Balanced equation: Same number of each type of atom on both sides

Five Main Types of Reactions

1. Synthesis (Combination) Reactions

Definition: Two or more substances combine to form a single product

General form:

$A + B \rightarrow AB$

Examples:

Metal + Nonmetal → Ionic compound:

$2Na\text{(s)} + Cl2\text{(g)} \rightarrow 2NaCl\text{(s)}$

$2Mg\text{(s)} + O2\text{(g)} \rightarrow 2MgO\text{(s)}$

Nonmetal + Nonmetal → Molecular compound:

$C\text{(s)} + O2\text{(g)} \rightarrow CO2\text{(g)}$

$N2\text{(g)} + 3H2\text{(g)} \rightarrow 2NH3\text{(g)}$ (Haber process)

Metal oxide + Water → Metal hydroxide:

$CaO\text{(s)} + H2O\text{(l)} \rightarrow Ca(OH)2\text{(aq)}$

$Na2O\text{(s)} + H2O\text{(l)} \rightarrow 2NaOH\text{(aq)}$

Nonmetal oxide + Water → Acid:

$CO2\text{(g)} + H2O\text{(l)} \rightarrow H2CO3\text{(aq)}$ (carbonic acid)

$SO3\text{(g)} + H2O\text{(l)} \rightarrow H2SO4\text{(aq)}$ (sulfuric acid)

Characteristics:

• Multiple reactants → One product
• Often exothermic (releases energy)
• Opposite of decomposition

2. Decomposition Reactions

Definition: Single compound breaks down into two or more simpler substances

General form:

$AB \rightarrow A + B$

Examples:

Heat decomposition:

$2HgO\text{(s)} \xrightarrow{\Delta} 2Hg\text{(l)} + O2\text{(g)}$

$CaCO3\text{(s)} \xrightarrow{\Delta} CaO\text{(s)} + CO2\text{(g)}$

Electrolysis (electrical energy):

$2H2O\text{(l)} \rightarrow[\text{electricity] 2H2(g) + O2(g)}$

$2NaCl\text{(l)} \rightarrow[\text{electricity] 2Na(l) + Cl2(g)}$

Catalytic decomposition:

$2H2O2\text{(aq)} \xrightarrow{MnO_2} 2H2O\text{(l)} + O2\text{(g)}$

Binary compound decomposition:

$2KClO3\text{(s)} \xrightarrow{\Delta} 2KCl\text{(s)} + 3O2\text{(g)}$

Characteristics:

• One reactant → Multiple products
• Often requires energy input (endothermic)
• Need heat, electricity, or light
• Opposite of synthesis

3. Single Replacement (Displacement) Reactions

Definition: One element replaces another in a compound

General forms:

Metal replaces metal:

$A + BC \rightarrow AC + B$

Nonmetal replaces nonmetal:

$A + BC \rightarrow BA + C$

Examples:

More reactive metal displaces less reactive:

$Zn\text{(s)} + CuSO4\text{(aq)} \rightarrow ZnSO4\text{(aq)} + Cu\text{(s)}$

$Mg\text{(s)} + 2HCl\text{(aq)} \rightarrow MgCl2\text{(aq)} + H2\text{(g)}$

$2Al\text{(s)} + 3CuCl2\text{(aq)} \rightarrow 2AlCl3\text{(aq)} + 3Cu\text{(s)}$

More reactive halogen displaces less reactive:

$Cl2\text{(g)} + 2NaBr\text{(aq)} \rightarrow 2NaCl\text{(aq)} + Br2\text{(l)}$

$Br2\text{(l)} + 2KI\text{(aq)} \rightarrow 2KBr\text{(aq)} + I2\text{(s)}$

Predicting if reaction occurs:

Use Activity Series (reactivity order)

Metal Activity Series (most → least reactive):

• Li > K > Ba > Ca > Na > Mg > Al > Zn > Fe > Ni > Sn > Pb > H > Cu > Ag > Au

Rule: Metal higher in series displaces metal lower in series

Halogen Activity Series:

• F₂ > Cl₂ > Br₂ > I₂

Rule: Halogen higher in series displaces halogen lower in series

Will this reaction occur?

