Chapter 8: Basic Concept of Chemical Bonding

Valence Electrons• The electrons in the highest occupied level of an element’s

atoms. • The number of valence electrons largely determines the

chemical properties of an element. • Can be determined by looking at the electron configuration

or Lewis structure (dot diagram) for an element. • Alkali metals have 1, Alkaline Earth metals have 2,

Transition metals have 2 or 3, Boron group has 3, Carbon group has 4, Nitrogen group has 5, Chalcogens have 6, Halogens have 7, and the Noble Gases have 8 (an octet).

• Chemist Gilbert Lewis explained that atoms form certain kinds of bonds due to atoms tendency to achieve the electron configuration of a noble gas (octet rule-Atoms tend to gain lose, or share electrons until they are surrounded by eight valence electrons).

Electronegativity

• The tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of other elements.

Ionic Bonding• Two elements with very different electronegativities

(difference between 4.0 and 1.7).• Elements with high electronegativity pull electron(s) more

strongly and therefore steal the electron(s) from the element with a low electronegativity.

• The stronger elements therefore becomes negatively charged (anion) since it has gained electrons and the weaker elements becomes positively charged (cation) since it has lost electrons.

• Since they have opposite charges, they attract one another by electrostatic forces which creates the bond.

• Usually between one non-metal and one metal. • Lewis Structure for Ionic Bonds:

Na Cl become Na+[ Cl ]-

Ionic Bond Characteristics• Compounds are called formula units• Most are crystalline solids at room temperature (pg 422 in regular chem

text has some great pictures)• In solids NaCl, each sodium cation is surrounded by six chloride anions

and each chlorine anion is surrounded by six sodium cations. Giving them both a coordination number (the number of ions of opposite charge that surrounds the ion in a crystal) of 6.

• In this arrangement, each ion is attracted strongly to each of its neighbors and repulsions are minimized, which results in a very stable structure.

• High melting point (usually above 300 °C)• High solubility (dissolving ability) in water since water is polar.• Good conductor of electricity either when melted (the orderly crystal

structure breaks down allowing cations and anions to move freely through the liquid) or when mixed in water (ions also become free to move around allowing electrons to flow).

• Brittle, not very malleable because when struck with a hammer, the blow tends to push ions of like charge into contact and they repel, shattering the crystal.

Energetics of Ionic Bond Formation• When atoms lose electrons to form ions, energy is required (taken in) so it

is an endothermic process. For example, when Na becomes Na+ it requires 496 kJ/mol.

• When atoms gain electrons to form ions, energy is given off so it is an exothermic reaction. For example, when Cl becomes Cl- it gives off 349 kJ/mol.

• Therefore, if the transfer of an electron from one atom to another were the only factor in forming an ionic bond, the overall process would be an endothermic process the requires 496-349= 147 kJ/mol. But this is only if the ions formed far apart from each other.

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

• A measure of just how much stabilization results from arranging oppositely charged ions in an ionic solid is given by the lattice energy, which is the energy required to completely separate a mole of a solid ionic compound into its gaseous state. For example, the lattice energy for NaCl is 788kJ/mol, which means that NaCl needs to take in that much energy to be able to make the ions break apart. Therefore, when they come together the opposite occurs…788 kJ/mol of energy is given off so the formation process is actually very exothermic. (See picture on pg 300)

• Lattice energies of given in Table 8.2 on pg 301 actually increase as the charges on the ions increase and as their radii decrease.

Covalent Bonding

• Two elements with similar electronegativities (difference is less than 1.7).

• Since they have similar pull on the electron(s) they end up both holding onto it and the electron(s) and sharing them, which creates the bond.

• Usually both non-metals.• Lewis Structure for covalent bond:

H H becomes H H

Covalent Bonding Characteristics

• Compounds are called molecules.• Can be solid, liquid or gas at room temperature.• Low melting point (usually below )• High to low solubility in water (depending on the

molecules’ polarity)• Poor to nonconducting when dissolved in water

(aqueous). • A greater range of physical properties is found due to the

widely varying intermolecular attractions (polar and nonpolar bonding)

• Can have single, double or triple covalent bonds between two atoms.

Polarity

• Polar bonds are when one atoms has a slightly stronger pull on the electrons that are being shared (difference between 1.7 and 0.3). Therefore, one end of the bond is slightly more negative (the one that is slightly stronger) and one is slightly more positive (the one that is slightly weaker).

• All the diatomic elements (Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, Iodine and Hydrogen) have non-polar covalent bonds since they share the electrons equally.

HFH2

Polar Bonds vs. Polar Molecules

• The presence of a polar bond in a molecule often makes the entire molecule polar, where one end of the molecule has a slightly more negative charge and the other end is slightly more positive (called a dipole).

(Ex: H20 and NH3)• However, if the polar bonds are arranged

symmetrically around the molecule the entire molecule is actually non-polar.

