Ionic Compounds Ionic Compounds Ionic Compounds Ionic Compounds Ch.6 & 7

Ionic CompoundsIonic CompoundsIonic CompoundsIonic Compounds

Ch.6 & 7Ch.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. Na+ + Cl- NaCl2. Ca2+ + Cl- CaCl23. Mg2+ + O2- MgO4. Fe3+ + S2- ______5. 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. NaCl ____________________2. CaCl2 ____________________

3. MgO ____________________4. Fe2S3 ____________________

5. BeI2 ____________________

Ionic Compound Characteristics

• A formula unit is the smallest whole number ratio of cation to anion

ex. NaCl• 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 CompoundCharacteristics

• 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 IonsNames 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: ClO4

-, ClO3-, ClO2

-, ClO-, Cl- (p.226)

Polyatomic Ions to Study

• Phosphate PO4-3

• Chlorate ClO3-

• Nitrate NO3-

• Sulfate SO3-2

• Carbonate CO3-2

Oxidation NumbersDefinition: 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 KMnO4 K = ___ Mn = ___ O =

___

Oxidation Number Practice

• NaH Na = _____ H = _____

• KClO2 K = ____ Cl = ____ O = ____

• KClO3 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 MoleculesCovalent MoleculesCovalent MoleculesCovalent Molecules

Ch. 6 & 7Ch. 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 (0-0.3)– Polar-covalent (0.3-1.7)– Ionic (over 1.7)

Lewis Structures• Definition – formulas in which

atomic symbols represent nuclei and inner-shell 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 Prefixesmono 1di 2tri 3tetra 4penta 5

hexa 6hepta 7octa 8nona 9deca 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.

PracticeName the following covalent

compounds:1. CO2 _________________

2. P2O5 _________________

3. OF3 _________________

4. SO2 _________________

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

PracticeName the following acids:1. HCl ________________2. HClO4 ________________

3. H2SO3 ________________

4. HI ________________