Chemical Bonding I: The Lewis Model

Out

The

Serum Albumin

* Some important nutrients, are not soluble in water, so special carriers must be made to chaperone them to hungry cells. * Serum albumin, is the carrier of fatty acids in the blood. Fatty acids are essential for two major things in your body. They are the building blocks for lipids, which form all of the membranes around and inside cells. * They are also rich sources of energy, and may be broken down inside cells to form ATP.

Metallic

Metal and Metal; Electrons pooled

Table

for

Lewis Theory of Ionic Bonding

---- Title Page ---

Lewis Theory of Covalent Bonding

---- Title Page -----

Lewis Theory: An Overview

1. Valence e- play a fundamental role in chemical bonding. 2. e- transfer leads to ionic bonds. 3. Sharing of e- leads to a covalent bond. 4. e- are transferred or shared to give each atom a noble gas configuration, the octet.

Periodic Trends

Across the period nuclear charge increases so they get smaller. Energy level changes between anions and cations.

Lewis Bonding Theory

Atoms bond because it results in a more stable electron configuration. More stable = lower potential energy Atoms bond together by either transferring or sharing electrons. Usually this results in all atoms obtaining an outer shell with eight electrons. Octet rule There are some exceptions to this rule—the key to remember is to try to get an electron configuration like a noble gas.

Why Do Atoms Bond?

Chemical bonds form because they lower the potential energy between the charged particles that compose atoms. A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms. To calculate this potential energy, you need to consider the following interactions: Nucleus-to-nucleus repulsions Electron-to-electron repulsions Nucleus-to-electron attractions

Small (0 - 4)

Covalent; ex. Cl

Bond Dipole Moments

Dipole moment, m, is a measure of bond polarity. A dipole is a material with a + and − end. It is directly proportional to the size of the partial charges and directly proportional to the distance between them. m = (q)(r) Not Coulomb's law Measured in Debyes, D Generally, the more electrons two atoms share and the larger the atoms are, the larger the dipole moment.

Predicting Ionic Formulas Using Lewis Symbols

Electrons are transferred until the metal loses all its valence electrons and the nonmetal has an octet Numbers of atoms are adjusted so the electron transfer comes out even

Covalent Bonding: Bonding and Lone Pair Electrons

Electrons that are shared by atoms are called bonding pairs. Electrons that are not shared by atoms but belong to a particular atom are called lone pairs. Also known as nonbonding pairs

Crystal Lattice

Electrostatic attraction is nondirectional! No direct anion-cation pair Therefore, there is no ionic molecule. The chemical formula is an empirical formula, simply giving the ratio of ions based on charge balance.

Bonding Theories

Explain how and why atoms attach together to form molecules Explain why some combinations of atoms are stable and others are not - Why is water H2O, not HO or H3O? Can be used to predict the shapes of molecules Can be used to predict the chemical and physical properties of compounds

Dipole Moments

Fill

Lewis Electron-Dot Symbols

For main group elements - The A group number gives the number of valence electrons. Place one dot per valence electron on each of the four sides of the element symbol. Pair the dots (electrons) until all of the valence electrons are used. Example: Nitrogen, N, is in Group 5A and therefore has 5 valence electrons.

Electronegativity Difference and Bond Type

If the difference in electronegativity between bonded atoms is 0, the bond is pure covalent. Equal sharing If the difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent. If the difference in electronegativity between bonded atoms is 0.5 to 1.9, the bond is polar covalent. If difference in electronegativity between bonded atoms is larger than or equal to 2.0, the bond is ionic.

Large (2.0+)

Ionic; ex. NaCl

Energetics of Ionic Bond Formation

Lattice Energy - Energy released when the solid crystal forms from separate ions in the gas state.

Lewis Theory and Ionic Bonding

Lewis symbols can be used to represent the transfer of electrons from metal atom to nonmetal atom, resulting in ions that are attracted to each other and therefore bond.

Lewis Theory of Covalent Bonding

Lewis theory implies that another way atoms can achieve an octet of valence electrons is to share their valence electrons with other atoms. The shared electrons would then count toward each atom's octet. The sharing of valence electrons is called covalent bonding.

Ionic Bonding Model versus Reality

Lewis theory implies that if the ions are displaced from their position in the crystal lattice, repulsive forces should occur. This predicts the crystal will become unstable and break apart. Lewis theory predicts ionic solids will be brittle. Ionic solids are brittle. When struck they shatter.

