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Chemical Bonding I: The Covalent Bond

Chemical Bonding I: The Covalent Bond. Lewis Structures. Atoms react with one another to form molecules in order to achieve a more stable electronic configuration The valence electrons are those of greatest interest because as they are shared between two atoms in a covalent bond

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Chemical Bonding I: The Covalent Bond

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  1. Chemical Bonding I: The Covalent Bond

  2. Lewis Structures • Atoms react with one another to form molecules in order to achieve a more stable electronic configuration • The valence electrons are those of greatest interest because as they are shared between two atoms in a covalent bond • In the Lewis approach, an atom is represented by its chemical symbol surrounded by points that represent the valence electrons • Because members of the same group in the periodic table have the same number of valence electrons, they share the same Lewis representation (except for He within the noble gases where it has two valence electrons as opposed to eight) • The Lewis approach does not work well for transition metals, lanthanides, or actinides

  3. Lewis Dot Symbols

  4. The Covalent Bond • A covalent bond is a bond in which two electrons are shared by two atoms • In a covalent bond, each of the electrons in the shared pair is attracted by the nuclei of the two atoms • These attractions help keep the two atoms together • Only valence electrons participate in covalent bonds

  5. Lewis Structures • Electrons not participating in the formation of a covalent bond are called non-bonding electrons or lone pairs • In Lewis structures, covalent bonds are illustrated by dashes or by pairs of points between two atoms and the lone pairs are illustrated by pairs of points associated with individual atoms

  6. The Octet Rule • Lewis proposed that any atom, except hydrogen, tends to form bonds until it is surrounded by eight valence electrons • this is the octet rule • hydrogen seeks to have two electrons rather than eight • Beyond the second period (Li to F), there are exceptions to the octet rule due to the availability of the d orbitals within the same shell as the s and p valence orbitals

  7. Multiple Bonds • If two atoms share a pair of electrons, they form a single bond • If two atoms share two pairs of electrons, they form a double bond • If two atoms share three pairs of electrons, they form a triple bond • For a given pair of atoms, triple bonds are shorter and more stable than the double bonds, which in turn are shorter and more stable than single bonds

  8. Polar Covalent Bond • In a homonuclear diatomic molecule (like H2 or F2), bonding electrons are perfectly shared • In a heteronuclear diatomic molecule (like HF), the sharing is not done equally, i.e., the electrons are closer to one atom than the other • The bond is said to be polar covalent (or simply, polar) • In an ionic bond, the transfer of the electron is almost complete • In a polar bond, there is still significant sharing

  9. Electronegativity • Electronegativity is the tendency of an atom to attract the electrons in a chemical bond towards itself • Electronegativity is a relative value, and is therefore unitless • The more electronegative an element is, the more this element tends to attract electrons • An element that has a strong electron affinity and high ionization energy tends to have a high electronegativity • Pauling first established a method for estimating the electronegativity of most elements

  10. Electronegativity

  11. Polar Covalent Bonds and Ionic Bonds • A bond between a metal and a non-metal tends to be ionic • A bond between two non-metallic elements tends to be polar covalent • General Rule: • If the difference in electronegativity is equal to or greater than 2.0, the bond is ionic • If the difference in electronegativity is less than 2.0, the bond is polar covalent • If the bond is between two atoms of the same element, the bond is pure covalent

  12. Electronegativity and Oxidation State • The oxidation state indicates the charge that an atom in a molecule would have if the electrons were transferred completely to the more electronegative of the atoms participating in the bond • e.g.; In water, O is more electronegative than H, therefore O takes the electron from the H in each H:O bond and has a charge of -2, and each H has a charge of +1 • e.g.; In hydrogen peroxide, each O has an oxidation state of -1 since each O takes the electron from the H in the H:O bond, but they share the electrons perfectly in the O:O bond • In a compound that contains F, the F is always -1 since it is the most electronegative element

  13. The Rules for Writing Lewis Structures • Step 1: Establish the “skeleton” structure of the compound using the chemical symbols and placing bonded atoms side by side • In cases of doubt, the least electronegative atom occupies the central position • Step 2: Calculate the total number of valence electrons • In the case of an anion, add the absolute value of the negative charge to the total electron count • In the case of a cation, subtract the value of the positive charge from the total electron count

  14. The Rules for Writing Lewis Structures • Step 3: Draw a single covalent bond between the central atom and each of the atoms surrounding it • As much a possible, complete the octets of the atoms bound to the central atom (except for hydrogen, which only needs two electrons) • The number of electrons at the end of this step must be the same as the number that was determined in step 2 • Step 4: If an atom does not satisfy the octet rule, displace lone pairs from atoms to which it is bound so as to form double and triple bonds

