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8.S: Basic Concepts of Chemical Bonding (Summary)

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  • These is a summary of key concepts of the chapter in the Textmap created for "Chemistry: The Central Science" by Brown et al.


    • Ionic bond: bond formed on the basis of electrostatic forces that exist between oppositely charged ions. The ions are formed from atoms by transfer of one or more electrons
    • Covalent bond: bond formed between two or more atoms by a sharing of electrons
    • Metallic bond: bonding in which the bonding electrons are relatively free to move throughout the 3D structure
    • Electrondot symbols: aka Lewis symbols; simple and convenient way of showing the valence electrons of atoms and keeping tack of them in the course of bond formation
    • The number of valence electrons of any representative element is the same as the column number of the element in the periodic table
    • Octet rule: atoms tend to lose or gain electrons until they are surrounded by 8 valence electrons


    Energies of Ionic Bond Formation

    The formation of ionic compounds is very exothermic

    Removing an electron form an atom, such as Na, is endothermic because energy needs to be used to overcome the attractive forces within the atom. Adding an electron is the opposite process and releases lots of energy

    The principal reason that ionic compounds are stable is the attraction between ions of unlike charge. This attraction draws the ions together, releasing energy and causing the ions to form a solid array (lattice)

    Lattice energy: energy required to separate completely a mole of a solid ionic compounds into its gaseous ions

    Large values of lattice energy mean that the ions are strongly attracted to one another

    Energy released by the attraction between the ions of unlike charges more than makes up for the endothermic nature of ionization energies, making the formation of ionic compounds an exothermic process

    \[E = k \dfrac{Q_1Q_2}{ d}\]

    • E = potential energy of two interacting charged particles
    • Q1 and Q2 = charges on the particles
    • D= distance between the particles
    • K = constant; 8.99 X 109 Jm/C2

    For a given arrangement of ions, the lattice energy increases as the charges of ions increase and as their radii decrease. The magnitude of lattice energies depends primarily on the ionic charges because ionic radii do not vary over a wide range.

    Electron Configurations of Ions

    Many ions tend to have noblegas electron configurations. This is why Na can have a +1 charge, but not a +2 one. Once an ion has reached noblegas configuration, it wants to stay there.

    • Na: 1s2 2s2 2p6 3s1 = [Ne] 3s1
    • Na+: 1s2 2s2 2p6 = [Ne]
    • Na2+: 1s2 2s2 2p5

    Similarly, addition of electrons to nonmetals is either exothermic or slightly endothermic as long as electrons are being added to the valence shell. Further addition of electrons requires tremendous amounts of energy; more than is available form the lattice energy

    • Cl: 1s2 2s2 2p6 3s2 3p5 = [Ne] 3s2 3p5
    • Cl: 1s2 2s2 2p6 3s2 3p6 = [Ar]
    • Cl2: 1s2 2s2 2p6 3s2 3p6 4s1 = [Ar]

    The lattice energies of ionic compounds are generally large enough to compensate for the loss of up to only 3 electrons from atoms. Thus we find cations only having charges of +1, +2, or +3.

    Because most transition metals have more than 3 electrons beyond a noblegas core, attainment of a noblegas configuration for these ions is usually impossible.

    When a positive ion is formed from an atom, electrons are always lost first from the subshell with the largest value of n. Thus, a transition metal always loses the outer s electrons before it loses electrons from the underlying d subshell.

    8.3 Sizes of Ions

    Sizes of ions are important in determining both the way in which the ions pack in a solid and the lattice energy of the solid. It is also a major factor governing the properties of ions in solution

    The size of an electron depends on its nuclear charge, the number of electrons it possesses, and the orbitals in which the outershell electrons reside

    Positive ions are formed by removing 1 or more electrons from the outermost region of the atom. Thus, the formation of a cation not only vacates the most spatially extended orbitals, it also decreases the total electronelectron repulsions. Hence, cations are smaller than the original atoms from which they came.

    The opposite happens when speaking of negative ions. An added electron increases electronelectron repulsions and causes the electrons to spread out more in space.

    For ions of the same charge, size increases as we go down a group


    Ionic substances are usually brittle with high melting points. They are usually crystalline, meaning that they have flat surfaces that make characteristic angles with one another.

    Covalent bond: chemical bond formed by sharing a pair of electrons

    Lewis structure: structure that represents bonding using dots for unpaired electrons and lines for bonds

    For nonmetals, the number of valence electrons is the same as the group number

    Knowing this, we can predict that an element in Group 7A would need one covalent bond in order to get an octet, an element in Group 6A would need two, and so on.

    Multiple Bonds

    • Single bond: sharing of one pair of electrons, one covalent bond
    • Double bond: two shared electrons

    Distance between bonded atoms decreases as the number of shared electron pairs increases


    • Bond polarity: measure of how equally the electrons are shared between the atoms in a chemical bond
    • Nonpolar bond: one in which the electrons are shared equally between the two atoms
    • Polar covalent bond: one of the atoms exerts a greater attraction for the electron than the other


    Used to estimate whether a bond will be nonpolar, polar covalent, or ionic

    Electronegativity: ability of an atom in a molecule to attract electrons to itself

    An atom with a very negative electron affinity and high ionization energy will both attract electrons from other atoms and resist having its electrons attracted away; it will be highly electronegative

    Highest electronegativity = 4.0 (Fluorine), lowest = 0.7 (Cesium)

    Electronegativity increases form left to right, and usually decreases with increasing atomic number in any one group

    Electronegativity and Bond Polarity

    Differences in electronegativities:

    Nonpolar = 0 – 0.4

    Polar covalent = 0.4 – 1.6

    Ionic = > 1.6 (> 50% = ionic)

    δ+ and δ : "delta sign"; represent partial positive and negative charge. The atom with the δ is the more electronegative one


    1. Sum the valence electrons from all atoms. For an anion, add an electron to the total negative charge. For a cation, subtract an electron.
    2. Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond.
    3. Complete the octets of the atoms bonded to the central atom.
    4. Place any leftover electrons on the central atom, even if doing so results in more than an octet
    5. If there are not enough electrons to give the central atom an octet, try multiple bonds


    Resonance structures (resonance forms) are individual Lewis structures in cases where two or more Lewis structures are equally good descriptions of a single molecule. If a molecule (or ion) has two or more resonance structures, the molecule is a blend of these structures. The molecule does not oscillate rapidly between two or more different forms.


