Section 1: Stability in Bonding

• Combined Elements

  • Some of the matter around you is in the form of uncombined elements such as copper, sulfur, and oxygen.

  • An observation you will make is that the compound formed when elements combine often has chemical and physical properties that aren’t anything like those of the individual elements.

  • Sodium is a soft, silvery metal that combines with chlorine, a greenish yellow gas, to form sodium chloride, which is a white crystalline solid.

• Formulas

  • Chemical Formula: tells what elements a compound contains and the exact number of the atoms of each element in a unit of that compound.

  • A subscript written after a symbol tells how many atoms of that element are in a unit of the compound.

    • If a symbol has no subscript, the unit contains only one atom of that element.

• Atomic Stability

  • The electric forces between electrons and protons, which are oppositely charged, hold atoms and molecules together, and thus they are the forces that cause compounds to form.

  • Atoms of noble gases are unusually stable.

  • To understand the stability of the noble gases, it is helpful to look at electron dot diagrams across a period.

    • Electron dot diagrams show only the electrons in the outer energy level of an atom.

  • The number of electrons in each group’s outer level increases across the table until the noble gases in Group 18, where each has a complete outer energy level.

  • An atom is chemically stable when its outermost energy level has the maximum number of electrons.

  • Electron dot diagrams of noble gases show that they each have a stable outer energy level.

  • Hydrogen and helium, the elements in period 1 of the periodic table, can hold a maximum of two electrons in their outer energy levels.

  • When you look at the elements in Groups 13 through 17, you see that none of the elements has a stable energy level.

  • Atoms with partially stable outer energy levels can lose, gain, or share electrons to obtain stable outer energy levels.

  • Sodium, as a Group1 element, will lose one electron. It will have the same number of electrons as neon. Chlorine, which belongs to Group 17, will gain one electron. It will have the same number of electrons as argon.

  • In water, hydrogen and oxygen each contribute one electron to each hydrogen-oxygen bond. The atoms share those electrons instead of giving them up.

  • Ion: a charged particle because it now has either more or fewer electrons than protons.

  • When a potassium atom loses an electron, the atom becomes positively charged because there is one electron less in the atom than there are protons in the nucleus.

  • An iodine atom has seven electrons in its outer energy level.

  • Potassium and iodine must perform a transfer of one electron. Potassium and iodine end up with stable outer energy levels.

• Ionic Bond: the force of attraction between the opposite charges of the ions in an ionic compound.

  • In an ionic bond, a transfer of one or more electrons takes place.

  • A magnesium atom gives an electron to each of two chlorine atoms to form MgCl2.

  • The positive charge of the magnesium ion is equal to the negative charge of the two chloride ions.

  • Metals and nonmetals usually combine by forming ionic bonds.

• Sharing Electrons

  • Some atoms of nonmetals are unlikely to lose or gain electrons.

  • Each time an electron is removed, the nucleus holds the remaining electrons even more tightly.

  • Covalent Bond: The attraction that forms between atoms when they share electrons

  • Molecule: A neutral particle that forms as a result of electron sharing

  • A single covalent bond is made up of two shared electrons

  • A water molecule contains two single bonds.

  • A covalent bond also can contain more than one pair of shared electrons.

  • Covalent bonds form between nonmetallic elements.

  • Many covalent compounds are liquids or gases at room temperature.

  • Electrons are not always shared equally between atoms in a covalent bond.

  • The strength of the attraction of each atom to its electrons is related to the size of the atom, the charge of the nucleus, and the total number of electrons the atom contains.

  • The chlorine atom exerts the greater pull on the electrons in hydrogen chloride, which forms hydrochloric acid in water.

  • Polar Molecule: one that has a slightly positive end and a slightly negative end, although the overall molecule is neutral.

  • The polarity of water is responsible for many of its unique properties

  • Nonpolar Molecule: one in which electrons are shared equally in bonds.

• Properties of Compounds

  • A compound whose atoms are held together by covalent bonds is a covalent compound.

  • A compound that is composed of ions is an ionic compound.

  • Table sugar is an example of a covalent compound that is a crystalline solid soluble in water. Ionic compounds, such as table salt, are also soluble in water.

  • The chemical and physical properties of covalent compounds and ionic compounds are different.

  • The covalent bonds between atoms in molecules are strong.

  • Melting and boiling points of covalent compounds are relatively lower when compared to ionic compounds.

  • The ionic bonds between ions are relatively strong.

  • Some common ionic compounds used in everyday life are potassium chloride as a salt substitute; potassium iodide added to table salt; and sodium fluoride added to toothpaste.

Section 3: Writing Formulas and Naming Compounds

• Binary Ionic Compounds

  • Binary Compound: one that is composed of two elements.

  • Because all elements in a given group have the same number of electrons in their outer energy levels, they must gain or lose the same number of electrons.

  • Oxidation Number: The charge on the ion'

  • When naming these compounds, the oxidation number is expressed in the name with a roman numeral.

  • When writing formulas, it is important to remember that although the individual ions in a compound carry charges, the compound itself is neutral.

  • A formula must have the same number of positive ions and negative ions so the charges balance.

  • After you’ve learned how to find the oxidation numbers and their least common multiple, you can write formulas for ionic compounds.

• Compounds with Polyatomic Ions

  • Not all ionic compounds are binary.

  • Some ionic compounds are composed of more than two elements. They contain polyatomic ions.

  • Polyatomic Ion: a positively or negatively charged, covalently bonded group of atoms.

  • To write formulas for these compounds, follow the rules for binary compounds, with one addition. When more than one polyatomic ion is needed, write parentheses around the polyatomic ion before adding the subscript.

• Compound with Added Water

  • Hydrate: a compound that has water chemically attached to its ions and written into its chemical formula.

  • To write the formula for a hydrate, write the formula for the compound and then place a dot followed by the number of water molecules.

• Naming Binary Covalent Compounds

  • Covalent compounds are those formed between elements that are nonmetals

  • The main ionic compound is named the regular way, but the number of water molecules in the hydrate is indicated by the Greek prefix.