  15. Lewis Structures • Example: Write the Lewis structure of the carbonate ion (CO32-).

  16. Lewis Structures • Example Write the Lewis structure of the nitrite ion (NO2-).

  17. Formal Charges in Lewis Structures • For an isolated atom, the number of electrons associated with it simply corresponds to the number of valence electrons that it possesses as an isolated atom • In a molecule, both electrons in a lone pair are assigned to the atom on which it is located while the electrons involved in a bond are split equally amongst the two atoms involved in the bond • The formal charge of an atom in a molecule is the difference between the number of valence electrons contained in an isolated atom and the number of electrons assigned to the same atom in a Lewis structure of that molecule

  18. Formal Charges in Lewis Structures • Formal Charge = number of valence electrons in the isolated atom - number of non-bonding electrons - 1/2 (number of bonding electrons) • In the case of a neutral molecule, the sum of the formal charges must equal zero • In the case of an ion, the sum of the formal charges must equal the charge of the ion • N.B. The formal charges do not represent the real distribution of charge in a molecule • e.g.; For H2O, we do not have O2- and H+ ions , but we know that the electronegativity of O is greater than that of H, so O has a slightly negative charge and the H’s have slightly positive charges

  19. Formal Charges in Lewis Structures • Example: Determine the formal charges in a carbonate ion. • Solution: • For the C atom: • For the O atom in the C=O bond • For the O atoms in a C-O bond • N.B. The sum of the formal charges (-2) is the charge of the ion Formal charge Formal charge Formal charge

  20. Formal Charges in Lewis Structures • If there are several Lewis structures that obey the octet rule, we can use the formal charges to determine which structure is the best • In the case of a neutral molecule, a Lewis structure which has no formal charges is preferable to another that does have formal charges • A Lewis structure that has higher formal charges (2, 3, etc.) is not as good as one that has lower formal charges • If Lewis structures have a similar distribution of formal charges, the best structure is the one in which negative formal charges are placed on the most electronegative atoms

  21. The Formal Charge and Lewis Structures • Example: Which of these Lewis structures for N2O is the best? • Solution: The third structure is the worst since the terminal N atom has a formal charge of -2. The first structure is better than the second structure because O is more electronegative than N and the first has the negative formal charge on the O while the second puts is on the N.

  22. Resonance • If we look at, for example, a molecule like ozone, O3, there are two different but equivalently good Lewis structures • Each Lewis structure predicts that ozone has a double O=O and a single O-O bond • However, it has been found experimentally that the two O-O bonds in ozone are equivalent (same length, same strength) • In reality, ozone is a hybrid of these two Lewis structures

  23. Resonance • Resonance is the use of two or more Lewis structures to represent a given molecule • Each of the Lewis structures is called a resonance structure • The symbol  indicates that the illustrated structures are resonance structures • Experimentally, we see that the bonds in ozone are shorter and stronger than the O-O single bonds in H2O2 but they are longer and weaker than the O=O double bond in O2

  24. Resonance • Do not believe that a molecule such as the carbonate ion passes successively and quickly from one resonance structure to another • None of the resonance structures adequately represents the real molecule • We use the concept of resonance to explain why the three C-O bonds in the carbonate ion are all identical and intermediate in strength between single and double CO bonds • N.B. If you change the position of atoms between two Lewis structures, this is not a resonance structure (rather, these are two distinct molecules )

  25. Exceptions to the Octet Rule: Incomplete Octet • In some compounds, it is impossible to fill the octet of an atom • e.g.; In the gaseous state, BeH2 forms distinct molecules: H-Be-H where only four electrons surround the Be and it is impossible to fill its octet as we only have 4 valence electrons • e.g.; BF3 is a relatively stable molecule • Even if the last three resonance structures fill the B atom’s octet, experiments indicate that the first structure is the dominant resonance structure

  26. Exceptions to the Octet Rule: Molecules with an Odd Number of Electrons • Certain molecules contain an odd number of electrons • These molecules are free radicals and are typically very reactive • With an odd electron count, it is impossible to obey the octet rule • Examples of there molecules are: • Nitric oxide: • Nitrogen dioxide: • Superoxide Anion:

  27. Exceptions to the Octet Rule: Expanded Octet • If the central atom is within the third period, or beyond, the d orbitals of the central atom can participate in covalent bonds • The d orbitals allow the central atom to accommodate more than eight electrons, allowing for an expanded octet • N.B. Even if an atom can have an expanded octet, it can always choose to obey the octet rule • Examples of these molecules are: Sulfur hexafluoride Phosphorus pentafluoride

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