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Quizzes > Science > Chemistry

Master Ionic, Covalent & Metallic Bonds - Take the Quiz!

Think you can ace covalent vs ionic vs metallic bonding? Start now!

Difficulty: Moderate
2-5mins
Learning OutcomesCheat Sheet
Paper art illustration for chemistry quiz on Ionic, Covalent, Metallic Bonds on dark blue background

This Ionic, Covalent & Metallic Bonds quiz helps you tell ionic, covalent, and metallic bonds apart, pick the right examples, and recall traits like electron sharing, charge, and conductivity. Use it to check for gaps before a test, and for a quick review try the bonding warm-up and the short practice .

Which type of chemical bond results from the transfer of electrons between atoms?
Hydrogen bond
Metallic bond
Covalent bond
Ionic bond
Ionic bonds form when one atom donates electrons to another, creating positively and negatively charged ions that attract each other. This transfer leads to electrostatic attraction between ions. In contrast, covalent bonds share electrons and metallic bonds involve delocalized electrons.
What type of bond involves the sharing of electron pairs between atoms?
Van der Waals bond
Covalent bond
Metallic bond
Ionic bond
Covalent bonds occur when two atoms share one or more pairs of electrons to attain a full valence shell. This shared pair holds the atoms together. Ionic bonds instead involve full electron transfer, and metallic bonds involve a sea of delocalized electrons.
Which bond is characterized by a lattice of positive ions surrounded by delocalized electrons?
Metallic bond
Dipole - dipole interaction
Covalent bond
Ionic bond
Metallic bonds consist of a lattice of metal cations immersed in a 'sea' of delocalized valence electrons, allowing metals to conduct electricity. This electron sea also explains malleability and ductility. Ionic bonds have fixed ions in a lattice but lack free electrons for conductivity.
How are ions held together in an ionic bond?
Hydrogen bonding
Electrostatic attraction
Electron sharing
Van der Waals forces
Ions in an ionic bond attract one another due to opposite charges, generating strong electrostatic forces. This attraction forms a crystal lattice in many salts. Electron sharing describes covalent bonds, while hydrogen bonds and van der Waals forces are weaker intermolecular interactions.
What type of bond is present in table salt (NaCl)?
Coordinate covalent bond
Ionic bond
Covalent bond
Metallic bond
NaCl forms through the transfer of an electron from sodium to chlorine, creating Na+ and Cl? ions. These oppositely charged ions attract each other to form an ionic lattice. Covalent bonds share electrons rather than transfer them.
What type of bond is found in the diatomic molecule H2?
Nonpolar covalent bond
Metallic bond
Ionic bond
Polar covalent bond
In H2, two hydrogen atoms share their electrons equally, forming a nonpolar covalent bond. Since both atoms have the same electronegativity, the shared pair is not displaced. Polar covalent bonds occur when atoms differ in electronegativity.
Which type of bond imparts malleability and ductility to metals?
Ionic bond
Metallic bond
Hydrogen bond
Covalent bond
The delocalized electrons in metallic bonds allow metal ions to slide past each other when stress is applied, giving metals their malleability and ductility. Ionic and covalent bonds are more rigid, leading to brittleness in many solids.
Which bond type generally has the highest melting point?
Hydrogen bond
Covalent bond
Metallic bond
Ionic bond
Ionic compounds often have very high melting points due to strong electrostatic attractions between ions in their lattice. Covalent and metallic bonds vary but are generally lower. Hydrogen bonds are much weaker, giving low melting points.
The strength of an ionic bond increases with:
Increasing ionic radius
Increasing temperature
Decreasing ionic charge
Increasing ionic charge and decreasing ionic radius
Lattice energy, a measure of ionic bond strength, increases when ions carry higher charges and are closer together (smaller radii). Larger ions spread charge over a bigger volume, reducing attraction. Temperature affects kinetic energy, not bond strength.
An electronegativity difference greater than approximately 1.7 usually indicates what bond type?
Ionic bond
Covalent bond
Metallic bond
Van der Waals interaction
Pauling's scale suggests that when the difference in electronegativity between two atoms exceeds about 1.7, electron transfer is favorable, forming an ionic bond. Smaller differences favor covalent bonding. Metallic bonds are not defined by electronegativity differences.
What is the molecular geometry of CO2?
Tetrahedral
Bent
Trigonal planar
Linear
CO2 has two regions of electron density around the central carbon atom with no lone pairs, resulting in a linear shape with a bond angle of 180°. Bent and trigonal geometries arise with different electron pair counts.
Why are metals good conductors of electricity?
Because of ion migration
Because of mobile delocalized electrons
Because of fixed covalent bonds
Because of hydrogen bonding
In metals, valence electrons become delocalized across the entire lattice, forming an 'electron sea' that can move freely under an electric field. This mobility allows metals to conduct current. Ionic solids conduct only when ions can move.
