Chapter 1 Electronic Structure and Bonding
- The Structure of an Atom
Protons are positively charged. Neutrons have no charge. Electrons are negatively charged.
atomic number = # of protons
atomic number of carbon = 6
Neutral carbon has six protons and six electrons.
An atom consists of electrons, positively charged protons, and neutral neutrons
• Electrons form chemical bonds
• Atomic number: numbers of protons in its nucleus
• Mass number: the sum of the protons and neutrons of an atom
• Isotopes have the same atomic number but di erent mass numbers
• The atomic weight: the average weighted mass of its atoms
• Molecular weight: the sum of the atomic weights of all the atoms in the molecule
- The Distribution of Electrons in an Atom
Quantum mechanics uses the mathematical equation of wave motions to characterize the motion
of an electron around a nucleus
Wave functions or orbitals tell us the energy of the electron and the volume of space around the
nucleus where an electron is most likely to be found
The atomic orbital closer to the nucleus has the lowest energy
Degenerate orbitals have the same energy
The rst shell is closest to the nucleus.
The closer the atomic orbital is to the nucleus, the lower its energy.
Within a shell, energy of s < p.
- Relative Energies of the Atomic Orbitals
Aufbau principle: An electron goes into the atomic orbital with the lowest energy.
Pauli exclusion principle: No more than two electrons can be in an atomic orbital.
Hund’s rule: An electron goes into an empty degenerate orbital rather than pairing up.
-> Core and valence electrons
Lewis’s theory: an atom will give up, accept, or share electrons in order to achieve a lled outer
shell or an outer shell that contains eight electrons. + is lost electron, - is gained.
- A Hydrogen Atom Can Lose or Gain an Electron
A hydrogen atom achieves an empty shell by losing an electron or a lled outer shell by gaining an
electron.
- Bonding
Ionic Bonds Are Formed by the Transfer of Electrons
An ionic bond is the attraction between ions of opposite charge
Attractive forces between opposite charges are called electrostatic attractions
Achieving a lled outer shell by sharing electrons. A bond formed by sharing electrons is called a
covalent bond.
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,- Nonpolar and Polar Covalent Bonds
Nonpolar covalent bond = bonded atoms are the same or have similar electronegativities.
Polar covalent bond = bonded atoms have di erent electronegativities.
Equal sharing of electrons: nonpolar covalent bond (e.g., H2)
Sharing of electrons between atoms of di erent electronegativities: polar covalent bond (e.g., HF)
- The Greater the Di erence in Electronegativity, the More Polar the Bond
Nonpolar covalent bond: electonegativity di erence < 0.5
Polar covalent bond: electonegativity di erence 0.5 – 1.9
Electronegativity di erence > 1.9: electrons are not shared; atoms are held together by the
attraction of opposite charges
- Dipole moment
Dipole moment = size of the charge x the distance between the charges
The greater the di erence in electronegativity, the greater the dipole moment and the more polar
the bond.
Red has the most negative electrostatic potential; attracts positive charge. Blue has the most
positive electrostatic potential; attracts negative charge.
- Formal charge
Formal Charge = the # of valence electrons – (the # of lone-pair electrons + the # of bonds)
- Carbon forms four bonds
If carbon does not form four bonds, it has a charge (or it is a radical).
- Nitrogen forms three bonds
Nitrogen has one lone pair. If nitrogen does not form three bonds, it is charged.
- Oxygen forms two bonds
Oxygen has two lone pairs. If oxygen does not form two bonds, it is charged.
- Hydrogen and the Halogens Form One Bond
A halogen has three lone pairs. If hydrogen or halogen does not form one bond, it has a charge (or
it is a radical).
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, - Organic chemistry
Organic compounds are compounds containing carbon
Carbon neither readily gives up nor readily accepts electrons
Carbon shares electrons with other carbon atoms as well as with several di erent kinds of atoms
- Shape of atomic orbitals
An orbital tells us the volume of space around the nucleus where an electron is most likely to be
found
The s Orbitals
A node is at a region where a standing wave has an amplitude of zero
The p Orbitals
- Molecular orbitals
Molecular orbitals belong to the whole molecule
sigma bond: formed by overlapping of two s orbitals or by end-on overlap of two p orbitals
A sigma bond is stronger than alpha bond
Bond strength/bond dissociation: energy required to break a bond or energy released to form a
bond
- Atomic orbitals can combine in two ways
In-phase overlap forms a bonding MO; out-of-phase overlap forms an antibonding MO
Pi bond (p) is formed by sideways overlap of two parallel p orbitals
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- The Structure of an Atom
Protons are positively charged. Neutrons have no charge. Electrons are negatively charged.
atomic number = # of protons
atomic number of carbon = 6
Neutral carbon has six protons and six electrons.