    1. Molecules with an odd number of electrons
    2. Molecules in which an atom has less than an octet
    3. Molecules in which an atom has more than an octet

    Odd Number of Electrons

    In a few molecules, such as ClO2, NO, and NO2, the number of electrons is odd. In NO for example, there are 5+6 = 11 valence electron. Hence, complete pairing of these electrons is impossible and an octet around each atom cannot be achieved.

    Less Than an Octet

    Second type of exception occurs when there are fewer than eight electrons around an atom in a molecule or ion. Relatively rare situation; most often encountered in compounds of Boron and Beryllium. For example, let’s consider Boron Trifluoride, BF3

    There are 6 electrons around the Boron atom. We can form a double bond between Boron and any of the 3 Fluorine atoms (3 possible resonance structures)

    However, by doing so, we forced a Fluorine atom to share additional electrons with Boron. This would make the F atom to have a +1 charge, and the Boron atom to have a –1 charge, which is extremely unfavorable.

    We then conclude that the structures containing the double bonds are less important than the one illustrated on the right. Since in this case Boron has only 6 valence electrons, it will react violently with molecules that have an unshared pair of electrons.

    More Than an Octet

    The octet rule works as well as it does because the representative elements usually employ only an ns and three np valenceshell orbitals in bonding, and these hold eight electrons.

    Because elements of the second period have only 2s and 2p orbitals, they cannot have more than an octet of electrons in their valence shells. However, from the third period on, the elements have unfilled nd orbitals that can be used in bonding.

    Size also plays an important role in determining whether an atom can accommodate more than eight electrons. The larger the central atom, the larger the number of atoms that can surround it. The size of the surrounding atoms is also important. Expanded valence shells occur most often when the central atom is bonded to the smallest and most electronegative atoms.


    Bonddissociation energy: aka bond energy; enthalpy change, ΔH, required to break a particular bond in a mole of gaseous substance.

    For polyatomic molecules, we must often utilize average bond energies.

    Bond energy is always positive, the greater the bond energy, the stronger the bond

    A molecule with strong bonds generally has less tendency to undergo chemical change than does one with weak bonds

    Bond Energies and the Enthalpy of Reactions

    ΔH = Σ (bond energies of bonds broken) – Σ (bond energies of bonds formed)

    If ΔH > 0, the reaction is endothermic

    If ΔH < 0, the reaction is exothermic

    Cl – Cl (g) + H – CH3 (g) → H – Cl (g) + Cl – CH3 (g)

    Bonds broken: 1 mol Cl – Cl, 1 mol C – H

    Bonds made: 1 mol H – Cl, 1 mol C – Cl

    ∆H = [D (Cl – Cl) + D(C – H)] [D (H – Cl) + D (Cl – Cl)]

    = (242 kJ + 413 kJ) – (431kJ + 328 kJ)

    = 104 kJ

    Bond Strength and Bond Length

    As the number of bonds between a given element increase, the bond energy increases and the bond length decreases. Hence, the atoms are held more tightly and closely together. In general, as the number of bonds between two atoms increases, the bond grows shorter and stronger.

    8.10 Oxidation Numbers

    Oxidation Numbers: aka Oxidation states; a positive or negative whole number assigned to an element in a molecule or ion on the basis of a set of normal rules; to some degree it reflects the positive or negative character of an atom

    Oxidation numbers do NOT correspond to real charges on the atoms, EXCEPT in the special case of simple ionic substances

    1. The oxidation form of an element in its elemental form is zero.
    2. The oxidation number of a monoatomic ion is the same as its charge. Fore example, the oxidation number of sodium in Na+ is +1, and that of sulfur in S2 is –2
    3. In binary compounds (those with two different elements), the element with greater electronegativity is assigned a negative oxidation number equal to its charge in simple ionic compounds the element. For example, consider the oxidation state of Cl in PCl3. Cl is more electronegative than P. In its simple ionic compounds, Cl appears as the ion Cl. Thus, in PCl3, Cl is assigned an oxidation number of –1.
    4. The sum of the oxidation numbers equals zero for an electrically neutral compound and equals the overall charge of the ionic species. For example, PCl3 is a neutral molecule. Thus, the sum of oxidation number of the P and Cl atoms must equal zero Because the oxidation number of each Cl in this compound is –1 (rule 3), the oxidation number of P must be +3.

    Group 1A elements are +1, Group 2A are +2, and Aluminum is +3.

    The most electronegative element, F, is always found in the –1 oxidation state. Oxygen is usually in the –2 state, however it can be –1 in peroxides.

    Hydrogen has an oxidation number of +1 when it is bonded to a more electronegative element (most nonmetals), and of –1 when bonded to less electronegative elements (most metals)

    Oxidation Numbers and Nomenclature

    Name of the less electronegative element is given first, followed by the name of the more electronegative element modified to have an –ide ending

    Compounds of metals in higher oxidation states tend to be molecular rather than ionic

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