Chapter 6 Chemical Bonding
I. Formation of Ions
A. When an atom gains or loses electrons
B. Ion: An atom that has a net positive or negative electric charge
1. Represented by a + or a –
2. Example:
3. Anion: An ion with a negative charge
4. Cation: An ion with a positive charge
II. Chemical bond: The force that holds atoms or ions together as a unit
A. Ionic bond: Achieve stable electron configurations through the transfer of electrons
between atoms
1. Octet: When the outer energy level is filled with 8 electrons (8 valence
electrons)
B. Ionic Compounds: Compounds that contain ionic bonds (Metal and a non-metal)
1. Chemical formula: A notation that shows what elements a compound
contains and
the ratio of the atoms or ions of these elements in the compound
-NaCl, MgCl2
III. Naming Compounds and Writing Formulas
A. Binary ionic compounds (metal and non-metal)
a. Name of cation followed by name of anion
1. Change suffix on non-metal to “ide”
B. Stock System
a. Many Transition Metals with multiple ions
b. Roman numeral indicates charge
C. Polyatomic ion: A covalently bonded group of atoms that has a positive or
negative charge and acts as a unit
Ex:
D. Steps for writing formulas for Ionic compounds
1. Write symbol of cation
2. Write symbol of anion
3. Determine charges
4. Cross over
5. Check using algebra
Ex: Calcium Chloride
Barium fluoride
Lithium Oxide
SrI2
K2S
Fr2O
Examples using Polyatomic Ions:
Magnesium Nitrate
Calcium Sulfate
Beryllium Phosphate
Aluminum hydroxide
Li2CO3
BaMnO4
Examples Using the Stock System (Roman Numerals)
Iron II chloride
Tin IV Chloride
Copper I Sulfate
Nickel I Carbonate
FeSO4
NiC2H3O2
Pb(NO3)2
B. Crystal lattice: Set of attractions that keeps ions in fixed positions in a rigid framework
Ex: NaCl; Each Cl- is surrounded by 6 Na+
1. Crystals: Solids whose particles are arranged in a lattice structure
2. Properties of Ionic compounds
a. High melting point (NaCl 801oC)
b. Good conductor of electricity when melted (molten state)
c. Shatter when struck
C. Ionization energy: The amount of energy used to remove an electron form an atom
a. Increases across period
b. Decreases down group
II. Covalent Bonding (2 non-metals)
A. Covalent Bonds: A chemical bond in which two atoms share a pair of valence
electrons
1. Sharing electrons
a. Achieve a stable electron configuration
2. Molecule: A neutral group of atoms that are joined together by one or
more covalent bonds (2 non-metals)
3. Diatomic molecules: Two atoms
a. H2, N2, O2, F2, Cl2, Br2, I2
4. Multiple covalent bonds
a. Double
b. Triple
5. Polar Covalent bond: Covalent bond in which electrons are not shared
equally
B. Polar and Non-polar molecules
1. Polar (H2O and HCl)
a. High boiling points
b. Surface tension
c. High attraction for one another
2. Non-polar Molecules (CO2)
a. Low boiling points
b. Low surface tension
c. Low attraction for one another
III. Naming Molecular Compounds (Between 2 non-metals)
A. Write most metallic atom first (left side of table)
B. 2nd element ends in “ide”
C. Greek prefixes indicate number of atoms
a. If there is only one of the first element do not use “mono”
Covalently Bonded atoms; molecules (2 non-metals)
Examples:
P5O10
N3Cl7
C4Cl8
Diphosphorus tetrafluoride
Nitrogen dioxide
Hexasilicon triphosphide
IV. The Structure of Metals
A. Metallic Bond: The attraction between a metal cation and the shared electrons that
surround it
1. Electron sea model
B. Alloy: A mixture of two or more elements, at least one of which is a metal
1. Copper alloys
a. Bronze; copper and tin
b. Brass; copper and zinc
2. Steel Alloys; iron and carbon
3. Aluminum Alloys: Airplanes (Cu, Mn)
Chapter 7 Chemical Reactions
I. Chemical Equations: A representation of a chemical reaction in which the reactants and
products are expressed as formulas
1. Reactants: The substances that undergo change
a. Left side of the equation
2. Products: The new substances that formed as a result of that change
a. Right side of the equation
II. The Law of the Conservation of Mass: Mass is neither created nor destroyed in a
chemical reaction
1. Lavoisier (1743-1794)
III. Balancing Equations:
1. Coefficients: The numbers that appear before the formulas
2. First balance the atoms that appear only once on each side of the equation
3. Balance polyatomic ions that appear on both sides as single units
4. Balance H and O atoms last
5. Count atoms to be sure that equation is balanced
IV. Types of Reactions
A. Classification Reactions
1. Synthesis: A reaction in which two or more substances react to form a single
substance
A + B  AB
2. Decomposition: a reaction in which a compound breaks down into two or more simpler
substances
AB  A + B
3. Single Replacement: A reaction in which one element takes the place of another element
in a compound
A + BC B + AC
4. Double Replacement: A reaction in which two different compounds exchange positive
ions and form two new compounds
AB + CD AD + CB
5. Combustion: A reaction in which a substance reacts rapidly with oxygen, often
producing heat and light
CH4 + O2  H2O + CO2
V. Energy Changes in Reactions
A. Chemical Energy: The energy stored in chemical bonds of a substance
1. Breaking of bonds in reactants
2. Forming of bonds in products
Ex: Lighting a gas grill
B. Exothermic and Endothermic reactions
1. Exothermic reaction: A chemical reaction that releases energy to its
surroundings
a. Gets warmer
C. Endothermic reaction: A chemical reaction that absorbs energy from its
surroundings
a. Gets colder
D. Chemical change: occurs when a substance reacts and forms one or more new
substances
1. Four types of evidence for a chemical change
a. Color change
1. Silver tarnishing
b. Production of gas (bubbles)
c. The formation of a precipitate
1. Precipitate: A solid that forms and separates from a liquid mixture
d. Energy change (Change in temp)
1. Release or absorb heat (temperature change)
2. Production of light