Atoms bond to become more stable, usually by achieving a full outer electron shell. Ionic bonds form when electrons are transferred from a metal to a non-metal, creating oppositely charged ions that attract each other. Covalent bonds form when non-metal atoms share electrons. Metallic bonds occur between metal atoms, where electrons move freely in a shared "sea," giving metals properties like conductivity and malleability.
Example
Sodium chloride (table salt) forms an ionic bond: sodium transfers an electron to chlorine, creating a positively charged sodium ion and a negatively charged chloride ion that attract each other, while water (H₂O) forms covalent bonds, with oxygen and hydrogen atoms sharing electrons.
Key terms
Ionic bond:
A bond formed by the transfer of electrons between a metal and a non-metal.
Covalent bond:
A bond formed by non-metal atoms sharing electrons.
Metallic bond:
A bond between metal atoms sharing a "sea" of free-moving electrons.
Questions
1. An ionic bond forms by:
The transfer of electrons between a metal and non-metal
Sharing electrons between two non-metals
A "sea" of free electrons between metals
No electron interaction at all
2. A covalent bond forms by:
Non-metal atoms sharing electrons
A metal transferring electrons to another metal
No electron interaction at all
Only occurring between identical atoms of a metal
3. A metallic bond involves:
A "sea" of free-moving electrons between metal atoms
The transfer of electrons to a non-metal only
No electrons involved at all
Sharing electrons between two non-metals only
4. Atoms bond mainly to become:
More stable
Less stable
Radioactive
Invisible
5. Sodium chloride (table salt) is an example of:
An ionic bond
A covalent bond only
A metallic bond only
No bond at all
6. Water (H₂O) is an example of:
A covalent bond
An ionic bond only
A metallic bond only
No bond at all
7. Metals are known for properties like conductivity and malleability due to:
Metallic bonding
Ionic bonding only
Covalent bonding only
Having no bonding at all
8. In an ionic bond, the resulting ions have:
Opposite charges that attract each other
Identical charges that repel each other
No charge at all
A charge that is always neutral
9. A "full outer electron shell" is generally associated with:
Greater atomic stability
Greater atomic instability
No effect on stability at all
Only radioactive elements
10. Why do metals conduct electricity well, according to metallic bonding theory?
Free-moving electrons in the "sea" can carry an electric current through the structure
Metals contain no electrons at all
Metallic bonds prevent any electron movement whatsoever
Conductivity has no connection to electron behaviour
11. In sodium chloride, sodium becomes a positively charged ion because it:
Loses an electron to chlorine
Gains an extra electron from chlorine
Shares its electrons equally with chlorine
Loses a proton, not an electron
12. Why might a covalent compound like water not conduct electricity well in its pure form, unlike an ionic compound dissolved in water?
Covalent molecules generally do not have free-moving charged particles the way dissolved ions do
All covalent compounds always conduct electricity perfectly
Water molecules have no electrons involved in their bonding at all
Conductivity is entirely unrelated to the presence of charged particles
13. Why is metallic bonding often described using a "sea of electrons" model rather than fixed bonds between specific atom pairs?
Electrons in a metal are delocalised and move freely across many atoms rather than being tied to one bond
All bonding types, including metallic, always involve electrons fixed to a single atom pair
Metallic bonding has no meaningful electron behaviour at all
The "sea of electrons" model has been fully disproven and is never used
14. Why do ionic compounds like salt typically have high melting points compared to many covalent compounds?
Strong electrostatic attraction between oppositely charged ions throughout the structure requires significant energy to break
Ionic compounds have no attractive forces holding their structure together
Covalent compounds always have stronger bonds than ionic compounds
Melting point has no relationship to the type of bonding present
15. Why might the malleability of metals (being able to be bent or shaped without breaking) be explained by metallic bonding?
Layers of metal atoms can slide past each other while the shared electron sea maintains the bonding throughout
Metals are always too brittle to bend under any circumstances
Malleability has no connection to how electrons are arranged in a metal
Metallic bonds break completely every time a metal is bent
16. Why might scientists predict whether a bond between two elements will be ionic or covalent based on their position on the periodic table?
Metals bonding with non-metals tend to form ionic bonds, while two non-metals tend to form covalent bonds
Bond type can never be predicted from an element's position on the periodic table
All elements form identical bond types regardless of their properties
Ionic and covalent bonds are indistinguishable from one another
17. Why is understanding chemical bonding foundational to explaining the physical properties of everyday materials?
The type of bonding directly influences properties like melting point, conductivity and hardness
Bonding type has no connection to any physical property of a material
Physical properties are always identical regardless of bonding type
Only appearance, not bonding, determines a material's physical properties
18. Why do noble gases (like neon and argon) rarely form bonds with other elements?
They already have a full outer electron shell, giving them little tendency to gain, lose or share electrons
Noble gases always form the strongest possible chemical bonds
These elements have no electrons at all in their outer shell
Bonding tendency has no connection to outer shell electron configuration
19. Carbon dioxide (CO₂), where carbon shares electrons with two oxygen atoms, is an example of:
Covalent bonding
Ionic bonding only
Metallic bonding only
A substance with no bonding at all
20. Why might diamond (a covalent network of carbon atoms) be extremely hard, unlike many simpler covalent molecules?
Its atoms are held in a rigid, extensively interconnected lattice of strong covalent bonds throughout the structure
Diamond contains no chemical bonds of any kind
Covalent bonding always produces very soft, weak materials
Diamond's hardness has no connection to its bonding structure
21. Why might alloys (mixtures of metals, like bronze) often be harder than the pure metals they are made from?
Mixing different-sized atoms disrupts the regular layers in metallic bonding, making it harder for them to slide past each other
Alloys always have identical properties to their individual pure metals
Mixing metals together always weakens metallic bonding completely
Hardness has no connection to how atoms are arranged within a metal
Answer key (parent copy)
1. The transfer of electrons between a metal and non-metal
2. Non-metal atoms sharing electrons
3. A "sea" of free-moving electrons between metal atoms
4. More stable
5. An ionic bond
6. A covalent bond
7. Metallic bonding
8. Opposite charges that attract each other
9. Greater atomic stability
10. Free-moving electrons in the "sea" can carry an electric current through the structure
11. Loses an electron to chlorine
12. Covalent molecules generally do not have free-moving charged particles the way dissolved ions do
13. Electrons in a metal are delocalised and move freely across many atoms rather than being tied to one bond
14. Strong electrostatic attraction between oppositely charged ions throughout the structure requires significant energy to break
15. Layers of metal atoms can slide past each other while the shared electron sea maintains the bonding throughout
16. Metals bonding with non-metals tend to form ionic bonds, while two non-metals tend to form covalent bonds
17. The type of bonding directly influences properties like melting point, conductivity and hardness
18. They already have a full outer electron shell, giving them little tendency to gain, lose or share electrons
19. Covalent bonding
20. Its atoms are held in a rigid, extensively interconnected lattice of strong covalent bonds throughout the structure
21. Mixing different-sized atoms disrupts the regular layers in metallic bonding, making it harder for them to slide past each other