice energy • The more negative the lattice energy, the stronger the force of attraction. • Lattice energy is directly related to the size of the ions that are bonded. • Smaller ions generally have a more negative value for lattice energy, and require more energy to separate. • Electrostatic force of attraction is inversely related to the distance between the opposite charges.
Energy and the Ionic Bond (cont.) • The value of lattice energy is also affected by the charge of the ion. • The ionic bond formed from the attraction of ions with larger positive or negative charges has a more negative lattice energy. • Ex. The lattice energy of MgO is almost 4 times greater than that of NaF. • The lattice energy of SrCl2 is between those of MgO and NaF.
8.3 Names and Formulas for Ionic Compounds 7. Formulas for Ionic Compounds • When writing names and formulas for ionic compounds, the cation appears first followed by the anion. • Chemists around the world need to communicate with one another, so a standardized system of naming compounds was developed. • A formula unit represents the simplest ratio of the ions involved. • Monatomic ionsare one-atom ions. • Oxidation number, or oxidation state, is the charge of a monatomic ion. • .
Formulas for Ionic Compounds (cont.) • The symbol for the cation is always written first, followed by the symbol of the anion. • Subscripts represent the number of ions of each element in an ionic compound. • The total charge must equal zero in an ionic compound.
Formulas for Ionic Compounds (cont.) • Polyatomic ionsare ions made up of more than one atom. • Never change subscripts of polyatomic ions, place in parentheses and write the appropriate subscript outside the parentheses
8. Names for Ions and Ionic Compounds • An oxyanionis a polyatomic ion composed of an element (usually a non-metal), bonded to one or more oxygen atoms.
Chemical nomenclature is a systematic way of naming compounds. • Name the cation followed by the anion. • For monatomic, cations use the element name. • For monatomic anions, use the root element name and the suffix –ide. • To distinguish between different oxidation states of the same element, the oxidation state is written in parentheses after the name of the cation. • When the compound contains a polyatomic ion, name the cation followed by the name of the polyatomic ion. Names for Ions and Ionic Compounds (cont.)
8.4 Metallic Bonds and Properties of Metals 9. Metallic Bonds and the Properties of Metals • Metals are not ionic but share several properties with ionic compounds. • Metals also form lattices in the solid state, where 8 to 12 other atoms closely surround each metal atom. • Within the crowded lattice, the outer energy levels of metal atoms overlap. • The electron sea modelproposes that all metal atoms in a metallic solid contribute their valence electrons to form a "sea" of electrons.
Metallic Bonds and the Properties of Metals (cont.) • The electrons are free to move around and are referred to as delocalized electrons, forming a metallic cation. • A metallic bondis the attraction of an metallic cation for delocalized electrons. • Boiling points are much more extreme than melting points because of the energy required to separate atoms from the groups of cations and electrons.
Metallic Bonds and the Properties of Metals (cont.) • Metals are malleable because they can be hammered into sheets. • Metals are ductile because they can be drawn into wires. • Mobile electrons around cations make metals good conductors of electricity and heat. • As the number of delocalized electrons increases, so does hardness and strength.
10. Metal Alloys • An alloy is a mixture of elements that has metallic properties. • The properties of alloys differ from the elements they contain. • Substitutional alloys are formed when some atoms in the original metallic solid are replaced by other metals of similar atomic structure. • Interstitial alloys are formed when small holes in a metallic crystal are filled with smaller atoms.
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