An Introduction to Chemical Reactions

A chemical reaction is a process that leads to the chemical transformation of one set of chemical substances into another. At the heart of chemistry is understanding how and why these reactions occur. We can classify most chemical reactions into one of five main types.

1. Synthesis (or Combination) Reaction

In a synthesis reaction, two or more simple substances combine to form a more complex product. Think of it as building something bigger from smaller pieces.

General Formula: A + B ⟶ AB

Example: The reaction between sodium and chlorine gas to produce sodium chloride (table salt).

2Na + Cl₂ ⟶ 2NaCl

2. Decomposition Reaction

A decomposition reaction is the opposite of a synthesis reaction. A complex substance breaks down into two or more simpler substances.

General Formula: AB ⟶ A + B

Example: The decomposition of hydrogen peroxide into water and oxygen gas.

2H₂O₂ ⟶ 2H₂O + O₂

3. Single-Displacement (or Single-Replacement) Reaction

In this type of reaction, a single uncombined element replaces another element in a compound. Essentially, one element "kicks out" another.

General Formula: A + BC ⟶ AC + B

Example: Zinc metal reacts with copper(II) chloride. The zinc replaces the copper.

Zn + CuCl₂ ⟶ ZnCl₂ + Cu

4. Double-Displacement (or Double-Replacement) Reaction

In a double-displacement reaction, the positive and negative ions of two ionic compounds exchange places to form two new compounds. It's like the compounds "swap partners."

General Formula: AB + CD ⟶ AD + CB

Example: When silver nitrate reacts with sodium chloride, they form silver chloride and sodium nitrate.

AgNO₃ + NaCl ⟶ AgCl + NaNO₃

5. Combustion Reaction

A combustion reaction is a rapid reaction between a substance with an oxidant, usually oxygen, to produce heat and light. When the substance is a hydrocarbon (made of carbon and hydrogen), the products are always carbon dioxide and water.

General Formula (for hydrocarbons): Hydrocarbon + O₂ ⟶ CO₂ + H₂O

Example: The combustion of methane (natural gas).

CH₄ + 2O₂ ⟶ CO₂ + 2H₂O

Practice Problems

Instructions: For each of the following reactions, identify which of the five types of chemical reactions it is.

1. H₂ + O₂ ⟶ H₂O

• Type of Reaction: _________________________

2. Mg + 2HCl ⟶ MgCl₂ + H₂

• Type of Reaction: _________________________

3. CaCO₃ ⟶ CaO + CO₂

• Type of Reaction: _________________________

4. C₃H₈ + 5O₂ ⟶ 3CO₂ + 4H₂O

• Type of Reaction: _________________________

5. 2KOH + H₂SO₄ ⟶ K₂SO₄ + 2H₂O

• Type of Reaction: _________________________

Chemical Reactions - ANSWER KEY

1. 2H₂ + O₂ ⟶ 2H₂O

Type of Reaction: Synthesis

Explanation: Two simple substances (hydrogen and oxygen) combine to form a more complex one (water).

2. Mg + 2HCl ⟶ MgCl₂ + H₂

Type of Reaction: Single-Displacement

Explanation: Magnesium (a single element) replaces the hydrogen in hydrochloric acid (HCl).

3. CaCO₃ ⟶ CaO + CO₂

Type of Reaction: Decomposition

Explanation: A complex substance (calcium carbonate) breaks down into two simpler substances (calcium oxide and carbon dioxide).

4. C₃H₈ + 5O₂ ⟶ 3CO₂ + 4H₂O

Type of Reaction: Combustion

Explanation: A hydrocarbon (propane, C₃H₈) reacts with oxygen to produce carbon dioxide and water.

5. 2KOH + H₂SO₄ ⟶ K₂SO₄ + 2H₂O

Type of Reaction: Double-Displacement

Explanation: The potassium (K⁺) from KOH swaps with the hydrogen (H⁺) from H₂SO₄. The positive ions have "swapped partners."

Chemistry: Polar and Nonpolar Covalent Bonds // Polar and Nonpolar Covalent Molecules

What Are Covalent Bonds?

First, let's quickly review what a covalent bond is. Atoms are most stable when their outer orbital level of electrons is full. To achieve this, nonmetal atoms will share electrons with each other. This sharing of electrons creates a strong bond that holds the atoms together, forming a molecule.

Think of it like two people who each have one dog, but they both want to be walking two dogs. So, they agree to share their dogs and walk them together. Each person now gets to interact with two dogs. Similarly, atoms share electrons to get a full outer shell.

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Part 1: Polar vs. Nonpolar Covalent Bonds

The key to understanding polar and nonpolar bonds is a concept called electronegativity.

Electronegativity: This is just a fancy word for how much an atom "pulls" on shared electrons in a bond. Think of it as an atom's "greed" for electrons. Some atoms, like Oxygen and Fluorine, are very greedy (high electronegativity), while others, like Hydrogen and Carbon, are less greedy (lower electronegativity). 

This difference in "greed" creates two types of covalent bonds:

1. Nonpolar Covalent Bonds (Equal Sharing)

This happens when two atoms share electrons equally. It's like a perfectly balanced tug-of-war.