(Ex: CO2 and CH4)

Drawing Lewis Structures for Covalent Bonds

• Count the total valence electrons for the molecule: To do this, find the number of valence electrons for each atom in the molecule, and add them up.

• Determine the center atom(s). This will be the first atom listed in the formula, unless, the first atom listed is Hydrogen.

• Put other elements in formula around central atom and connect them with single bonds. If more than 4 outer atoms they will have to be bonded to others outer atoms.

• Fill in eight electrons for each non-central atom. There are a few exceptions to this rule we will learn later.

• Fill in eight electrons around central atom. Also some exceptions we will learn later.

• Count how many total electrons you now have on your drawing. If count is less the valence electrons determined earlier add the extras to the central atom. If count is more than the valence electrons determined earlier change some unshared pairs into double or triple bonds.

Formal Charge• The formal charge of any atom in a molecule is the charge

the atom would have if all the atoms in the molecule had the same electronegativity-if all electrons were shared equally.

• To calculate formal charge, we assign the electrons to the atom as follows:– All unshared (nonbonding) electrons are assigned to the atom on

which they are found.– For any bond, single double or triple, half of the bonding electrons are

assigned to each atom in the bond.• The formal charge of each atom is then calculated by

subtracting the number of electrons assigned to the atom from the number of valence electrons in the isolated atom originally.

• Practice website• We generally choose the Lewis structure in which the atoms

bear formal charges closet to zero.• We also generally choose the Lewis structure in which any

negative charges reside on the more electronegative atoms.

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Resonance Structures

• We sometimes encounter molecules and ions in which the experimentally determined arrangement of atoms is not adequately described by a single Lewis structure.

• Consider a molecule of ozone, O3, which is a bent molecules with two equal O—O bond lengths.

• See example on page 319 in book. • See website example

Exceptions to the Octet Rule• Three main types of exceptions for covalent

bonding:– Molecules and polyatomic ions containing an

odd number of electrons– Molecules and polyatomic ions in which an

atom has fewer than an octet of valence electrons.

– Molecules and polyatomic ions in which an atom has more than an octet of valence electrons.

Molecules and polyatomic ions containing an odd number of electrons

• In some molecules, the total number of valence electrons is odd when two atoms bond. For example, NO since N has 5 and O has 6, the total is 11 so an octet could never be achieved as you need to have an even number for that.

Molecules and polyatomic ions in which an atom has fewer than an octet of

valence electrons.• Hydrogen, Beryllium and Boron have two few valence electrons

to ever obtain a full octet. See this. • Hydrogen can have at most 2 valence electrons after it shares

its electron with another atom. • Beryllium will have 4 valence electrons after it has finished

bonding.

• Boron will have 6 valence electrons after it shares its valence

electrons with other atoms.

Molecules and polyatomic ions in which an atom has more than an octet of

valence electrons• Elements in periods 3, 4, 5, 6 and 7 can

expand their octet to have 10, 12, or 14 valence electrons. See how this can happen.

Strengths of Covalent Bonds

• The stability of a molecule is related to the strengths of the covalent bonds it contains.

• The strength of a covalent bond between two atoms is determined by the energy required to break that bond.

• It is easiest to relate bond strength to the enthalpy change in reactions in which bonds are broken.

• The bond enthalpy is the enthalpy change, ΔH, for the breaking of a particular bond in one mole of a gaseous substance… which we will cover in Chapter 5 (Section 5.3).

Metallic Bonding

• Metals are made up of closely packed cations rather than neutral atoms. The cations are surrounded by mobile valence electrons (delocalized electrons), which drift freely from one part of the metal to another.

• The attractions of the free-floating valence electrons for the positively charged metal cations creates the bond that holds the metal together.

Metallic Bond Characteristics• Good conductors of electrical current because electrons can flow

freely in them. As electrons enter one end of a bar of metal, an equal number leave the other end.

• Ductile and malleable due to the metals ability to be forced through a narrow opening to produce a wire or shaped with a hammer without shattering. A sea of drifting valence electrons insulates the metal cations from one another. When a metal is subjected to pressure, the metal cations easily slide past one another like ball bearing immersed in oil.

• Also crystalline solids at room temperature (except for mercury). Three arrangements of the atoms– Body-centered cubic– Face-centered cubic– Hexagonal close-packed

Other Intermolecular Forces

• Molecules are often attracted to each other by a variety of other forces, although weaker than ionic or covalent bonds. – Van der Waals forces:

• Dispersion forces-the weakest of all molecular interactions caused by the motion of electrons.

• Dipole interactions-occurs when polar molecules are attracted to one another

– Hydrogen bonds: • Attractive forces in which a hydrogen covalently bonds to a very

electronegative atom is also weakly bonded to an unshared pair of electrons of another electronegative atom. Since hydrogen only has one electron when it bonds with a highly electronegative atom, the nucleus of the hydrogen atom becomes very electron deficient. Therefore, the sharing of a nonbonding electron pair on a nearby electronegative atom compensates for the deficiency.