Lewis Dot Structures

l Valence electrons are represented as dots around the element symbol

– X’s and O’x can also be used

l Chemical Bonds: are the forces that link atoms together

– The electron configuration determines how atoms bond

l Ionic Compound: a compound that is composed entirely of ions

– Cations: positively charged ions

l Metals form cations

– Anions: negatively charged ions

l Nonmetals form anions

l Almost any combination of anions and cations can form an ionic bond

Oxidation Numbers

l Show the combining capacity of an element.

l Tells how many electrons an atom gains or loses when forming a compound.

l Remember the Octet Rule: atoms gain or lose electrons to form a full set of valence electrons

Group 1A 2A 3A 4A 5A 6A 7A

Oxidation # +1 +2 +3 +/- 4 -3 -2 -1

Monatomic Ions

l Monatomic Ions: one atom ions

^– Examples: Li⁺¹, O^-2

– Elements in Group 1A to 7A always form the same type of ion

– Monatomic cations have the same name as the element

l Ca = Calcium

l Ca⁺² = Calcium Ion

– To name a monatomic anion replace the suffix of the element name with suffix “ide”

l Cl = Chlorine Cl^-1 = Chloride ion

l Br = Bromine Br ^–1 = Bromide ion

l O = Oxygen O^-2 = Oxide ion

Binary Ionic Compounds: contain ions of only 2 elements

l Empirical formulas: denote the ratio of ions in a compound

– Element Symbols indicate ions or elements in compound and subscripts indicate the ratio of atoms or ions

– Formulas for ionic compounds usually give the smallest whole number ratio of ions

– If no subscript assume 1 atom or ions

Naming Binary Ionic Compounds

l Name the cation followed by the anion

– Don’t forget the “ide” ending for the anion

l CaO = Calcium Oxide

l MgCl₂ = Magnesium Chloride

l Al₂S₃ = Aluminum Sulfide

Criss Cross Method to write formulas for Binary Ionic Compounds

More Criss Cross Examples

1. Barium Sulfide Ba S

^2. Ba⁺² S^-2

_3. Ba₂S₂ - simplify

_4. Ba₁S₁ – remove 1s

5. BaS

Calcium Nitride

1. Ca N

^2. Ca⁺² N^-3

_3. Ca₃N₂

Transition Metals:

· do not follow the octet rule and commonly form more than one monatomic cation

· Use Roman Numerals to distinguish the cations of transition metals

l Fe⁺² = Iron (II) Fe⁺³ = Iron (III)

l Cu⁺¹ = Copper (I) Cu⁺² = Copper (II)

Writing Formulas for Transition Metals

– Iron (II) Sulfide

– Fe S

^– Fe⁺² S^-2

_– Fe₂S₂ - simplify

_– Fe₁S₁ – remove 1s

– FeS

– Iron (III) Sulfide

– Fe S

^– Fe⁺³ S^-2

– Fe₂S₃

Polyatomic Ions:ions formed from a group of elements

Writing Formulas with Polyatomic ions

– Magnesium Nitrate

_– Mg NO₃

^– Mg⁺² NO₃^-1

– Mg₁ (NO₃)₂

Note: Use parantheses around polyatomic ion to show subscript is for the entire ion

_– Mg (NO₃)₂

– Ammonium Carbonate

_– NH₄ CO₃

^– NH₄⁺¹ CO₃^-2

– (NH₄) ₂ (CO₃)₁

_– (NH₄)₂ CO₃

No parantheses needed around

CO₃ because there is only one CO₃ ion

Covalent Bonds

· A Covalent Bond is formed by a shared pair of electrons between 2 atoms

· Covalent Bonds usually form between 2 nonmetals

· Molecule: a group of atoms that are united by covalent bonds

· Molecular Substance: a substance that is made of molecules

Molecular Formulas

l Molecular Formula: tells how many atoms are in a single molecule of a compound

– F₂: contains 2 F atoms

– C₆H₁₂O₆: contains 6 C, 12 H, and 6 O atoms

– You can write an empirical formula by simplifying molecular formula to the simplest whole numbers:

l C₆H₁₂O₆ Þ CH₂O

l C₁₂H₂₄O₁₂ Þ CH₂O

l C₁₈H₃₆O₁₈ Þ CH₂O

– Many different molecular substances can have the same empirical formula

Describing Covalent Bonds

· The Octet Rule Applies: atoms share electrons to obtain 8 valence electrons

· The shared pair completes the octet for both atoms

· H only needs 2 electrons to fill its outer shell

· The group number of the atom determines the number of shared pairs that an atom needs