Covalent Bonding: Model versus Reality

Lewis theory implies that some combinations should be stable, whereas others should not. stable combinations result in "octets" Using these ideas of Lewis theory allows us to predict the formulas of molecules of covalently bonded substances. Hydrogen and the halogens are all diatomic molecular elements, as predicted by Lewis theory. Oxygen generally forms either two single bonds or a double bond in its molecular compounds, as predicted by Lewis theory.

Ionic Bonding Model versus Reality

Lewis theory implies that the attractions between ions are strong. Lewis theory predicts ionic compounds should have high melting points and boiling points because breaking down the crystal should require a lot of energy. The stronger the attraction (larger the lattice energy), the higher the melting point. Ionic compounds have high melting points and boiling points. MP generally > 300 °C All ionic compounds are solids at room temperature.

Ionic Bonding Model versus Reality

Lewis theory implies that the positions of the ions in the crystal lattice are critical to the stability of the structure. Lewis theory predicts that moving ions out of position should therefore be difficult, and ionic solids should be hard. Hardness is measured by rubbing two materials together and seeing which "streaks" or cuts the other. The harder material is the one that cuts or doesn't streak. Ionic solids are relatively hard. Compared to most molecular solids

Ionic Bonding Model versus Reality

Lewis theory implies that, in the liquid state or when dissolved in water, the ions will have the ability to move around. Lewis theory predicts that both a liquid ionic compound and an ionic compound dissolved in water should conduct electricity. Ionic compounds conduct electricity in the liquid state or when dissolved in water.

Covalent Bonding: Model versus Reality

Lewis theory of covalent bonding implies that the attractions between atoms are directional. The shared electrons are most stable between the bonding atoms. Therefore, Lewis theory predicts covalently bonded compounds will be found as individual molecules. Rather than an array like ionic compounds Compounds of nonmetals are made of individual molecule units.

Covalent Bonding: Model versus Reality

Lewis theory predicts that neither molecular solids nor liquids should conduct electricity. There are no charged particles around to allow the material to conduct. Molecular compounds do not conduct electricity in the solid or liquid state. Molecular acids conduct electricity when dissolved in water, but not in the solid or liquid state, due to them being ionized by the water.

Covalent Bonding: Model versus Reality

Lewis theory predicts that the hardness and brittleness of molecular compounds should vary depending on the strength of intermolecular attractive forces. The kind and strength of the intermolecular attractions varies based on many factors. Some molecular solids are brittle and hard, but many are soft and waxy.

Covalent Bonding: Model versus Reality

Lewis theory predicts that the melting and boiling points of molecular compounds should be relatively low. This involves breaking the attractions between the molecules, but not the bonds between the atoms. The covalent bonds are strong, but the attractions between the molecules are generally weak. Molecular compounds have low melting points and boiling points. MP generally < 300 °C Molecular compounds are found in all three states at room temperature.

Covalent Bonding: Model versus Reality

Lewis theory predicts that the more electrons two atoms share, the shorter the bond should be. Bond length is determined by measuring the distance between the nuclei of bonded atoms. In general, triple bonds are shorter than double bonds, and double bonds are shorter than single bonds.

Covalent Bonding: Model versus Reality

Lewis theory predicts that the more electrons two atoms share, the stronger the bond should be. Bond strength is measured by how much energy must be added into the bond to break it in half. In general, triple bonds are stronger than double bonds, and double bonds are stronger than single bonds. However, Lewis theory would predict that double bonds are twice as strong as single bonds, but the reality is they are less than twice as strong.

Lewis Theory Predictions for Ionic Bonding

Lewis theory predicts the number of electrons a metal atom should lose or a nonmetal atom should gain in order to attain a stable electron arrangement. The octet rule This allows us to predict the formulas of ionic compounds that result. It also allows us to predict the relative strengths of the resulting ionic bonds from Coulomb's law.

Ionic

Metal and Nonmetal; Electrons Transferred

Bond Polarity

Most bonds have some degree of sharing and some degree of ion formation to them. Bonds are classified as covalent if the amount of electron transfer is insufficient for the material to display the classic properties of ionic compounds. If the sharing is unequal enough - polar covalent. The larger the difference in electronegativity, the more polar the bond. Negative end toward more electronegative atom.