What is a coordinate covalent bond?
A bond in a metallic lattice
A bond with ionic character
A bond where one atom provides both electrons
A bond involving hydrogen
In a coordinate covalent (dative) bond, both electrons in the shared pair originate from the same atom. After formation, it is indistinguishable from a regular covalent bond. This occurs in complexes and some Lewis acids/bases.
Which property is typical of covalent network solids like diamond?
Solubility in water
Good electrical conductivity
Malleability
High melting point
Covalent network solids like diamond have each atom bonded strongly in a 3D lattice, leading to extremely high melting points. They are poor electrical conductors and are generally hard and insoluble. Malleability and solubility are not typical.
Diamond consists of what type of bonding?
Metallic bonding
Hydrogen bonding
Ionic bonding
Network covalent bonding
Diamond is a covalent network solid where each carbon atom is sp³-hybridized and bonded tetrahedrally to four others, creating a three-dimensional lattice. There are no free electrons or ions, so diamond is a poor conductor.
In ionic compounds, lattice energy is most directly related to:
Molecular weight
Number of lone pairs
Bond polarity
Ionic charges and internuclear distance
Lattice energy increases with the product of the charges on the ions and decreases with the distance between ion centers (Coulomb's law). Molecular weight and lone pairs are not direct factors.
According to band theory, what allows metals to conduct electricity?
Strong ionic interactions
Covalent network formation
Overlapping energy bands forming a conduction band
Fixed valence electrons
Band theory describes that in metals, the valence and conduction bands overlap or are partially filled, allowing electrons to move freely under an applied potential. This delocalization enables conductivity.
Which molecule has the most polar bond?
NH3
HF
H2O
HCl
HF has the largest electronegativity difference (~1.9) between hydrogen and fluorine, making its bond the most polar among these. HCl and H2O are polar but less so, and NH3 is moderately polar.
Compare the bond lengths of single, double, and triple bonds between the same two atoms. Which is shortest?
They are the same length
Triple bond
Double bond
Single bond
Higher bond orders (more shared electron pairs) pull atoms closer together, reducing bond length. Therefore, a triple bond is shortest, followed by double, then single. Bond length trends inversely with bond strength.
CaF2 forms what type of crystal lattice?
Metallic
Covalent network
Molecular
Ionic lattice
Calcium fluoride is composed of Ca²? and F? ions arranged in a fluorite-type ionic lattice. There are no covalent networks or metallic bonds. It crystallizes in a cubic ionic structure.
Brass is an example of:
A molecular solid
An ionic compound
A covalent network
A metallic alloy
Brass is an alloy of copper and zinc, bound by metallic bonds with delocalized electrons. It's not ionic or covalent network, and molecular solids involve discrete molecules held by intermolecular forces.
Why do solid ionic compounds not conduct electricity, but their molten state does?
Ions are fixed in the solid lattice but mobile when molten
Covalent bonds form upon melting
Electrons are delocalized only in the solid
Hydrogen bonding increases in the liquid
In the solid state, ions are locked into a rigid lattice and cannot move to conduct electricity. Upon melting, the lattice breaks down, freeing ions to move and carry charge. This ionic mobility in the liquid phase enables electrical conductivity.
Which relationship between bond order (n) and bond length (d) is correct?
Bond order and length are independent
Bond length depends only on atom size
As bond order increases, bond length decreases
As bond order increases, bond length increases
Greater bond order means more shared electron pairs, which pull bonded atoms closer together, shortening the bond distance. This also correlates with higher bond energy. Atomic size influences length but does not override the bond order effect.
In transition metals, how do d orbitals contribute to metallic bonding?
d orbitals prevent electron mobility
d orbitals form localized covalent bonds
Overlap of d orbitals allows additional delocalized electrons
d orbitals only interact in ionic compounds
Transition metals have partially filled d orbitals that overlap, contributing extra electrons to the conduction band. This delocalization strengthens metallic bonding and explains properties like high conductivity and variable oxidation states.
In a Born - Haber cycle for NaCl, which data must be known to calculate the lattice energy?
All of the above (sublimation energy, ionization energy, bond dissociation energy, electron affinity, enthalpy of formation)
Only enthalpy of formation
Only ionization energy and electron affinity
Only bond dissociation energies
A Born - Haber cycle for NaCl requires summing multiple steps: sublimation of Na, ionization energy, half the Cl2 bond dissociation, electron affinity of Cl, and the enthalpy of formation. This allows calculation of the lattice energy by Hess's law.
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Study Outcomes

