Chapter 7 Chemistry
Monoatomic Natural State
*All elements not classified in another group are monoatomic and everything in group 8 on the periodic table in monoatomic.* This means one atom per molecule, which in simple terms mean it is written with a subscript of 1. The one is not needed to be written.
Diatomic Natural State
*Group 7 elements* and *Hydrogen, Oxygen* and *Nitrogen* are diatomic. This means they are written with a subscript of 2
Tetratomic Natural State
*Phosphorus* is the only known element to be tetratomic. It is written with a subscript of 4
Octatomic Natural State
*Sulfur* is the only known octatomic element. it is written with a subscript of 8.
*Subscripts in Acids*
*the charge on the H matches the charge the ion would have if it were an ion* Even though it is not an ion you have to have the right amount of H+1's to balance the negative ion in the formula before you can name the acid.
Oxidation numbers rule 1
1) oxidation number for an atom in an element is 0
if the anion ends in *-ate*
1) take the root of the anion 2) add the suffix -ic 3) add the word acid *-ate changes to -ic*
If the anion ends in *-ite*
1) take the root of the anion 2) add the suffix -ous 3) add the word acid *-ite changes to -ous*
If the anion in the acid ends in *-ide*
1) take the root of the anion 2) add the prefix hydro- 3) add the suffix -ic 4)add the word acid *-ide changes to hydro- -ic*
Steps for balancing equations by using oxidation numbers
1)Assign oxidation numbers to all elements 2)Identify oxidized and reduced elements 3)Multiply the changes in oxidation numbers by small integers to equalize the decrease
Oxidation numbers rule 2
2) oxidation number for a monatomic ion is the same as its charge
Metallic elements combine with nonmetallic elements to yield ionic compounds (ionic compounds are called salts)
2Al+3Cl₂→2AlCl₃ 2K+I₂→2KI
one compound plus oxygen forms oxides
2C₂H₆+7O₂→4CO₂+6H₂O
Oxides of less active metals decompose to form the metal and oxide when heated
2HgO→2Hg +O₂
Electrolysis (adding electricity) of compounds decomposes into two or more elements or compouds
2H₂O→2H₂+O₂ 2NaCl→2Na+Cl₂
Metallic (any positive ion) chlorates decompose into metallic chlorides and oxygen when heated
2KClO₃→2KCl+3O₂
Peroxides decompose into oxides and oxygen
2K₂O₂→2K₂O + O₂ 2H₂O₂→2H₂O + O₂
Oxygen combines with most elements to form oxides
2Mg+O₂→2MgO C+O₂ → CO₂
Very active metals (only the first 5) can replace one (AND ONLY ONE) of the hydrogens from the stable compound water.
2Na+2H₂O→2NaOH+H₂ *a metal hydroxide is always one of the products*
Oxidation numbers rule 3
3)sum of oxidation in all numbers of a neutral compound is zero
Oxidation numbers rule 4
4)the sum of oxidation numbers in an ion is equal to the charge of the ion
Oxidation numbers rule 5
5)In compounds, fluorine is always assigned an oxidation number of -1 (the most electronegative element in a compound always has a negative oxidation number.)
Oxidation numbers rule 6
6) Hydrogen's oxidation number will be - +1 when bonded to a nonmetal (HCl) - -1 when bonded to a metal (NaH)
Oxidation numbers rule 7
7) Oxygen usually has an oxidation number of -2. Exceptions: in peroxides, oxygen will be -1 and when combined with only F, it will be +2 and in O2 it will be 0
Oxidation numbers rule 8
8) Halogens usually have an oxidation number of -1.
Precipitate
A solid product in a reaction
Precipitation Reaction- one of the possible products must be a precipitate or no reaction occurs.
