Topics Chemical Bonds Force that holds atom together Stability in Bonding Valence Electrons Reactivity Types of Bonds Ionic, Metallic, Covalent, Hydrogen bond Writing Formulas and Naming Compounds Binary compound Polyatomic ion Hydrate Reactivity Definition: Reactivity describes how likely an element is to form bonds with other elements. Reactivity of an element is determined by its valence electrons. Valence Electrons Definition: Valence electrons are electrons in the outermost energy shell (level) that are available to be lost, gained or shared when elements form chemical bonds. Elements in the same group of the Periodic Table have similar chemical properties because of their valence electrons. Element Groups And Valence Electrons Stability in Bonding Group 1 13 14 15 16 17 Valence Electron 1 3 4 5 6 7 Category Alkali Metal Alkali Earth Halogen Reactivity Very High High High Very High Some elements combine chemically and no longer have the same properties they did before forming a compound. A chemical formula is composed of symbols and subscripts indicating the number of atoms of an element in a compound. E.g. NaCl, H O, CaCO 3, CO 18 8 Noble Gas Stable 1
Types of Bonds Atoms form compounds when the compound is more stable than the separate atoms. Noble gas are more chemically stable. Elements that do not have full outer energy shell are more stable in compounds. Atoms can lose, gain, or share electrons to get a stable outer energy level A chemical bond is the force that holds atoms together in a compound. Intramolecular (Strong) Intermolecular (Weak) Chemical Bonds Covalent Ionic Metallic Hydrogen Other noncovalent Intramolecular Force Intermolecular Force Definition: force and its property within the structure of a single molecule. Ionic, Covalent, and Metallic bond Definition: force that act between stable molecules or between functional groups of macromolecules. Example: dipole-dipole interaction, hydrogen bonds, di-sulphide bonds Cystein-Cystein Chemical bond Ionic Bond Definition: The force that holds atoms together in a compound. Ions are charged particles because it has more or fewer electrons than protons. When an atom loses an electron, it becomes a positively charged ion; a superscript indicates the charge, e.g., H +, Ca +, Li + When an atom gains an electron, it becomes a negatively charged ion, e.g., HCO 3-, OH - Definition: Force of attraction between opposite charges of the ions, often forming (crystal) lattice. An ionic compound is held together by the ionic bond, often joining a metal to a nonmetal The result of this bond is neutral compound. The sum of the charges on the ions is zero.
Covalent Bond Definition: the force of attraction between atoms sharing electrons. Molecules are neutral particles formed as a result of sharing electrons. Atoms can form double or triple bonds depending on whether they share two or three pairs of electrons. Covalent bonds are usually involved when two nonmetals form a compound or when a nonmetal bond with a metalloid. Electrons shared in a molecule are held more closely to the atoms with the larger nucleus. A polar molecule has one end that is slightly negative and one end that is slightly positive although the overall molecule is neutral. In a nonpolar molecule electrons are shared equally. Main Difference Ionic bonds form when atoms lose or gain electrons; covalent bonds form when atoms share electrons. Metallic Bond electromagnetic interaction between delocalized electrons, called conduction electrons and gathered in an "electron sea", and the metallic nuclei within metals. Form metallic bond joins metal to metal. This type of bonding is collective in nature. There is no single metallic bond. Electronegativity aka Electron Affinity Electronegativity measures the ability of an atom to attract electrons. Fluorine has the highest electronegativity. Ionization Energy Definition: The amount of energy needed to remove an electron from a neutral atom is called ionization energy. Moving from left to right, ionization energy increases across a period. Moving from top to bottom in a group, it generally decreases. 3
Drawing Chemical Bond Bohr s Model Valence shells are drawn in circles for s suborbital (Group 1, ) and p suborbital (Group 13 thru 18) Protons and Neutrons are indicated in nucleus. Electrons in shells travel in pairs. Lewis (Electron) Dot Diagram Easier way to illustrate electrons on outermost energy shell and their reactivity. Charged atoms and molecules are placed in brackets with their charges indicated at top right corner. Octet Rule Octet rule is a simple chemical rule of thumb that states atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas. This rule is applicable to main group elements. Valence Shells Bohr s Model Bond Using Bohr s Model 4
Electron Dot (Lewis) Diagram Chemical Nomenclature Nomenclature: The system of naming in arts or science Writing chemical formulas and naming compounds Binary Metal-Nonmetal Compounds Binary Nonmetal-Nonmetal Compounds Polyatomic ions Hydrate Oxidation Number Oxidation Number # of e - gained, lost, or shared to become stable Determine the oxidation number from the periodic table. 5
Balancing Oxidation Number Ba+ F Binary Metal-Nonmetal Compounds Composed of two elements When writing binary compounds: Use the least common multiples of oxidation number. Compound is neutral. Charges on the ions must be balanced. Use first element, the root name of the second element, and the suffix -ide. More examples If the metal has more than one possible charge, indicate the ion with charge in roman numerals FeCl - Iron (II) chloride. Alternatively, common names may be used. Add ous for the lower charge and ic for the higher charge: FeCl ferrous chloride FeCl 3 ferric chloride Compound FeF FeF 3 Hg Br HgBr Systematic Name Iron (II) fluoride Iron (III) fluoride Mercury (I) Mercury (II) Common Name Ferrous fluoride Ferric fluoride Mercurous Mercuric Binary Nonmetal-Nonmetal Compounds Add ide to the second element Use Greek prefixes for number of atoms: 1 mono 6 hexa di 7 hepta 3 tri 8 octa Examples: CO carbon monoxide CO carbon dioxide N O 5 dinitrogen pentoxide 4 tetra 9 nona 5 penta 10 - deca Common names: ous and ic (-ic has greater charge, OR has fewer atoms) Formula NO N O NO N O 5 N O 3 Systematic Name Nitrogen monoxide Dinitrogen monoxide Nitrogen dioxide Dinitrogen pentoxide Dinitrogen trioxide Common Name Nitric oxide Nitrous oxide Nitrogen peroxide Nitric anhydride Nitrous anhydride 6
Polyatomic Ion The compound contains three or more elements. When writing names: Positive ions first, followed by negative ion Use the least common multiple of oxidation number, and put parenthesis around the polyatomic ion before adding a subscript Balancing Oxidation Number K + OH Ca + NO 3 Polyatomic Compounds Ammonium ion NH 4 1+ -ide ions: CN cyanide, OH hydroxide Oxyanions -ate are more oxygen NO nitrite NO 3 nitrate Some oxyanions have extra hydrogen SO 4 - sulfate HSO 4 hydrogen sulfate (or bisulfate) Oxyanions (continued) If more than two possibilities Formula ClO ClO ClO 3 ClO 4 Name Hypochlorite Chlorite Chlorate Perchlorate Naming compounds with polyatomic ions Positive charge species on left Negative charge species on right Use parenthesis as needed Formula BaSO 4 Ca(NO 3 ) Ca(NO ) Fe(NO 3 ) Ions Ba + and SO 4 - Ca + and NO 3 Ca + and NO Fe + and NO 3 Name Barium sulfate Calcium nitrate Calcium nitrite Iron (II) nitrate or ferrous nitrate Hydrate Compound with water chemically attached to its ions Two types: Hydro Acids: Hydro + Halogen name + ic HCl hydrochloric acid HF hydrofluoric acid Oxoacids: polyatomic ion + acid HNO 3 nitric acid, from nitrate HNO nitrous acid, from nitrite 7