Chemistry Chapter 8

Ionic Bonds

Due to the electrostatic attractions between oppositely charged ions.

How does the electronegativity of an element differ from its electron affinity?

EA can either increase or decrease across a period, while electronegativity only increases.

Bonding Pair

Each shared electron pair, shown as a line.

Electrostatic Potential Energy

Eel = (kQ1Q2)/d Q1 & Q2 - The charges on the particles (Coulombs) d - The distance between their centers in meters k - constant = 8.99 x 10^9 J-m/C^2

For the group 6A elements, what is the trend in electronegativity with increasing atomic number?

Electronegativity generally decreases with increasing atomic number.

Polar Molecule

A molecule (such as HF) in which the centers of positive and negative charge do not coincide. - Both bonds and entire molecules are described as being polar and nonpolar. - Polar molecules align themselves with respect to one another, with the negative end of one molecule and the positive end of another attracting each other. - Polar molecules are attracted to ions. - The negative end of a polar molecule is attracted to a positive ion, and the positive end is attracted to a negative ion.

Single Bond

A shared electron pair which consists of a single covalent bond.

Nonpolar Covalent Bond

A bond in which the electrons are shared equally.

Key Characteristics of Ionic Substances

- Usually brittle. - High melting points. - Usually crystalline. - Can be cleaved.

Predict whether the N-O bonds in NO3- are stronger or weaker than the N-O bond in NO+.

- In all resonance Lewis structures, only the electron placement differs. Factors such as bond lengths stay exactly the same.

Lattice Energy Periodic Trend

- The attractive interaction between two oppositely charged ions increases as the magnitudes of their charges increase and as the distance between their centers decreases. - For a given arrangement of ions, the lattice energy increases as the charges on the ions increase and as their radii decrease. - The variation in the magnitude of lattice energies depends more on ionic charge than on ionic radius because ionic radii vary over only a limited range compared to charges.

The magnitude of the lattice energy of an ionic solid depends on:

- The charges of the ions. - Their sizes. - Their arrangement in the solid.

Measurement of the dipole moments can provide us with valuable information about the charge distribution in molecules.

- The electronegativity difference decreases as the bond length increases. - The varying degree of electronic charge shift can be visualized based on calculations of electron distribution. - In figure 8.11, the change in the electronegativity difference for these molecules has a greater effect on the dipole moment than does the change in bond length.

Electron Configurations of Ions of the s- and p- Block Elements

- The energetics of ionic bond formation helps explain why many ions tend to have noble-gas electron configuration. Na 1s2 2s2 2p6 3s1 = [Ne] 3s1 Na+ 1s2 2s2 2p6 = [Ne] - Lattice energy increases with increasing ionic charge. - The increase in lattice energy is NOT enough to compensate for the energy needed to remove an inner-shell electron. - Adding electrons is either exothermic or only slightly endothermic AS LONG AS the electrons are added to the valence shell. Cl 1s2 2s2 2p6 3s2 3p5 = [Ne] 3s2 3p5 Cl- 1s2 2s2 2p6 3s2 3p6 = [Ne] 3s2 3p6 = [Ar] - Noble gases are stable, so it is VERY UNFAVORABLE to form a Cl2- ion.

Energetics of Ionic Bond Formation - What factors make the formation of ionic compounds so exothermic?

- The energy released by the attraction between ions of unlike charge more than makes up for the endothermic nature of ionization energies, making the formation of ionic compounds an exothermic process. - The strong interactions cause most ionic materials to be hard, brittle, materials with high melting points. - The attraction between ions of opposite charge makes ionic compounds stable, which causes the ions to draw together, releasing and causing many ions to form a solid array, or lattice. - The energy released by the attraction between ions of unlike charge more than makes up for the endothermic nature of ionization energies, making the formation of ionic compounds an exothermic process. - Most ionic materials are hard and brittle. - High melting points (NaCl at 801 degrees C).