An atom consists of electrons, positively charged protons, and neutral neutrons
• Electrons form chemical bonds
• Atomic number: numbers of protons in its nucleus
• Mass number: the sum of the protons and neutrons of an atom
• Isotopes have the same atomic number but di erent mass numbers
• The atomic weight: the average weighted mass of its atoms
• Molecular weight: the sum of the atomic weights of all the atoms in the molecule
- The Distribution of Electrons in an Atom
Quantum mechanics uses the mathematical equation of wave motions to characterize the motion
of an electron around a nucleus
Wave functions or orbitals tell us the energy of the electron and the volume of space around the
nucleus where an electron is most likely to be found
The atomic orbital closer to the nucleus has the lowest energy
Degenerate orbitals have the same energy
The rst shell is closest to the nucleus.
The closer the atomic orbital is to the nucleus, the lower its energy.
Within a shell, energy of s < p.
- Relative Energies of the Atomic Orbitals
Aufbau principle: An electron goes into the atomic orbital with the lowest energy.
Pauli exclusion principle: No more than two electrons can be in an atomic orbital.
Hund’s rule: An electron goes into an empty degenerate orbital rather than pairing up.
-> Core and valence electrons
Lewis’s theory: an atom will give up, accept, or share electrons in order to achieve a lled outer
shell or an outer shell that contains eight electrons. + is lost electron, - is gained.
- A Hydrogen Atom Can Lose or Gain an Electron
A hydrogen atom achieves an empty shell by losing an electron or a lled outer shell by gaining an
electron.
- Bonding
Ionic Bonds Are Formed by the Transfer of Electrons
An ionic bond is the attraction between ions of opposite charge
Attractive forces between opposite charges are called electrostatic attractions
Achieving a lled outer shell by sharing electrons. A bond formed by sharing electrons is called a
covalent bond.
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,- Nonpolar and Polar Covalent Bonds
Nonpolar covalent bond = bonded atoms are the same or have similar electronegativities.
Polar covalent bond = bonded atoms have di erent electronegativities.
Equal sharing of electrons: nonpolar covalent bond (e.g., H2)
Sharing of electrons between atoms of di erent electronegativities: polar covalent bond (e.g., HF)
- The Greater the Di erence in Electronegativity, the More Polar the Bond
Nonpolar covalent bond: electonegativity di erence < 0.5
Polar covalent bond: electonegativity di erence 0.5 – 1.9
Electronegativity di erence > 1.9: electrons are not shared; atoms are held together by the
attraction of opposite charges
- Dipole moment
Dipole moment = size of the charge x the distance between the charges
The greater the di erence in electronegativity, the greater the dipole moment and the more polar
the bond.
Red has the most negative electrostatic potential; attracts positive charge. Blue has the most
positive electrostatic potential; attracts negative charge.
- Formal charge
Formal Charge = the # of valence electrons – (the # of lone-pair electrons + the # of bonds)
- Carbon forms four bonds
If carbon does not form four bonds, it has a charge (or it is a radical).
- Nitrogen forms three bonds
Nitrogen has one lone pair. If nitrogen does not form three bonds, it is charged.
- Oxygen forms two bonds
Oxygen has two lone pairs. If oxygen does not form two bonds, it is charged.
- Hydrogen and the Halogens Form One Bond
A halogen has three lone pairs. If hydrogen or halogen does not form one bond, it has a charge (or
it is a radical).
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, - Organic chemistry
Organic compounds are compounds containing carbon
Carbon neither readily gives up nor readily accepts electrons
Carbon shares electrons with other carbon atoms as well as with several di erent kinds of atoms
- Shape of atomic orbitals
An orbital tells us the volume of space around the nucleus where an electron is most likely to be
found
The s Orbitals
A node is at a region where a standing wave has an amplitude of zero
The p Orbitals
- Molecular orbitals
Molecular orbitals belong to the whole molecule
sigma bond: formed by overlapping of two s orbitals or by end-on overlap of two p orbitals
A sigma bond is stronger than alpha bond
Bond strength/bond dissociation: energy required to break a bond or energy released to form a
bond
- Atomic orbitals can combine in two ways
In-phase overlap forms a bonding MO; out-of-phase overlap forms an antibonding MO
Pi bond (p) is formed by sideways overlap of two parallel p orbitals
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