$Cu\text{(s)} + 2AgNO3\text{(aq)} \rightarrow ?$

Check: Cu is above Ag in activity series → YES, reaction occurs

$Cu\text{(s)} + 2AgNO3\text{(aq)} \rightarrow Cu(NO3)2\text{(aq)} + 2Ag\text{(s)}$

$Ag\text{(s)} + CuSO4\text{(aq)} \rightarrow ?$

Check: Ag is below Cu in activity series → NO, no reaction

$Ag\text{(s)} + CuSO4\text{(aq)} \rightarrow \text{No Reaction}$

4. Double Replacement (Metathesis) Reactions

Definition: Two compounds exchange partners

General form:

$AB + CD \rightarrow AD + CB$

Cations and anions "switch partners"

Examples:

Precipitation reactions:

$AgNO3\text{(aq)} + NaCl\text{(aq)} \rightarrow AgCl\text{(s)} + NaNO3\text{(aq)}$

$Pb(NO3)2\text{(aq)} + 2KI\text{(aq)} \rightarrow PbI2\text{(s)} + 2KNO3\text{(aq)}$

Acid-base neutralization:

$HCl\text{(aq)} + NaOH\text{(aq)} \rightarrow NaCl\text{(aq)} + H2O\text{(l)}$

$H2SO4\text{(aq)} + 2KOH\text{(aq)} \rightarrow K2SO4\text{(aq)} + 2H2O\text{(l)}$

Gas formation:

$2HCl\text{(aq)} + Na2CO3\text{(aq)} \rightarrow 2NaCl\text{(aq)} + H2O\text{(l)} + CO2\text{(g)}$

$2HCl\text{(aq)} + Na2S\text{(aq)} \rightarrow 2NaCl\text{(aq)} + H2S\text{(g)}$

Predicting if reaction occurs:

Double replacement reactions occur if one of these forms:

1. Precipitate (insoluble solid)
2. Water (in neutralization)
3. Gas (escapes from solution)

Use solubility rules to predict precipitates!

5. Combustion Reactions

Definition: Substance reacts rapidly with oxygen, releasing energy as heat and light

General form (hydrocarbon combustion):

$C_xH_y + O2 \rightarrow CO2 + H2O + \text{energy}$

Complete combustion examples:

$CH4\text{(g)} + 2O2\text{(g)} \rightarrow CO2\text{(g)} + 2H2O\text{(g)}$ (methane)

$C3H8\text{(g)} + 5O2\text{(g)} \rightarrow 3CO2\text{(g)} + 4H2O\text{(g)}$ (propane)

$2C8H18\text{(l)} + 25O2\text{(g)} \rightarrow 16CO2\text{(g)} + 18H2O\text{(g)}$ (octane)

Complete combustion products:

• CO₂ (carbon dioxide)
• H₂O (water)
• Energy (heat and light)

Incomplete combustion (insufficient O₂):

$2CH4\text{(g)} + 3O2\text{(g)} \rightarrow 2CO\text{(g)} + 4H2O\text{(g)}$ (produces CO)

$CH4\text{(g)} + O2\text{(g)} \rightarrow C\text{(s)} + 2H2O\text{(g)}$ (produces soot)

Other combustion reactions:

$S\text{(s)} + O2\text{(g)} \rightarrow SO2\text{(g)}$

$4Fe\text{(s)} + 3O2\text{(g)} \rightarrow 2Fe2O3\text{(s)}$ (rusting)

Characteristics:

• Always involves O₂ as reactant
• Exothermic (releases energy)
• Often produces flame
• Important for energy production

Solubility Rules

Essential for predicting precipitation in double replacement reactions

Soluble Compounds (form aqueous solutions)

1. Group 1 (alkali metals) compounds:

• All Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺ salts are soluble
• Example: NaCl, K₂SO₄, LiNO₃

2. Ammonium (NH₄⁺) compounds:

• All NH₄⁺ salts are soluble
• Example: (NH₄)₂SO₄, NH₄Cl

3. Nitrates (NO₃⁻), acetates (C₂H₃O₂⁻), perchlorates (ClO₄⁻):

• All soluble
• Example: AgNO₃, Ca(NO₃)₂, Pb(C₂H₃O₂)₂

4. Most chlorides (Cl⁻), bromides (Br⁻), iodides (I⁻):

• Soluble EXCEPT with Ag⁺, Pb²⁺, Hg₂²⁺
• AgCl, PbBr₂, Hg₂I₂ are insoluble

5. Most sulfates (SO₄²⁻):

• Soluble EXCEPT with Sr²⁺, Ba²⁺, Pb²⁺
• SrSO₄, BaSO₄, PbSO₄ are insoluble

Insoluble Compounds (form precipitates)

1. Carbonates (CO₃²⁻), phosphates (PO₄³⁻), chromates (CrO₄²⁻):

• Insoluble EXCEPT with Group 1 and NH₄⁺
• CaCO₃, Ag₃PO₄, PbCrO₄ are insoluble
• But: Na₂CO₃, K₃PO₄, (NH₄)₂CrO₄ are soluble

