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The energy required to break one mole of covalent bonds in gaseous state is called bond energy
For Example
H=H Bond Energy = +436 KJ/mol
I-I Bond Energy = +151 KJ/mol
Factors affecting Bond Energy
• between atoms on Left Hand Side and atoms on Right Hand Side of Periodic Table
• electrons are TRANSFERRED between atoms
• atoms end up as ions
• strong electrostatic attraction between ions of opposite charge
• giant ionic crystal lattice structure
• Properties ... high melting points.....brittle.....water soluble....conduct when molten or in aqueous solution
Ionic Bonding; a video demonstration:
Drawing dot and cross diagrams
Covalent Bonding
• between atoms of the same element; (e.g. in N2, O2, diamond, graphite)
• between atoms of different elements on the RHS of table; (e.g. CO2, SO2)
• when one of the elements is in the middle of the table; (e.g. C, Si)
• head-of-the-group elements with high I.E.’s , (e.g. Be in BeCl2)
• consists of a shared pair of electrons, one electron coming from each atom
• atoms share to try and get an ‘octet’ of electrons
• leads to the formation of simple molecules and giant molecules (e.g. silica)
Polar Covalent Bonds
In bonds between atoms of the same element the sharing of the electrons is equal between the two atoms. When two atoms of different elements make a bond, the electrons will not usually be shared equally. The electrons are pulled more toward the more electronegative element.
What is electronegativity?
Electronegativity is the measure of the ability of an atom in a molecule to draw bonding electrons to itself. In general, electronegativity increases from bottom to top and left to right on the periodic table. Fluorine is the most electronegative element since it has a tendency to pick up electrons easily and hold on to them strongly. An element like cesium has a low electronegativity. The unequal sharing of electrons is called a polar covalent bond. The definition of a polar covalent bond is a covalent bond in which the bonding electrons spend more time near one atom than the other.
Dative Covalent Bonding
• consists of a shared pair of electrons, both electrons from one atom
• one species is a lone pair donor - LEWIS BASE
• other species has space in outer shell to accept a lone pair - LEWIS ACID
• once the bond has been formed it is the same as a covalent bond
Metallic Bonding
• metal atoms arranged in regular lattice give up outer shell electrons
• electrons form a mobile ‘cloud’ which binds metal ions together
• strength of bond depends on number of electrons and size of ions
• mobile electrons ... allow electricity to be conducted
Strength of a Metallic Bond:
As we go down the group the atomic size increases, the attraction between the nucleus and free electrons decreases and therefore the melting melting point also decreases down the group.
As we move across the period, atomic size decreases and the strength of the metallic bond increases therefore melting point also increases
Sources:
Images and definitions: www.knockhardy.org.uk
Videos: Youtube.com
The VSEPR theory
VSEPR theory proposes that the geometric arrangement of terminal atoms, or groups of atoms about a central atom in a covalent compound, or charged ion, is determined solely by the repulsions between electron pairs present in the valence shell of the central atom. To view more about VSEPR theory visit this page
The Sigma and Pie bonds
- The Bond Energy
- Factors affecting the Bond energy
- The Four Main Types of Bonding
- Ionic Bonding
- Covalent Bonding
- Dative Covalent Bonding
- Metallic Bonding
- Section 2
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The energy required to break one mole of covalent bonds in gaseous state is called bond energy
For Example
H=H Bond Energy = +436 KJ/mol
I-I Bond Energy = +151 KJ/mol
Factors affecting Bond Energy
- Number of Bonds: Triple bonds are stronger than double bonds and double bonds are stronger than single bonds.
- Bond Length: Bond length is the distance between two nuclei of covalently bonded atoms as the bond length increases bond strength decreases, hence bond energy decreases
• between atoms on Left Hand Side and atoms on Right Hand Side of Periodic Table
• electrons are TRANSFERRED between atoms
• atoms end up as ions
• strong electrostatic attraction between ions of opposite charge
• giant ionic crystal lattice structure
• Properties ... high melting points.....brittle.....water soluble....conduct when molten or in aqueous solution
Ionic Bonding; a video demonstration:
Drawing dot and cross diagrams
http://www.youtube.com/watch?v=8BZBzFwVXl4
Covalent Bonding
• between atoms of the same element; (e.g. in N2, O2, diamond, graphite)
• between atoms of different elements on the RHS of table; (e.g. CO2, SO2)
• when one of the elements is in the middle of the table; (e.g. C, Si)
• head-of-the-group elements with high I.E.’s , (e.g. Be in BeCl2)
• consists of a shared pair of electrons, one electron coming from each atom
• atoms share to try and get an ‘octet’ of electrons
• leads to the formation of simple molecules and giant molecules (e.g. silica)
Polar Covalent Bonds
In bonds between atoms of the same element the sharing of the electrons is equal between the two atoms. When two atoms of different elements make a bond, the electrons will not usually be shared equally. The electrons are pulled more toward the more electronegative element.
What is electronegativity?
Electronegativity is the measure of the ability of an atom in a molecule to draw bonding electrons to itself. In general, electronegativity increases from bottom to top and left to right on the periodic table. Fluorine is the most electronegative element since it has a tendency to pick up electrons easily and hold on to them strongly. An element like cesium has a low electronegativity. The unequal sharing of electrons is called a polar covalent bond. The definition of a polar covalent bond is a covalent bond in which the bonding electrons spend more time near one atom than the other.
Dative Covalent Bonding
• consists of a shared pair of electrons, both electrons from one atom
• one species is a lone pair donor - LEWIS BASE
• other species has space in outer shell to accept a lone pair - LEWIS ACID
• once the bond has been formed it is the same as a covalent bond
Metallic Bonding
• metal atoms arranged in regular lattice give up outer shell electrons
• electrons form a mobile ‘cloud’ which binds metal ions together
• strength of bond depends on number of electrons and size of ions
• mobile electrons ... allow electricity to be conducted
Strength of a Metallic Bond:
As we go down the group the atomic size increases, the attraction between the nucleus and free electrons decreases and therefore the melting melting point also decreases down the group.
As we move across the period, atomic size decreases and the strength of the metallic bond increases therefore melting point also increases
Sources:
Images and definitions: www.knockhardy.org.uk
Videos: Youtube.com
The VSEPR theory
VSEPR theory proposes that the geometric arrangement of terminal atoms, or groups of atoms about a central atom in a covalent compound, or charged ion, is determined solely by the repulsions between electron pairs present in the valence shell of the central atom. To view more about VSEPR theory visit this page
The Sigma and Pie bonds
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