AgNO₃+HCl→AgCl+HNO₃ *AgCl is a precipitate*
Methane
CH₄
Bases (metallic hydroxides) decompose into water and metallic oxides
Ca(OH)₂ → CaO + H₂O
Metallic Carbonates decompose into metallic oxides and carbon dioxide when heated
CaCO₃→CaO+CO₂
Oxides of active metals combine with water to form compounds called bases (Metallic hydroxides) Remember: WATER IS AN OXIDE
CaO+H₂O→Ca(OH)₂ Na₂O+H₂O→2NaOH
More active nonmetals replace less active metals from their compounds
Cl₂+2NaI→2NaCl+I₂
More active metals replace less active metals from their compounds. (hydrogen can act like a metal)
Cu+2AgNO₃→Cu(NO₃)₂+2Ag Zn+H₂SO₄→ZnSO₄+H₂
Hydrated compounds decompose to yield the anhydrous compound and water when heated
CuSO₄ × 5H₂O → CuSO₄ + +5H₂O
Decane
C₁₀H₂₂
Ethane
C₂H₆
Propane
C₃H₈
Butane
C₄H₁₀
Pentane
C₅H₁₂
Hexane
C₆H₁₄
Heptane
C₇H₁₆
Octane
C₈H₁₈
Nonane
C₉H₂₀
Acid Base Reaction-an acid plus a base yields an ionic compound (a salt) and water (a molecular compound).
HCl+NaOH→NaCl + H₂O 2H₃PO₄+3Ca(OH)₂→Ca₃(PO₄)₂+6H₂O
Acids (hydrogen combined with some negative ion) decompose into water and nonmetallic oxides
H₂CO₃→H₂O + CO₂ 2H₃PO₄→P₂O₅ + 3 H₂O
Molecular reaction- the reaction occurs of one of the products is a molecular compound
H₂SO₄+Na₂O→Na₂SO₄ + H₂O (H₂O is molecular) CaCO₃+2HCl→CO₂+H₂O+CaCl₂ (CO₂ is molecular)
Oxides of metals combine with oxides of nonmetals to yield ionic compounds. These compounds will contain an -ate ion.
Na₂O+SO₃→Na₂SO₄
Oxides of nonmetals combine with water to form compounds called acids. These acids will be made from an -ate ion.
P₂O₅+3H₂O→2H₃PO₄ SO₃+H₂O→H₂SO₄
Nonmetallic elements tend to combine with each other to form covalent (molecular) compounds
S₈+8 F₂→8SF₂ P₄+6Cl₂→4PCl₃
Coefficeient
The number at the beginning of the compound/element/symbol
Subscript
The number in the lower right corner from the element letter
Reducing Agent
What is being oxidized in the equation. The oxidation of this causes the reduction of another part of the equation. Provides electrons when oxidized
Oxidizing Agent
What is being reduced in the equation. The reduction of this causes the oxidation of another part of the equation. Takes the electrons from the reducing agent.
redox reaction
a process where electrons are transferred from one substance to another
SALT
an ionic compound
all organic compounds contain
carbon and hydrogen
formula for organic compounds
carbon, then hydrogen, then any other elements
Saturated hydrocarbon
hydrocarbon made of all single covalent bonds
unsaturated hydrocarbon
hydrocarbon that contains double bonds
Acid
hydrogen combined with either oxygen or something considered a negative ion. *Acids are not ionic compounds*
net ionic equation
includes only the participating ions in the reaction
Ionic Reactions
ions exchange respective positive and negative ions *reactants are 2 compounds AB + CD -> AD + CB*
Combustion Reaction
most compounds combine with oxygen when they burn to form oxides *1 compound plus oxygen AB + O2 -> oxide of A + oxide of B*
Decomposition Reaction
one compound decomposes into two or more elements or compounds *only one reactant AB -> A + B*
Replacement Reaction
one element replaces another in a compound- more active metals replace less active metals OR more active metals replace less active nonmetals *A + BC -> AC + B*
Hydrocarbon
organic compounds that only contain carbon and hydrogen
when the oxidation number increases
oxidation (loss of electrons) has occurred
Oxidation
oxidation is loss of electrons (oil)
when the oxidation number decreases
reduction (gain of electrons) has occurred
Reduction
reduction is gain of electrons (rig)
Chemical Equation
representation of a chemical reaction A + B -> AB
Alkanes
saturated hydrocarbons that (with the exception of methane) contain carbon atoms linked in chains
molecular equation
shows the complete formula of all reactants and products and the physical state of each compound.
Reactants
the left side of a chemical equation
Products
the right side of a chemical equation
Natural State of Elements
the state in which an element exists when it is not in a compound
Organic Chemistry
the study of carbon compounds
Composition Reaction
two or more elements unite to form one new compound. There is only one product. *A + B -> AB*
complete ionic equation
what the equation looks like when the soluble compounds are separated into ions