Dipole

Movements of

Covalent

Nonmetal and Nonmetal; Electrons shared

Lewis Bonding Theory

One of the simplest bonding theories is called Lewis theory. Lewis theory emphasizes valence electrons to explain bonding. Using Lewis theory, we can draw models, called Lewis structures. Also known as electron dot structures Lewis structures allows us to predict many properties of molecules. Such as molecular stability, shape, size, and polarity

Predictions of Molecular Formulas by Lewis Theory

Oxygen is more stable when it is singly bonded to two other atoms.

Intermediate (0.4 - 2.0)

Polar Covalent; ex. HCL

Size of Isoelectronic ions

Positive ions have more protons so they are smaller.

Example: Using Lewis theory to predict chemical formulas of ionic compounds

Predict the formula of the compound that forms between calcium and chlorine. 1. Draw the Lewis dot symbols of the elements. 2. Transfer all the valence electrons from the metal to the nonmetal, adding more of each atom as you go, until all electrons are lost from the metal atoms and all nonmetal atoms have eight electrons.

Born-Haber Cycle

Remember Hess's law states that the enthalpy change between two states is the same as the sum of enthalpies in a multistep process that goes between the same two states.

Identify the trend in the enthalpy values. Explain the reason for this trend.

Smaller ions will have a greater attraction for each other because of their higher charge density. They will have larger Lattice Enthalpies and larger melting points because of the extra energy which must be put in to separate the oppositely charged ions.

Electronegativity

The ability of an atom to attract bonding electrons to itself is called electronegativity. Increases across period (left to right) and decreases down group (top to bottom) Fluorine is the most electronegative element. Francium is the least electronegative element. Noble gas atoms are not assigned values. Opposite of atomic size trend.

Ionic Bonding and the Crystal Lattice

The extra energy that is released comes from the formation of a structure in which every cation is surrounded by anions, and vice versa. This structure is called a crystal lattice. The crystal lattice is held together by the electrostatic attraction of the cations for all the surrounding anions. The crystal lattice maximizes the attractions between cations and anions, leading to the most stable arrangement.

Energetics of Ionic Bond Formation

The ionization energy of the metal is endothermic. The electron affinity of the nonmetal is exothermic. Generally, the ionization energy of the metal is larger than the electron affinity of the nonmetal; therefore, the formation of the ionic compound should be endothermic. But the heat of formation of most ionic compounds is exothermic and generally large. Why?

Metallic Bonds

The relatively low ionization energy of metals allows them to lose electrons easily. The simplest theory of metallic bonding involves the metal atoms releasing their valence electrons to be shared as a pool by all the atoms/ions in the metal. An organization of metal cation islands in a sea of electrons Electrons delocalized throughout the metal structure Bonding results from attraction of cation for the delocalized electrons.

Ionic Bonding Model versus Reality

To conduct electricity, a material must have charged particles that are able to flow through the material. Lewis theory implies that, in the ionic solid, the ions are locked in position and cannot move around. Lewis theory predicts that ionic solids should not conduct electricity. Ionic solids do not conduct electricity.

Ionic Bonding

Transfer of electrons between two atoms. Each atom achieves a full outer level of electrons. Most atoms in the 2nd and 3rd period, has 8 electrons, known as the octet rule. Metal atom loses electrons it becomes a cation. Metals have low ionization energy, - relatively easy to remove electrons Nonmetal atom gains electrons it becomes an anion. Nonmetals have high electron affinities, - advantageous to add electrons

Covalent Bonding

Two electrons are shared between two atoms. Each atom achieves a full outer level of electrons. The pair of electrons used are called the shared or bonding pair. - this pair of electrons counts for both atoms in completing the octet. The electron pairs that are not involved in bonding belong only to the atom with which they are associated. These are called lone pairs. BOND ORDER - When only one pair of electrons are shared between two atoms, it's called a single bond.- If two pairs of electrons are shared covalently between two atoms, it's called a double bond; three pairs, triple bond.

Valence Electrons and Bonding

Valence electrons are held most loosely. Chemical bonding involves the transfer or sharing of electrons between two or more atoms. Because of the two previously listed facts, valence electrons are most important in bonding. Lewis theory focuses on the behavior of the valence electrons.

Single Covalent Bonds

When two atoms share one pair of electrons, it is called a single covalent bond. Two electrons One atom may use more than one single bond to fulfill its octet. To different atoms

Triple Covalent Bond

When two atoms share three pairs of electrons the result is called a triple covalent bond. Six electrons

Double Covalent Bond

When two atoms share two pairs of electrons the result is called a double covalent bond. Four electrons

Several Molecules

in the Gas Phase