  1. Distinguish Bond Types -

    Analyze the characteristics of ionic, covalent, and metallic bonds and distinguish among them based on electron transfer, sharing, and delocalization.

  2. Compare Physical Properties -

    Compare the conductivity, melting points, and structural traits of ionic vs covalent vs metallic compounds to predict their behavior under different conditions.

  3. Interpret Lattice Structures -

    Interpret ionic lattice and metallic crystal structures, emphasizing how these arrangements relate to chemical bond ionic and metallic properties.

  4. Assess Metalloid Bonding -

    Assess the propensity of metalloids to form covalent bonds by examining their periodic table positions and bonding tendencies.

  5. Apply Classification Skills -

    Apply knowledge of ionic covalent metallic bonds to classify real-world materials and explain how bonding types determine material properties.

  6. Evaluate Your Understanding -

    Evaluate your grasp of covalent vs ionic vs metallic bonds through quiz performance and identify areas for further study.

Cheat Sheet

  1. Electron Transfer in Ionic Bonds -

    Ionic bonds form when atoms with a large electronegativity difference (typically >1.7 on the Pauling scale) transfer electrons to achieve full valence shells. For instance, NaCl forms via Na → Na❺ + e❻ and Cl + e❻ → Cl❻, resulting in a stable ionic lattice. A handy mnemonic is "metals lose, nonmetals gain," straight from university-level chemistry curricula.

  2. Electron Sharing in Covalent Bonds -

    Covalent bonds arise when atoms share electron pairs to reach an octet configuration, as seen in H₂O (H - O - H) or O₂ (O=O). Bond polarity depends on electronegativity differences: <0.4 nonpolar, 0.4 - 1.7 polar covalent (source: Royal Society of Chemistry). Try picturing a tug-of-war rope - if both pull equally, it's nonpolar; if one wins, it's polar!

  3. Sea of Electrons in Metallic Bonds -

    Metallic bonds feature a delocalized "sea" of electrons flowing freely around positive metal ions, granting metals high electrical conductivity and malleability. Classic examples include Cu and Fe, where atoms in a lattice share electrons in all directions (source: American Chemical Society). Remember "metallic metals conduct" - a simple catchphrase highlighting conductivity.

  4. Crystal Lattice Structures -

    Ionic compounds adopt rigid, repeating crystal lattices like the face-centered cubic arrangement of NaCl, which explains their high melting points and brittleness. Covalent network solids, such as diamond (a C-C network), also form giant lattices but with directional bonds. Compare these to metallic crystals - each structure underpins unique properties, according to materials science research at top universities.

  5. Comparing Bond Types Mnemonic -

    To recall covalent vs ionic vs metallic bonds, use "Ionic Is Transferring, Covalent Is Sharing, Metallic Is Sea" as a memory trick. Ionic bonds feature electrostatic attraction, covalent bonds focus on electron sharing, and metallic bonds involve delocalized electrons. This concise phrase aligns with guidelines from educational institutions like Khan Academy and helps you ace any quiz on ionic, covalent and metallic bonds.

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