When does this happen? It happens when the two atoms are the same (like an O₂ oxygen molecule) or when their electronegativity values are very, very close (like a C-H bond in methane).

Analogy: Two identical twins pulling on a toy with the exact same strength. The toy doesn't move closer to one or the other.

Result: The electron pair is shared right in the middle. There are no positive or negative "ends" to the bond.

2. Polar Covalent Bonds (Unequal Sharing)

This is the more common type. It happens when one atom is "greedier" (more electronegative) than the other, causing the electrons to be shared unequally.

When does this happen? It happens when two different atoms are bonded (like an O-H bond in water). The oxygen atom is much "greedier" than the hydrogen atom.

Analogy: A big, strong adult and a small child playing tug-of-war. The adult (the more electronegative atom) pulls the rope (the electrons) much closer to themselves.

Result: The electrons spend more time around the "greedier" atom. This creates:  A slight negative charge on the "greedier" atom (e.g., Oxygen). And, a slight positive charge on the "weaker" atom (e.g., Hydrogen).

This separation of charge is called a dipole (meaning "two poles," like the poles of a magnet).

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Part 2: Polar vs. Nonpolar Covalent Compounds (Molecules)

Now we're looking at the entire molecule, not just one bond. A molecule's overall polarity depends on two things:

1. Does it have polar bonds?

2. What is its 3D shape (molecular geometry)?

The Big Rule: A molecule is nonpolar if all the polar bonds (dipoles) cancel each other out due to its symmetrical shape. A molecule is polar if its polar bonds do not cancel out, leaving one side of the molecule more negative and one side more positive. Trigonal, pyramidal, and bent shapes will always form a polar molecule. For molecules that take either a linear or a tetrahedral shape, you need to look at what is happening with the bonds to decide if it is a nonpolar or polar molecule. 

Let's look at the two examples:

1. Nonpolar Compound Example: Carbon Dioxide (CO₂)

Bonds: The O=C bond is polar. Oxygen is "greedier" than carbon, so it pulls the shared electrons toward itself. This creates two polar bonds.

Shape: CO₂ is a linear molecule. The atoms are in a straight line (O=C=O).

Result: Imagine the carbon atom in the middle. It's being pulled by one oxygen on the left and one oxygen on the right with the exact same strength. These two "tugs" are in perfectly opposite directions, so they cancel each other out. 

Conclusion: Even though it has two polar bonds, the overall CO₂ molecule is nonpolar because its symmetrical shape causes the polar bonds to cancel. 

2. Polar Compound Example: Water (H₂O) 

Bonds: The O-H bond is very polar. Oxygen is way "greedier" than hydrogen and pulls the electrons in. 

Shape: Water is not linear. Because of extra electrons on the oxygen atom, the molecule is forced into a bent shape (like a "V" or Mickey Mouse head). 

Result: The two polar O-H bonds are now pulling at an angle. They don't pull in opposite directions, so they do not cancel each other out. Instead, all the "pull" (the negative charge) gathers on the oxygen side, and the "weaker" hydrogen side is left with a positive charge. 

Conclusion: The water molecule has a negative "end" (the oxygen) and a positive "end" (the hydrogens). This makes it a polar molecule. This polarity is why water is so good at dissolving things!

Writing Formulas and Names for Ionic Compounds

How to Write Ionic Formulas

Ionic compounds are formed when a metal (which forms a positive ion, or cation) transfers electrons to a nonmetal (which forms a negative ion, or anion). The compound formed is neutral, which means the total positive charge must equal the total negative charge.

We can use the "Criss-Cross Method" to easily find the chemical formula.

Example: Aluminum Oxide

1. Write the symbols and charges for the ions.

• Aluminum is Al³⁺

• Oxygen (as an oxide ion) is O²⁻

2. "Criss-cross" the numbers. The number from the charge of one ion becomes the subscript for the other ion. Ignore the +/- signs.

• The 3 from Al³⁺ moves to O.

• The 2 from O²⁻ moves to Al.

3. Write the final formula.

• Al₂O₃

Important Note: If the subscripts can be simplified, you must reduce them to the lowest whole-number ratio. For example, for Calcium Oxide (Ca²⁺ and O²⁻), the criss-cross method gives Ca₂O₂, which simplifies to CaO.

Practice Problems

Write the correct chemical formula for each of the following ionic compounds.

Part 1: Simple Binary Compounds

1. Sodium Chloride: ___________________

2. Magnesium Sulfide: ___________________

3. Aluminum Bromide: ___________________

4. Potassium Oxide: ___________________

5. Calcium Nitride: ___________________

Part 2: Compounds with Polyatomic Ions

(Remember to use parentheses if you need more than one polyatomic ion!)

6. Sodium Nitrate: ___________________

7. Calcium Hydroxide: ___________________

8. Aluminum Sulfate: ___________________

9. Magnesium Phosphate: ___________________

10. Ammonium Chloride: ___________________

How to Name Ionic Compounds

Naming ionic compounds follows a set of rules. The key is to identify the type of ions involved: simple metal and nonmetal, a metal with multiple possible charges, or a polyatomic ion.