Transition Metal Ions

- The lattice energies of ionic compounds are generally large enough to compensate for the loss of up to only 3 electrons from atoms. - Cations with charges of 1+, 2+, 3+ in ionic compounds. - Most transition metals have more than 3 electrons beyond a noble gas core. Silver: [Kr] 4d10 5s1 - Metals of Group 1B (Cu, Ag, Au) often occur as 1+ ions. - Transition metals generally do not form ions that have a noble-gas configuration, which limits the octet rule. - In forming ions, transition metals lose the valence-shell s electrons first, then as many d electrons as required to reach the charge of the ion. Fe: [Ar] 3d6 4s2 Fe2+: [Ar] 3d6 Fe3+: [Ar] 3d5

Strengths and Lengths of Covalent Bonds

- The stability of a molecule is related to the strengths of its covalent bonds. - Multiple bonds are generally stronger than single bonds. - As the number of bonds between two atoms increases, the bond grows shorter and stronger.

Visualizing Figure 8.4

- The structure expands from within. - Distances between the ions increase until the ions are very far apart. NaCl(s) -> Na+(g) + Cl-(g) Delta Hlattice = +788 kJ/mol - Process is highly endothermic. - Reverse process is highly exothermic.

Comparing Ionic and Covalent Bonding

- There is a continuum between the extremes of ionic and covalent bonding. - When covalent bonding is dominant, we expect compounds to exist as molecules, having all the properties we associate with molecular substances. - Relatively low melting and boiling points. - Non-electrolyte behavior when dissolved in water. - When ionic bonding is dominant, we expect the compounds to be brittle, high-melting solids with extended lattice structures, exhibiting strong electrolyte behavior when dissolved in water. - Assume that the interaction between a metal and a nonmetal is ionic and that between two nonmetals is covalent. - Use the difference in electronegativity as the main criteria for determining whether ionic or covalent bonding will be dominant. - The electronegativity values (Fig. 8.8) do not take into account changes in bonding that accompany changes in the oxidation state of the metal. - In general, as the oxidation state of a metal increases, so does the degree of covalent bonding. - Metals in high oxidation states form molecular substances rather than ionic compounds.

What ions are formed by Groups 1A, 2A, 3A?

1+, 2+, 3+ cations.

How to Calculate Formal Charges of Atoms in Lewis Structures

1. All unshared (nonbonding) electrons are assigned to the atom on which they are found. 2. For any bond - single, double, or triple - half of the bonding electrons are assigned to each atom in the bond. 3. The formal charge of each atom is calculated by subtracting the number of electrons assigned to the atom from the number of valence electrons in the neutral atom: *Formal Charge = valence electrons - 0.5(bonding electrons) - nonbonding electrons*

Exceptions to the Octet Rule

1. Odd number of electrons. 2. An atom has fewer than an octet of valence electrons. 3. An atom has more than an octet of valence electrons.

How to Draw Lewis Structures

1. Sum the valence electrons from all atoms, taking into account overall charge. 2. Write the symbols for the atoms, show which atoms are attached to which, and connect them with a single bond (a line, representing two electrons). 3. Complete the octets around all the atoms bonded to the central atom. 4. Place any remaining electrons on the central atom. 5. If there are not enough electrons to give the central atom an octet, try multiple bonds.

How to Identify the Dominant Lewis Structure

1. The dominant Lewis structure is generally the one in which the atoms bear formal charges closest to zero. 2. A Lewis structure in which any negative charges reside on the more electronegative atoms is generally more dominant than one that has negative charges on the less electronegative atoms.

What ions are formed by Groups 5A, 6A, 7A?

3-, 2-, 1- anions.

Debyes

3.34 x 10^-30 C-M

Polar Covalent Bond

A bond in which one of the atoms exerts a greater attraction for the bonding electrons than the other.

Covalent Bond

A chemical bond formed by sharing a pair of electrons. - In a stable molecule, the attractive forces must overcome the repulsive ones. - Most substances seen daily are gases, liquids, or solids with low melting points. - By using quantum mechanical methods analogous to those used for atoms in section 6.5, we can calculate the distribution of electron density in molecules. - The shared pair of electrons in any covalent bond acts as a kind of "glue" to bind atoms together.

