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