Ionic Compounds Ch 6 7 Ionic Bonding strongest

Ionic Compounds Ch. 6 & 7

Ionic Bonding – strongest of all bonds • definiton – results from the electrical attraction between cations (+) and anions (-) • “ions” means charges in our case • Calculated by the difference between electronegativity values (over 1. 7) • Figure 2 -- p. 176

Ionic Bonding • Positive ions are formed from the metal elements when they lose electrons • Negative ions are formed from the nonmetal elements when they gain electrons • The two, (+) and (-) come together to form a new compound

Ionic Compounds • definition – composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal • In a neutral compound, the net charge will be zero • Balance the charges from cations and anions until the compound is neutral

Forming Ionic Compounds 1. 2. 3. 4. 5. Na+ + Cl- Na. Cl Ca 2+ + Cl- Ca. Cl 2 Mg 2+ + O 2 - Mg. O Fe 3+ + S 2 - ______ beryllium + iodine ______

Naming Ionic Compounds • Name the metal first (complete name) • Name the root of the nonmetal the change the ending to –ide • If there is a polyatomic (more than one atom carrying one charge) use the name as it is • Transition metals have a Roman numeral after it to tell its charge

Naming Ionic Compounds 1. 2. 3. 4. 5. Na. Cl Ca. Cl 2 Mg. O Fe 2 S 3 Be. I 2 ____________________ __________

Ionic Compound Characteristics • A formula unit is the smallest whole number ratio of cation to anion ex. Na. Cl • Ionic compounds mostly have a crystalline structure • Lattice energy is the energy released when one mole of an ionic crystalline compound is formed from gaseous ions

Ionic Compound Characteristics • Generally have a high melting point and boiling point • Hard but brittle • Not electrical conductors as solids • Most can dissolve in water • Strongest bond

Polyatomic Ions • Definition: a charged group of covalently bonded atoms • Lewis structures can be used to show the electron placement and where the charge of the ion comes from (p. 194)

Polyatomic Ions Names have a system: – ending -ate means the highest number of oxygen bonded for that nonmetal (sulfate, chlorate, iodate, etc) – ending –ite means one less oxygen than ate (sulfite, chlorite, iodite, etc) – Chlorine is a special case: Cl. O 4 -, Cl. O 3 -, Cl. O 2 -, Cl. O-, Cl- (p. 226)

Polyatomic Ions to Study • • • Phosphate Chlorate Nitrate Sulfate Carbonate PO 4 -3 Cl. O 3 NO 3 SO 3 -2 CO 3 -2

Oxidation Numbers Definition: the number of electrons that must be added to or removed from an atom in a combined state to convert the atom into the elemental form Example: H+ + Cl- HCl KMn. O 4 K = ___ Mn = ___ O = ___

Oxidation Number Practice • Na. H Na = _____ H = _____ • KCl. O 2 K = ____ Cl = ____ O = ____ • KCl. O 3 K = ____ Cl = ____ O = ____

Metallic Bonding • Definition: the chemical bonding that results from the attraction between two metal atoms and the surrounding sea of electrons • Properties include: – High electrical and thermal conductivity – Strong absorber and reflector of light – Conforms to shape easily • Malleability – hammered into sheets • Ductility – drawn into wires

Covalent Molecules Ch. 6 & 7

Covalent Bonding • Definition: bond formed by the sharing of electron pairs between two nonmetals • There are two types of covalent bonds: – Nonpolar: electrons are shared equally – Polar-covalent: unequally shared electrons

Electronegativities • Every element on the PT has an electronegativity value • Those values are subtracted to see whether is: – Nonpolar – Polar-covalent – Ionic (0 -0. 3) (0. 3 -1. 7) (over 1. 7)

Lewis Structures • Definition – formulas in which atomic symbols represent nuclei and innershell electrons, dot-pairs in covalent bonds • Sample Problem C - p. 185 • Practice - p. 186

Molecular Geometry • VSEPR theory – repulsion between the sets of valence electrons surrounding an atom causes these sets to be oriented as far as possible • Use VSEPR with Lewis structures to come up with shapes • Shapes are: – Linear, trigonal planar, tetrahedral, bent, and trigonal pyramidal, etc. (p. 200) • Sample Problem F – p. 201

Intermolecular Forces • Definition: forces in between molecules • Weaker than bond strength • Types: – Dipole-dipole: strongest in polar covalent – Hydrogen bonding: hydrogen bonded to highly electronegative atom – London dispersion forces: weakest force due to motion of atoms in compound

Naming Covalent Molecules • Use prefixes listed on p. 228 • Prefixes tell how many atoms of each nonmetal make up the molecule • The nonmetal farthest left on PT is written first, then the most electronegative atom (farthest right on the PT) • Mono does not appear on the first atom to notate one, it is understood • Omit vowels on the prefix if there is a vowel on the element • Second atom ends in -ide

Covalent Prefixes mono di tri tetra penta 1 2 3 4 5 hexa hepta octa nona deca 6 7 8 9 10 Two Rules of Thumb: 1. If two vowels end up next to each other, the vowel on the prefix will be deleted. 2. Mono is never used for the first element.

Practice Name the following covalent compounds: 1. CO 2 _________ 2. P 2 O 5 _________ 3. OF 3 _________ 4. SO 2 _________

Acids • Formed from H+ + a nonmetal or a polyatomic • If H+ is bonded to a nonmetal, use the prefix hydro- + the root of the nonmetal + ending with –ic acid • If H+ is bonded to a polyatomic, do NOT use the prefix hydro! Use the root of the polyatomic and use ending from –ate to –ic or –ite to –ous acid • Section Review p. 231 #4 f-h

Practice Name the following acids: 1. HCl ________ 2. HCl. O 4 ________ 3. H 2 SO 3 ________ 4. HI ________