Periodic Trends There are various trends on the periodic table that need to be understood to explain chemical bonding. These include: Atomic/Ionic Radius Ionization Energy Electronegativity Electron Affinity Effective Nuclear Charge Homework: Lewis Theory Do: Periodic Trends Handout Now, where did I leave my keys From 1916-1919, Gilbert N. Lewis made several important proposals on bonding which lead to the development of Lewis Bonding Theory. Elements of his theory: Valence electrons play a fundamental role in chemical bonding. Sometimes bonding involves the transfer of one or more electrons from one atom to another. This leads to the ion formation and IONIC BONDS. Sometimes bonding involves sharing electrons between atoms, this leads to COVALENT BONDS. LEWIS SYMBOLS: A common chemical symbol surrounded by up to 8 dots. The symbol represents the nucleus and the electrons of the filled inner shell orbitals. The dots represent the valence electrons. For Example: Electrons are transferred or shared such that each atom gains a more stable electron configuration. Usually this is that of a noble gas (having 8 outer shell electrons). This arrangement is called an OCTET. 1
Hold the phone Types of Bonding This only works well for the representative elements. Transition metals, actinides and lanthanides have incompletely filled inner shells - we can't write simple Lewis structures for them. Ionic (metal/non-metal) Covalent (non-metal/non-metal) Intermolecular Bonding (between molecules): Hydrogen Bonding London Dispersion Forces Ionic Bonds Ionic bonds are forces that hold ionic compounds together. Forming the ionic bond: Step 1: A cation forms by the LOSS of 1 or more e-. Representative elements become cations, isoelectronic with the nearest Noble gas. For others (transition metals), not necessarily. Step 2: A anion forms by GAINING sufficient e- to become isoelectronic with the nearest Noble gas. The Lewis symbol will show an OCTET (8) of electrons for the anion. Step 3: The oppositely charged ions come together to form an ionic compound. The electrostatic attraction between the ions forms the bond. In the solid state, each anion surrounds itself with cations, and each cation with anions, forming an ionic crystal. Step 4: A formula unit of an ionic compound is the smallest collection of ions that would be electrically NEUTRAL. The formula unit is automatically obtained when the Lewis structure of the compound is written. The ionic crystal then consists of each constituent ion bound together in the crystal, not of individual molecules. ATOMS COMPOUNDS, properties change Their reactivity decreases considerably. They become neutral overall. There are no unique molecules in many ionic solids. Ionic compounds become electrically conductive when melted or dissolved. 2
IONIC BONDS are STRONG so that compounds held together by ionic bonds have HIGH MELTING TEMPERATURES. Covalent Bonds Covalent bonds arise from the sharing of electrons between atoms (generally of groups IVA, VA, VIA, and VIIA). Each electron in a shared pair is attracted to both nuclei involved in the bond. The valence electrons involved in the bond are called the BONDING ELECTRONS or the BOND PAIR. Those not involved in the bond are called the NONBONDING ELECTRONS or the LONE PAIRS. Lone Pair Bond Pair COVALENT BONDS are VERY STRONG. OCTET RULE : An atom other than hydrogen tends to form bonds until it is surrounded by eight valence e-. The pairs repel each other and thus tend to stay as far away as possible. Multiple Bonds Polar Covalent Bonds Covalent compounds can form multiple bonds. Depending upon how many pairs of electrons are shared, different types of bonds are made. Single Bond: Share ONE pair of electrons. Double Bond: Share TWO pairs of electrons. Triple Bond: Share THREE pairs of electrons. Imagine we have one atom with a somewhat higher electronegativity than the other in a covalent bond: This will cause the electrons to be shared unevenly, such that the shared electrons will spend more time (on average) closer to the atom that has the higher ELECTRONEGATIVITY. This is done to fill their valence shell (Octet Rule). 3
The greater the difference in electronegativity in the bonding atoms, the greater the polarity of the bond. Atoms with widely different electronegativity values (ΔE 2.0) tend to form IONIC BONDS. True NON-POLAR COVALENT BONDS form only when diatomic molecules are formed with two identical atoms (ΔE 0.4) Everything else will form a POLAR COVALENT BOND (ΔE 0.5 1.9). Homework: Do: W.S. 7-1 Hydrogen Bonding The force of attraction, shown here as a dotted line, is called a HYDROGEN BOND: Polar molecules, such as water molecules, have a weak, partial negative charge at one region of the molecule (the oxygen atom in water) and a partial positive charge elsewhere (the hydrogen atoms in water). When water molecules are close together, their positive and negative regions are attracted to the oppositelycharged regions of nearby molecules. 4