2. Hydroxides (OH⁻):

• Insoluble EXCEPT with Group 1, Ba²⁺, Sr²⁺, Ca²⁺ (slightly)
• Fe(OH)₃, Al(OH)₃, Cu(OH)₂ are insoluble
• But: NaOH, KOH, Ba(OH)₂ are soluble

3. Sulfides (S²⁻):

• Insoluble EXCEPT with Group 1, Group 2, NH₄⁺
• CuS, FeS, PbS are insoluble
• But: Na₂S, (NH₄)₂S are soluble

Using Solubility Rules

Example: Will precipitation occur?

$AgNO3\text{(aq)} + NaCl\text{(aq)} \rightarrow ?$

Step 1: Identify potential products (switch partners)

• Ag⁺ with Cl⁻ → AgCl
• Na⁺ with NO₃⁻ → NaNO₃

Step 2: Check solubility

• AgCl: Cl⁻ compound with Ag⁺ → Insoluble (exception to Cl⁻ rule)
• NaNO₃: Group 1 compound → Soluble

Step 3: Write equation

• AgCl precipitates → (s)
• NaNO₃ stays dissolved → (aq)

$AgNO3\text{(aq)} + NaCl\text{(aq)} \rightarrow AgCl\text{(s)} + NaNO3\text{(aq)}$

Yes, precipitation occurs! (AgCl is white precipitate)

Identifying Reaction Types

Strategy: Look for patterns

Synthesis: Multiple → One

• 2 or more reactants, 1 product
• Example: $2H2 + O2 \rightarrow 2H2O$

Decomposition: One → Multiple

• 1 reactant, 2 or more products
• Often needs energy input
• Example: $2H2O \rightarrow 2H2 + O2$

Single Replacement: Element + Compound → Element + Compound

• Free element replaces element in compound
• Check activity series
• Example: $Zn + CuSO4 \rightarrow ZnSO4 + Cu$

Double Replacement: Compound + Compound → Compound + Compound

• Ions switch partners
• Forms precipitate, water, or gas
• Example: $AgNO3 + NaCl \rightarrow AgCl + NaNO3$

Combustion: Substance + O₂ → CO₂ + H₂O

• Always has O₂ as reactant
• Produces CO₂ and H₂O (if hydrocarbon)
• Exothermic, flame
• Example: $CH4 + 2O2 \rightarrow CO2 + 2H2O$

Note: Some reactions fit multiple categories!

$2H2\text{(g)} + O2\text{(g)} \rightarrow 2H2O\text{(g)}$

This is both synthesis AND combustion!

Driving Forces for Reactions

Why do reactions occur? What makes them "go"?

1. Formation of Precipitate

Ions in solution combine to form insoluble solid

$Pb^{2+(aq) + 2I^-(aq) -> PbI2(s)}$

Driving force: Decrease in solubility (ions prefer solid state)

2. Formation of Water

Neutralization reactions

$H^+\text{(aq)} + OH^-\text{(aq)} \rightarrow H2O\text{(l)}$

Driving force: Formation of very stable H₂O molecule

3. Formation of Gas

Gas escapes from solution

2H^+\text{(aq)} + CO3^{2-(aq) -> H2O(l) + CO2(g) ^}

Driving force: Entropy increase (gas more disordered than aqueous)

Common gases formed:

• CO₂ from carbonates + acid
• H₂S from sulfides + acid
• NH₃ from ammonium salts + base
• SO₂ from sulfites + acid

4. Transfer of Electrons (Redox)

Oxidation-reduction reactions

$Zn\text{(s)} + Cu^{2+(aq) -> Zn^{2+}(aq) + Cu(s)}$

Driving force: Electron transfer to more stable configuration

5. Energy Release

Exothermic reactions (ΔH < 0)

$CH4\text{(g)} + 2O2\text{(g)} \rightarrow CO2\text{(g)} + 2H2O\text{(g)} + \text{energy}$

Driving force: System goes to lower energy state

Note: Most spontaneous reactions have at least one driving force!

Summary Table

Tips for Success

1. Learn to recognize patterns

• Count reactants and products
• Look for free elements
• Identify ions switching

2. Use activity series

• Predicts single replacement
• Higher replaces lower

3. Memorize solubility rules

• Essential for double replacement
• Predicts precipitates

4. Check for driving forces

• Precipitate, water, gas, redox
• No driving force → likely no reaction

5. Balance equations

• Atoms must be conserved
• Check your work!