Group 4A 5A 6A 7A

Desired # of shared pairs 4 3 2 1

Naming Molecular Compounds

l Prefixes are used to indicate the subscripts in the molecular formula:

Mono = 1

di = 2

tri = 3

tetra = 4

penta = 5

hexa = 6

hepta = 7

octa = 8

nona = 9

deca = 10

l Write the name of the least electronegative atom first

l If there is more than one atom of that element include a prefix

l Write the prefix of the second element followed by the root of the element name with an “ide” ending

CO₂ Carbon Dioxide

PCl₃ Phosphorus Trichloride

N₂O₅ Pentoxide

Lewis Dot Structures and
Covalent Bonds

l Draw Lewis Structures for each atom

l Count the total number of valence electrons

l Place the atom that requires the most shared pairs in the center

l Draw one covalent bond to connect the central atom to each of the other atoms

l Check to see that all atoms have octet and total valence electrons is the same as step 2

If not, add double then triple bonds until conditions are met

Lewis Dot Example: NCl₃

l Lewis Structures

l Total Valence e^- = 26

l N needs 3 shared pairs, Cl needs 1, N is central

l Draw one shared pair between N and each Cl

l All atoms have a full octet and total e^- = 26

Lewis Dot Example: CO₂

l Draw Lewis Structures

l Total Valence e^- = 16

l C needs 4 shared pairs, O needs 2, C is central

l Draw one shared pair between C and each O

l Atoms do not have a full octet

l Move one e- from C and one from O to make a double bond – Octets are still not complete

l Draw another double bond between C and the other O

l Octets are complete, total e- = 16

Lewis Dot Example: C₂H₂

l Draw Lewis Structures

l Total Valence e^- = 10

l C needs 4 shared pairs, H needs 1, Both C’s are central

l Draw one shared pair between each C and 1 shared pair between each C and 1 H

l H atoms are full, both C atoms still need 2 more shared pairs

l Move 2e- from each C to form a triple bond

l Octets are complete, total e- = 10

Dashes can be used to represent bonds

Exceptions to the Octet Rule

l Atoms with less than an octet: Many compounds of Boron do not have a complete octet

l Atoms with more than an octet: Atoms with d electrons sometimes form bonds that give more than an octet

l Molecules with an odd number of electrons: compounds with an odd number of electrons cannot follow the octet rule

– Such compounds are usually unstable- they easily react to form more stable compounds

Polarity

l Polar compounds: are formed when 2 atoms with different electronegativities form a covalent bond

– The atoms do not share electrons equally

– The electrons are more strongly attracted to the most electronegative atom

l Nonpolar covalent compounds: Bonds between atoms with similar or equal electronegativities

Bond type by Electronegativity

Water is Polar

Properties of Ionic Compounds

· Electrically neutral: the electrical charges of the anion and cation must cancel each other

· Typically hard, high melting point, and brittle

· Most Dissolve in water: ions separate and move freely in water; thus solutions of ionic compounds conduct electricity even though ionic solids do not!

· Insoluble in nonpolar, organic solvents

Properties of Covalent Compounds

l There are several forces holding molecular compounds together. The relative strength of these forces determines properties. Generally covalent compunds:

– Are relatively soft with relatively low melting points

– can be liquid, solid, or gas at room temperature (RT)

– typically do not conduct heat and electricity

l Nonpolar covalent compounds are not soluble in water

l Polar covalent compounds do dissolve in water

l Carbon forms more compounds than any other element. Compounds of hydrogen and carbon are called organic compounds. There are so many organic compounds that such compounds have very diverse properties

Properties of Metallic Compounds

· Metallic bonds hold atoms together in a metallic substance. These bonds are very strong thus metals:

– Have high melting points (most are solids at RT)

– Are good conductors

– Are malleable and ductile