Bond Polarity

A measure of how equally or unequally the electrons in any covalent bond are shared. - When two identical atoms bond, the electron pairs must be shared equally.

Formal Charge and Alternative Lewis Structures

All the possible lewis structures of an atom can be thought of as contributing to the actual arrangement of the electrons in the molecule, but not all of them will contribute to the same context.

The Octet Rule

Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons. - An octet of electrons consists of full s- and p- subshells in an atom. - The rule provides a useful framework for introducing many important concepts of bonding.

Cleaved

Broken apart along smooth, flat surfaces.

The C-O bond length in carbon monoxide, CO, is 1.13 A, whereas the C-O bond length in CO2 is 1.24 A. Without drawing a Lewis structure, do you think that CO contains a single, double, or triple bond?

CO probably has a triple bond.

Lewis Symbol

Consist of an element's chemical symbol plus a dot for each valence electron. - Dots are placed top->bottom->left->right. - The number of valence electrons in any representative element is the same as the element's group number.

Triple Bond

Corresponds to the sharing of three pairs of electrons.

Chemical Bond

Created whenever two atoms/ions are held strongly together.

Linus Pauling

Developed the first and most widely used electronegativity scale, which is based on thermochemical data. - Electronegativity increases from left to right and decreases from up to down.

Dipole

Established whenever two electrical charges of equal magnitude but opposite sign are separated by a distance.

Metallic Bonds

Formed by electrons that are relatively free to move from one atom to another.

Ionic Substances

Generally result from the interaction of metals on the left side of the periodic table with nonmetals on the right (excluding noble gases). Na(s) + 1/2Cl2(g) -> NaCl(s) - Sodium Chloride is composed of Na+ and Cl- ions arranged in a 3-D array. - There has been an electron transfer from the Na atom to the Cl atom.

Chlorine mono fluoride, ClF, and iodine monofluoride, IF, are interhalogen compounds - compounds that contain bonds between different halogen elements. Which of these molecules has the larger dipole moment?

IF has the larger dipole moment because of the increased distance (on the periodic table) between the two molecules.

Ionizing an H2 molecule to H2+ changes the strength of the bond. Based on the description of covalent bonding given previously, do you expect the H-H bond in H2+ to be weaker or stronger than the H-H bond in H2?

It will be weaker because of the loss of an electron.

(Figure 8.1) If the white powder were sugar, C12H22O11, how would we have to change this picture?

It would change from Metallic and Ionic to Non-Metallic and Covalent because sugar, C12H22O11, is not a metal.

Hypervalent

Molecules and ions with more than an octet of electrons around the central atom. - Formed only for central atoms from period 3 and below in the periodic table. - The principle reason for their formation is the relatively larger size of the central atom. - In elements of the second period, only the 2s and 2p valence orbitals are available for bonding. - The presence of unfilled 3d orbitals in P and S has a relatively minor impact on the formation of hypervalent molecules. - the larger size of the atoms from period 3 through 6 is more important to explain hypervalency than is the presence of unfilled d orbitals. - In general, multiple Lewis structures can contribute to the actual electron distribution in an atom or molecule.

Covalent Bonds

Molecules are formed by the sharing of electrons between atoms.

(Figure 8.4) If no color key were provided, how would you know which color ball represented Na+ and which represented Cl-?

Na+ has lost an electron, so its shape would be smaller than normal. Cl-, however, has gained an electron, so its shape would be larger.

If you were to perform the reaction KCl(s) -> K+(g) + Cl-(g), would energy be released?

No, energy will not be released. Instead, energy would be absorbed because it takes energy to break apart chemical bonds.

Electron Transfer

Occurs when one atom readily gives up an electron (low IE - ionization energy) and another atom readily gains an electron (high EA - electron affinity).

Which element forms a 3+ ion that has the electron configuration [Kr] 4d6?

