Homework WS3 Covalent Bonding Answerkey - Name: Date: - Studocu - Free Printable

You've uploaded an answer key for a worksheet on Covalent Bonding, not a problem to solve — it already contains the correct answers and explanations. However, I’ll walk you through each question and explain the solutions clearly so you can understand the concepts behind them.

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## ✔ Question 1: What is the difference between a molecular formula, structural formula, and an electron dot formula? Give an example of each.

🧪 Definitions & Examples:

- Molecular Formula:
→ Tells you *how many atoms* of each element are in one molecule.
→ Does not show how atoms are connected or bonded.
→ Example: `C₂H₂` (acetylene) — tells us there are 2 carbon atoms and 2 hydrogen atoms per molecule.

- Structural Formula:
→ Shows how atoms are connected using lines to represent bonds (each line = a shared pair of electrons).
→ Gives more detail than molecular formula.
→ Example: `H–H` for H₂ — shows two hydrogen atoms bonded by a single bond.

- Electron Dot Formula (Lewis Dot Structure):
→ Uses dots to represent valence electrons.
→ Shared pairs are shown as dots between atoms (or sometimes lines, but strictly speaking, dots are used in "dot" formulas).
→ Example: `H:H` for H₂ — each H has 1 electron; they share them to form a bond.

> 💡 Note: Sometimes “electron dot” and “Lewis structure” are used interchangeably. In this worksheet, they’re distinguishing electron dot (dots between atoms) from structural (lines).

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## ✔ Question 2: What is the difference between single, double, and triple covalent bonds? How many shared electrons are in each?

🔗 Bond Types:

- Single Covalent Bond:
→ 1 pair of electrons shared = 2 electrons total.
→ Represented by one line (e.g., H–H).

- Double Covalent Bond:
→ 2 pairs of electrons shared = 4 electrons total.
→ Represented by two lines (e.g., O=O in O₂).

- Triple Covalent Bond:
→ 3 pairs of electrons shared = 6 electrons total.
→ Represented by three lines (e.g., N≡N in N₂).

> ✔ Summary:
> - Single = 2 shared e⁻
> - Double = 4 shared e⁻
> - Triple = 6 shared e⁻

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## ✔ Question 3: Draw the electron dot and structural formulas for the following molecules.

We’ll go one by one and explain the reasoning.

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a) Cl₂ (Chlorine gas)

- Each Cl atom has 7 valence electrons.
- They share 1 pair (2 electrons) to complete their octet.
- So, they form a single bond.

#### ➤ Electron Dot Formula:
```
:Cl : Cl:
••
```
(Each Cl has 3 lone pairs + 1 shared pair — represented as dots between atoms)

#### ➤ Structural Formula:
```
:Cl – Cl:
```
(Each Cl still has 3 lone pairs shown as dots, bond shown as a line)

> ✔️ Worksheet answer matches: `:Cl–Cl:` with 3 lone pairs on each Cl.

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b) H₂S (Hydrogen sulfide)

- Sulfur (S) has 6 valence electrons.
- Each H has 1 valence electron.
- S forms two single bonds with two H atoms.
- After bonding, S has 2 lone pairs left (since 6 – 2 = 4 electrons → 2 pairs).

#### ➤ Electron Dot Formula:
```
H
•
: S :
•
H
```
(But usually drawn flat: `H:S:H` with two lone pairs on S)

#### ➤ Structural Formula:
```
H
|
: S :
|
H
```
(Sulfur with two lone pairs, bonded to two H’s)

> ✔️ Worksheet answer matches: `:S–H` with another H below, and two lone pairs on S.

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c) BrI (Bromine iodide)

- Both Br and I are halogens → each has 7 valence electrons.
- They share 1 pair to form a single bond.
- Each will have 3 lone pairs after bonding.

#### ➤ Electron Dot Formula:
```
:Br : I:
••
```
(Shared pair between Br and I, each with 3 lone pairs)

#### ➤ Structural Formula:
```
:Br – I:
```
(Br and I each have 3 lone pairs shown as dots)

> ✔️ Worksheet answer matches: `:I–Br:` with 3 lone pairs on each atom.

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## 📌 Final Notes:

- The worksheet is designed to help students distinguish between different ways of representing molecules.
- Electron dot emphasizes valence electrons (dots).
- Structural emphasizes connectivity (lines).
- Molecular formula only gives atom counts.

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✔ You now understand not just the *answers*, but the *why* behind them. This is key for mastering chemical bonding!

Let me know if you want to practice drawing Lewis structures for other molecules!