Step 2: Identify pattern

• Multiple reactants → Single product
• This is COMBINATION pattern

Step 3: Classify

$\boxed{\text{Synthesis (Combination) Reaction}}$

Explanation:

• Two elements (Mg and O₂) combine to form one compound (MgO)
• Fits pattern: A + B → AB
• Also could be classified as combustion (metal + oxygen)

Additional notes:

• This is formation of ionic compound (metal + nonmetal)
• Exothermic (releases energy as heat and light)
• This is how magnesium burns with bright white flame

(b) CaCO₃(s) → CaO(s) + CO₂(g)

Step 1: Count reactants and products

• Reactants: 1 substance (CaCO₃)
• Products: 2 substances (CaO and CO₂)

Step 2: Identify pattern

• Single reactant → Multiple products
• This is BREAKDOWN pattern

Step 3: Classify

$\boxed{\text{Decomposition Reaction}}$

Explanation:

• One compound (CaCO₃) breaks down into two simpler substances
• Fits pattern: AB → A + B
• Requires heat (Δ) - thermal decomposition

Additional notes:

• This is how limestone (CaCO₃) is converted to lime (CaO)
• Important industrial process
• Endothermic (requires energy input)
• CO₂ gas escapes (driving force)

Real-world application:

• Making cement
• Occurs in caves forming stalactites/stalagmites (reverse reaction)

(c) Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

Step 1: Count reactants and products

• Reactants: 2 substances
• Products: 2 substances

Pattern so far: Could be single or double replacement

Step 2: Look more closely at reactants

• Zn(s): Free element (metal)
• HCl(aq): Compound

One reactant is free element!

Step 3: Identify what's happening

Rewrite to see the replacement:

$Zn + H-Cl \rightarrow Zn-Cl2 + H2$

• Zn replaces H in HCl
• H is displaced as H₂ gas

Step 4: Classify

$\boxed{\text{Single Replacement (Displacement) Reaction}}$

Explanation:

• Free element (Zn) replaces another element (H) in a compound
• Fits pattern: A + BC → AC + B
• Zn is more reactive than H (check activity series)

Verification using activity series:

• Activity series: ... Zn > Fe > Ni > Sn > Pb > H > Cu > Ag > Au
• Zn is above H → can displace H
• Reaction occurs ✓

Additional notes:

• This produces H₂ gas (bubbling)
• Common lab reaction
• Used to test for reactive metals

Observable evidence:

• Zn metal dissolves
• H₂ gas bubbles form
• Solution gets warm (exothermic)

(d) AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Step 1: Count reactants and products

• Reactants: 2 compounds
• Products: 2 compounds

Step 2: Identify what's happening

Look at ions switching:

Reactants:

• AgNO₃: Ag⁺ and NO₃⁻
• NaCl: Na⁺ and Cl⁻

Products:

• AgCl: Ag⁺ and Cl⁻ (Ag switched from NO₃⁻ to Cl⁻)
• NaNO₃: Na⁺ and NO₃⁻ (Na switched from Cl⁻ to NO₃⁻)

Both cations switched anions!

Pattern:

• AB + CD → AD + CB
• Ag⁺NO₃⁻ + Na⁺Cl⁻ → Ag⁺Cl⁻ + Na⁺NO₃⁻

Step 3: Classify

$\boxed{\text{Double Replacement (Metathesis) Reaction}}$

Explanation:

• Two compounds exchange ions
• Cations and anions "switch partners"
• Forms precipitate AgCl(s) - driving force!

Step 4: Verify precipitation using solubility rules

Check AgCl:

• Chloride (Cl⁻) compound
• Most chlorides soluble...
• BUT AgCl is exception (Ag⁺ is insoluble with Cl⁻)
• AgCl is insoluble → precipitates ✓

Check NaNO₃:

• Nitrate (NO₃⁻) compound
• All nitrates soluble
• NaNO₃ stays dissolved ✓

Additional notes:

• AgCl is white precipitate
• This is classic precipitation reaction
• Used in qualitative analysis to test for Cl⁻ ions

Observable evidence:

• White cloudy precipitate forms immediately
• Solution becomes opaque

Summary Table:

Key skills practiced:

1. Counting reactants/products
2. Recognizing free elements
3. Identifying ion exchanges
4. Using activity series
5. Applying solubility rules

Strategy for identifying reaction types:

1. Count substances (reactants and products)
2. Look for free elements
3. Check if ions are switching
4. Identify driving forces (precipitate, gas, water)