Original: [Kr] 5s2 3s7 Original3+: [Kr] 4d6 Answer: Rh, Rhodium

G.N. Lewis

Suggested a simple way of showing the valence electrons in an atom and tracking them during bond formation, using what are now known as either Lewis electron-dot symbols or simply Lewis symbols.

As you go from HF to HI, does the H-X bond become more or less polar?

The H-X bond becomes less polar. - Even in ionic compounds, there is still some covalent contribution to the bonding.

Electronegativity

The ability of an atom in a molecule to attract electrons to itself. - Used to estimate whether a given bond is nonpolar covalent, polar covalent, or ionic. - The greater an atom's electronegativity, the greater its ability to attract electrons to itself. - Related to the atom's ionization energy and electron affinity, which are both properties of isolated atoms. - An atom with a very negative electron affinity and a high ionization energy both attracts electrons from other atoms and resists having its electrons attracted away; therefore, it is highly electronegative.

Formal Charge

The charge an atom would have if each bonding electron pair in the molecule were shared equally between its two atoms. - The sum of the formal charges equals the overall charge on the ion. - The formal charges on a neutral molecule must add to zero, whereas those on an ion add to give the charge on the ion. - Most important lewis structure = dominant. - Formal charges do not represent real charges on atoms. (Think "book keeping.") - The actual charge distributions are instead determined by a number of other factors, including electronegativity differences between atoms.

(Figure 8.7)What would happen to the concentration of electron density between the nuclei in (b) if you pulled the nuclei further apart?

The concentration of electron density would decrease because there would be less tension between the two molecules (AKA magnets).

If the charged particles are moved closer together, does u increase, decrease, or stay the same?

The dipole moment would decrease because of the decreased distance between the two particles.

Valence Electrons

The electrons involved in chemical bonding.

Lattice Energy

The energy required to completely separate one mole of a solid ionic compound into its gaseous ions. - Gives a measure of how much stabilization results from arranging oppositely charged ions in an ionic solid.

Are all these Lewis symbols for Cl correct (2 showing 7 VE and the last showing 5 VE)?

The first two dot structures are correct, but the third is not because it shows Chlorine to have only 5 valence electrons when it should have 7.

(Figure 8.5) Using the figure, find the most likely range of values for the lattice energy of KF.

The lattice energy of KF will most likely fall between the values of 701 (the lattice energy for KCl) and 910 (the lattice energy for NaI).

Multiple Bonds

The length of the bond between two atoms decreases as the number of shared electron pairs increases.

Resonance Structures

The placement of the atoms in two alternative but completely equivalent Lewis structures, but the placement of electrons is different. *Mixing point analogy.*

Dipole Moment

The quantitative measure of the magnitude of a dipole. - The dipole moment increases as the magnitude of q increases and r increases. - The larger the dipole moment, the more polar the bond. - For a nonpolar molecule, such as F2, the dipole moment is zero because there is no charge separation.

Goal of Chapter 8

To examine the relationship between the electronic structure of atoms and the ionic and covalent chemical bonds they form.

Lone/Nonbonding Pairs

Unshared electron pairs, shown as dots. - For nonmetals, the number of valence electrons is the same as the group number.

Ionic Covalent Bond

When two atoms differ in electronegativity by more than 2.0. - The shift of electron density toward the more electronegative atom in a bond can be seen from the results of calculations of electron-density distributions. - The greater the difference in electronegativity between two atoms, the more polar their bond.

Double Bond

When two lines are drawn in a Lewis structure to represent two electron pairs shared by two atoms.

The bond between carbon and hydrogen is one of the most important types of bonds in chemistry. The length of an H-C bond is approximately 1.1 A. Based on this distance and differences in electronegativity, do you expect the dipole moment of an individual H-C bond to be larger or smaller than that of an H-I bond?

Yes, I would expect that outcome.

(Figure 8.3) Do you expect a similar reaction between potassium metal and elemental bromine?

Yes, because K and Br are also on opposite sides of the periodic table, similar to how Na and Cl are situated.