Rule 1: Simple Binary Compounds (Metal + Nonmetal)

1. Write the name of the metal cation first.

2. Write the base name of the nonmetal anion and change the ending to -ide.

• Example: NaCl

• Na = Sodium

• Cl = Chlorine → Chloride

• Name: Sodium Chloride

Rule 2: Compounds with Transition Metals

Transition metals (and some others, like lead and tin) can form more than one type of positive ion. We use a Roman numeral in parentheses to show the charge of the metal cation.

1. Name the metal cation.

2. Determine the charge on the metal by looking at the anion's charge. The total charge of the compound must be zero.

3. Write the charge as a Roman numeral in parentheses.

4. Name the anion with the -ide ending.

• Example: FeCl₃

• We know each Cl is Cl⁻. There are three of them, for a total negative charge of 3-.

• To balance this, the one Fe ion must be Fe³⁺.

• Name: Iron (III) Chloride

Rule 3: Compounds with Polyatomic Ions

1. Name the cation (metal or polyatomic cation like NH₄⁺).

2. Name the polyatomic anion exactly as it is written in the chart.

• Example: Ca(NO₃)₂

• Ca = Calcium

• NO₃⁻ = Nitrate

• Name: Calcium Nitrate

Practice Problems

Write the correct name for each of the following ionic compounds.

Part 1: Simple Binary Compounds

1. K₂O: ___________________

2. AlBr₃: ___________________

3. MgS: ___________________

4. Ca₃P₂: ___________________

Part 2: Compounds with Transition Metals

5. CuO: ___________________

6. FeCl₂: ___________________

7. PbS₂: ___________________

Part 3: Compounds with Polyatomic Ions

8. KNO₃: ___________________

9. Al(OH)₃: ___________________

10. (NH₄)₂SO₄: ___________________

Writing Formulas for Ionic Compounds - ANSWER KEY

Part 1: Simple Binary Compounds

1. Sodium Chloride:
   Ions: Na⁺ and Cl⁻
   Criss-cross: Na₁Cl₁
   Formula: NaCl

2. Magnesium Sulfide:
   Ions: Mg²⁺ and S²⁻
   Criss-cross: Mg₂S₂ (Simplifies to 1:1 ratio)
   Formula: MgS

3. Aluminum Bromide:
   Ions: Al³⁺ and Br⁻
   Criss-cross: Al₁Br₃
   Formula: AlBr₃

4. Potassium Oxide:
   Ions: K⁺ and O²⁻
   Criss-cross: K₂O₁
   Formula: K₂O

5. Calcium Nitride:
   Ions: Ca²⁺ and N³⁻
   Criss-cross: Ca₃N₂
   Formula: Ca₃N₂

Part 2: Compounds with Polyatomic Ions

6. Sodium Nitrate:
   Ions: Na⁺ and NO₃⁻
   Criss-cross: Na₁(NO₃)₁
   Formula: NaNO₃

7. Calcium Hydroxide:
   Ions: Ca²⁺ and OH⁻
   Criss-cross: Ca₁(OH)₂ (Parentheses are needed)
   Formula: Ca(OH)₂

8. Aluminum Sulfate:
   Ions: Al³⁺ and SO₄²⁻
   Criss-cross: Al₂(SO₄)₃ (Parentheses are needed)
   Formula: Al₂(SO₄)₃

9. Magnesium Phosphate:
   Ions: Mg²⁺ and PO₄³⁻
   Criss-cross: Mg₃(PO₄)₂ (Parentheses are needed)
   Formula: Mg₃(PO₄)₂

10. Ammonium Chloride:
    Ions: NH₄⁺ and Cl⁻
    Criss-cross: (NH₄)₁Cl₁
    Formula: NH₄Cl

Naming Ionic Compounds - ANSWER KEY

Part 1: Simple Binary Compounds

1. K₂O:
   K = Potassium, O = Oxide
   Name: Potassium Oxide

2. AlBr₃:
   Al = Aluminum, Br = Bromide
   Name: Aluminum Bromide

3. MgS:
   Mg = Magnesium, S = Sulfide
   Name: Magnesium Sulfide

4. Ca₃P₂:
   Ca = Calcium, P = Phosphide
   Name: Calcium Phosphide

Part 2: Compounds with Transition Metals

5. CuO:
   Oxygen is O²⁻, so the one copper ion must be Cu²⁺.
   Name: Copper (II) Oxide

6. FeCl₂:
   Each chlorine is Cl⁻, so two make a 2- charge. The one iron ion must be Fe²⁺.
   Name: Iron (II) Chloride

7. PbS₂:
   Each sulfide is S²⁻, so two make a 4- charge. The one lead ion must be Pb⁴⁺.
   Name: Lead (IV) Sulfide

Part 3: Compounds with Polyatomic Ions

8. KNO₃:
   K = Potassium, NO₃ = Nitrate
   Name: Potassium Nitrate

9. Al(OH)₃:
   Al = Aluminum, OH = Hydroxide
   Name: Aluminum Hydroxide

10. (NH₄)₂SO₄:
    NH₄ = Ammonium, SO₄ = Sulfate
    Name: Ammonium Sulfate