Unit 4

Physicalchange 

Involve changes in IMFs

Evidence of chemical change

Production of heat or light, formation of gad, formation of precipitate (solid), color change

Limiting Reactant

The reactance used up first; typically the greater coefficient

Percent yield

Actual Yield / Theoretical yield × 100

Electrolyte

Forms ions in solution and conducts electricity

Non electrolyte

A substance that does not form ions in solution or conduct electricity; molecular compounds

Strong electrolyte 

100% dissociation in water; stong acids, strong bases, ionic compounds

Weak electrolyte 

Not completely dissolved; weak acids and weak bases

Double Displacement reactions

Exchange ions to form a precipitate

Solubility guidelines

Salts containing Alkali Metals, Nitrates ( NO3 -1), Ammonium (NH4 +1) and strong electrolytes, group 1 A metals

Insoluble compounds

Carbonates (CO3 2-) phosphates (PO 3-) chromate (CrO4 2- ) sulfides (S 2-) hydroxide (OH-)

Strong acids

So (H2SO4) I (HI) BRought (HBr) No (HNO3) Clean (HCl) Clothes (HCl3) (HCl4). Conjugate bases are weak

Strong Bases

Metal with hydroxide ion, Alkali Metals, Calcium, strontium, barium. 

Arrhenuis Acids

Produces H+ in water

Arrhenius base

Produces OH- in water

Bronsted Acid

Proton donor

Bronsted base

Proton acceptor

Factors affecting Acid strength 

More electronegativity more acidic, more oxygen more acidic

Dilution

M1V1"=M2V2

Steps of Half Reaction Method 

Step 1: Identify what’s oxidized and what’s reduced

• Oxidation = loss of electrons (the oxidation number increases).
  
  

• Reduction = gain of electrons (the oxidation number decreases).

Write down the oxidation numbers for each atom in the reaction to see which ones changes

Step 2: Write two half-reactions

Split the full reaction into:

• One oxidation half-reaction
  
  

• One reduction half-reaction
  
  

Each should show only one element changing oxidation state, along with any atoms directly bonded to it.

 Step 3: Balance all atoms except oxygen and hydrogen

Balance the elements that aren’t O or H first on both sides of each half-reaction.

Step 4: Balance oxygen atoms

• In acidic solution, add H₂O to the side that needs more oxygen.
  
  

• In basic solution, also add H₂O, but you’ll fix H later with OH⁻ instead of H⁺.
  
  

 Step 5: Balance hydrogen atoms

• In acidic solution, add H⁺ where needed.
  
  

• In basic solution, add OH⁻ to balance H, and combine any resulting H⁺ and OH⁻ into H₂O.
  
  

 Step 6: Balance charge by adding electrons (e⁻)

• Add electrons to the more positive side (or less negative side) of each half-reaction so that both sides have the same net charge.
  
  

• The number of electrons lost in oxidation must equal the number gained in reduction.
  
  

Step 7: Equalize electrons between the two half-reactions

Multiply one or both half-reactions by appropriate factors so the electrons cancel when added together.

Step 8: Add the half-reactions together

Combine the two balanced half-reactions, canceling:

• Electrons
  
  

• Any identical species (like H₂O or H⁺) that appear on both sides.

What is the oxidation number of free elements and diatomic atoms?

0

What is the oxidation number of oxygen?

-2

What is the oxidation number of H2O2 and O2 2- (peroxide)? 

-1

What is the oxidation number of hydrogen?

+1

What is the oxidation number of hydrogen when bonded to metals

-1

What is the oxidation number of fluorine?

-1

What is the oxidation number of Group 1 and group 2 metals?

+1, +2