Bonding structure Summary
Bonding structure
Metals: metallic bonding
Ionic compound (metal +nonmetal): ionic bonding
Molecule (nonmetal +nonmetal): covalent bonding
3.1 Metallic Bonding
1) Definition
The electrostatic attraction between a lattice if positive ions and delocalized electrons.
2) The strength of metallic bonding (depend on)
Delocalized electrons (=valence electrons=Group number)
More valence electrons, stronger metallic bonding
Ionic radii (=distance between nucleus and e-)
Greater ionic radius, farther distance between nucleus and Ve-, stronger the metallic bonding
3) Physical properties
Malleable and ductile
The layer of cations can slide over each other without breaking the bond.
Q: Why alloys are less malleable and ductile?
A: Different sizes of atoms are introduced and less easy for cations to slide.
Good conductor of electricity and heat
Delocalized electrons are free to move from one side of lattice to the other.
Luster
Interaction between delocalized electrons and light
Melting Point
a. Same Group:
Same valence electrons
Ionic radii increase, attraction between nucleus and Ve- decrease, strength of metallic bonding decrease.
Melting point decrease down the group.
b. Same Period
Same energy level
Ve- increase, ionic radii decrease, attraction increase, strength of metallic bonding increase.
Melting point increase across the period.
c. Alloy have lower melting point
Different atoms are introduced; weaken the strength of metallic bond.
Melting point lower
3.2 Covalent bonding
1) Fundamental knowledge
It is between atoms that have little difference in electronegativities. Usually between nonmetals.
Electronegativity: the ability for an atom to attract bonding pair of electrons in a covalent bond
Nonmetal: high electronegativity
Metal: low electronegativity
Electronegativity value:
Increase across a period
Decrease down a group
Greatest: Fluorine (4.0) Smallest: Cs (0.7)
Electronegativity difference:
＜1.8 covalent
Metals: metallic bonding
Ionic compound (metal +nonmetal): ionic bonding
Molecule (nonmetal +nonmetal): covalent bonding
3.1 Metallic Bonding
1) Definition
The electrostatic attraction between a lattice if positive ions and delocalized electrons.
2) The strength of metallic bonding (depend on)
Delocalized electrons (=valence electrons=Group number)
More valence electrons, stronger metallic bonding
Ionic radii (=distance between nucleus and e-)
Greater ionic radius, farther distance between nucleus and Ve-, stronger the metallic bonding
3) Physical properties
Malleable and ductile
The layer of cations can slide over each other without breaking the bond.
Q: Why alloys are less malleable and ductile?
A: Different sizes of atoms are introduced and less easy for cations to slide.
Good conductor of electricity and heat
Delocalized electrons are free to move from one side of lattice to the other.
Luster
Interaction between delocalized electrons and light
Melting Point
a. Same Group:
Same valence electrons
Ionic radii increase, attraction between nucleus and Ve- decrease, strength of metallic bonding decrease.
Melting point decrease down the group.
b. Same Period
Same energy level
Ve- increase, ionic radii decrease, attraction increase, strength of metallic bonding increase.
Melting point increase across the period.
c. Alloy have lower melting point
Different atoms are introduced; weaken the strength of metallic bond.
Melting point lower
3.2 Covalent bonding
1) Fundamental knowledge
It is between atoms that have little difference in electronegativities. Usually between nonmetals.
Electronegativity: the ability for an atom to attract bonding pair of electrons in a covalent bond
Nonmetal: high electronegativity
Metal: low electronegativity
Electronegativity value:
Increase across a period
Decrease down a group
Greatest: Fluorine (4.0) Smallest: Cs (0.7)
Electronegativity difference:
＜1.8 covalent