Chemical Bonding review games for Chemistry.
This unit covers ionic bonds, covalent bonds, Lewis structures and electronegativity — essential concepts for Chemistry. Use our interactive study games to test your understanding, or review questions in traditional format below.
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This unit covers ionic bonds, covalent bonds, Lewis structures and electronegativity — essential concepts for Chemistry. Use our interactive study games to test your understanding, or review questions in traditional format below.
Key Concepts Breakdown
1 Ionic Bonds
Ionic bonds form between a metal and a nonmetal through the transfer of electrons, creating oppositely charged ions that attract each other. Students must know that metals lose electrons to form cations and nonmetals gain electrons to form anions. The resulting compound is electrically neutral overall.
Key Points
- Metal loses electrons → cation (positive charge); nonmetal gains electrons → anion (negative charge)
- Ionic compounds form crystal lattice structures and have high melting points
- Ionic compounds conduct electricity when dissolved in water or melted
- The charge of each ion is determined by how many electrons are lost or gained to reach a full outer shell
Predict the formula for the ionic compound formed between calcium (Ca) and chlorine (Cl).
Calcium is in Group 2, so it loses 2 electrons to form Ca²⁺. Chlorine is in Group 17, so it gains 1 electron to form Cl⁻. To balance the charges, you need 1 Ca²⁺ and 2 Cl⁻, giving the formula CaCl₂.
2 Covalent Bonds
Covalent bonds form between two nonmetals through the sharing of electron pairs. Students must distinguish between single (2 electrons shared), double (4 electrons shared), and triple bonds (6 electrons shared), and understand that more shared pairs create shorter, stronger bonds.
Key Points
- Nonmetal + nonmetal → covalent bond (electron sharing, not transfer)
- Single bond < double bond < triple bond in terms of bond strength and shortness
- Molecules formed by covalent bonds generally have lower melting points than ionic compounds
- Polar covalent bonds occur when electrons are shared unequally due to differing electronegativities
Classify the bond in O₂ and state how many electrons are shared.
Oxygen and oxygen are both nonmetals, so they form a covalent bond. Each oxygen needs 2 more electrons to complete its octet, so they share 2 pairs of electrons, forming a double bond with 4 electrons total shared between them.
3 Lewis Structures
Lewis structures are diagrams that show how valence electrons are arranged around atoms in a molecule, including bonding pairs and lone pairs. Students must be able to draw them correctly by counting total valence electrons, forming bonds, and satisfying the octet rule for each atom (with hydrogen limited to 2 electrons).
Key Points
- Count total valence electrons: use the periodic group number for each atom, adjust for charge if ion
- Hydrogen gets 2 electrons (duet rule); most other atoms follow the octet rule (8 electrons)
- Place lone pairs on outer atoms first, then on the central atom
- If the central atom lacks an octet, convert lone pairs to double or triple bonds
Draw the Lewis structure for CO₂.
Carbon has 4 valence electrons and each oxygen has 6, giving 4 + 6 + 6 = 16 total valence electrons. Carbon is the central atom; placing single bonds to each oxygen uses 4 electrons, leaving 12 for lone pairs. Filling oxygen octets uses all 12, but carbon only has 4 electrons — so each lone pair on oxygen is converted to a shared pair, forming two double bonds (O=C=O), satisfying all octets with 16 electrons used.
4 Electronegativity
Electronegativity is a measure of how strongly an atom attracts shared electrons in a bond. Students must know the general trend (increases across a period, decreases down a group) and use electronegativity differences to classify bonds as nonpolar covalent, polar covalent, or ionic.
Key Points
- Electronegativity increases across a period (left → right) and decreases down a group (top → bottom); fluorine is highest
- ΔEN = 0–0.4: nonpolar covalent bond; ΔEN = 0.5–1.7: polar covalent bond; ΔEN > 1.7: ionic bond
- In a polar covalent bond, the more electronegative atom carries a partial negative charge (δ⁻)
- Bond polarity affects molecular polarity, which affects physical properties like boiling point and solubility
Classify the bond between hydrogen (EN = 2.1) and fluorine (EN = 4.0). Which atom is δ⁻?
The electronegativity difference is 4.0 − 2.1 = 1.9, which is greater than 1.7, so this bond is ionic in character — though HF is typically classified as polar covalent in introductory courses because it is a molecule. Fluorine has the higher electronegativity, so it pulls the shared electrons closer to itself and carries the partial negative charge (δ⁻), while hydrogen is δ⁺.
Questions, answered.
What is Chemical Bonding?
Chemical Bonding is Unit 3 of Chemistry, covering ionic bonds, covalent bonds, Lewis structures and electronegativity.
How to study for Chemistry Unit 3?
Start with the Quick Summary above, review the Key Concepts, then test yourself with our interactive study games. Aim for 80%+ accuracy before moving on.
How many questions are in this unit?
This unit has 28+ review questions across 5 different game modes.