Each water molecule is hydrogen bonded to four others: London Dispersion Forces The London dispersion force is the weakest intermolecular force. It is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles. HYDROGEN BONDS are WEAK. Because of the constant motion of the electrons, an atom or molecule can develop a temporary (instantaneous) dipole when its electrons are distributed unsymmetrically about the nucleus. Dispersion forces are present between any two molecules (even polar molecules) when they are almost touching. LONDON DISPERSION FORCES are the WEAKEST. Writing Lewis Structures Simple Ionic Compounds Easy to do The overall charge on the compound must equal zero, that is, the number of electrons lost by one atom MUST EQUAL the number of electrons gained by the other atom. The Lewis Structure of each ion is used to construct the Lewis Structure for the ionic compound. For Example: KBr Li 2 S K 3 P Ionic compounds end up having ZERO valence electrons. 5
THE LEWIS STRUCTURES OF COVALENT COMPOUNDS THAT OBEY THE OCTET RULE Lewis structures show how the VALENCE electrons are distributed in a molecule. Covalent compounds share electrons to fill their valence shells. Therefore, the Lewis structures for these compounds are drawn a little differently. Rules of the Road! 1. Find the total number of valence e-. Go by the column or group that it s in. 2. After connecting your central atom to the terminal atoms with single bonds, begin adding the remaining valence e- as lone pairs. First around the terminal atoms. Then around the central atom (if you have any left over). 3. Satisfy the octet rule for your central atom by either: Replacing a lone pair on your terminal atoms with a bond (to make a double bond). OR: By replacing two lone pairs with two bonds (to make a triple bond). Remember: You can t just add double bonds without first removing a lone pair. Not only are you adding more electrons than you started with, but you re probably breaking the octet rule for the terminal atoms. NH + 4 O 2 NO - 2 C 2 H 4 CO THE LEWIS STRUCTURES OF COVALENT COMPOUNDS THAT VIOLATE THE OCTET RULE Incomplete Octets: In addition to H, Be, B, and Al are exceptions to the octet rule. Since they have very low electronegativities, they can only accept one electron for every one they donate. For Example: Hypervalent Compounds: Elements in the third and fourth periods MAY attain more than an octet of valence e- when they form covalent compounds (the electrons are placed in low lying d-orbitals). Other than the fact that the central atom will end up with more than eight valence electrons, the same rules apply. Examples include Xe, P, S, and I. BF 3 6
For Example: PCl 5 SF 4 Formal Charge When we draw Lewis structures, we may end up with several possible structures for the molecule. The question then becomes XeF 4 Calculating the FORMAL CHARGE allows us to figure this out: 1 F.C. = (Valence e ) - 2 (number of e in covalent bonds) + (number of e in lone pairs) Note: The sum of the formal charges equals the overall charge on the ion or molecule. The guidelines : We generally choose the Lewis structure in which the atoms bear formal charges closest to ZERO. We generally choose the Lewis structure in which any negative formal charge resides on the more electronegative atom. For Example: Draw the three possible Lewis structures for the thiocyanate ion, NCS -. Determine the formal charges of the atoms in each structure. Which Lewis structure is the preferred one? Resonance For Example: When we draw Lewis structures in which we must make a choice as to what gets a double bond, the structure is actually a blend of two or three structures. Draw three resonance structures for the polyatomic ion CO 3 2-. We say that the structure RESONATES (contains contributions from each of the resonance structures). Resonance occurs simply because the electron-dot model is too limited to show how electrons are being shared between the atoms. 7
Homework: VSEPR Do: Lewis Structures W.S. Study for your quiz! Lewis Diagrams The Valence Shell Electron Pair Repulsion model is not so much a model of chemical bonding, as a scheme for explaining the shapes of molecules. It is based on the quantum mechanical view that bonds represent electron clouds that repel each other and thus try to stay as far apart as possible. This minimizes the energy of repulsion, represents the lowest energy configuration of the molecule, and gives it a distinctive shape. Electron Group Geometries An ELECTRON GROUP is any collection of valence electrons, localized in a region around a central atom, that repels other groups of valence electrons. Can include: A single unpaired electron. A lone pair of electrons. One bonding pair of electrons in a single covalent bond. Two bonding pairs of electrons in a double covalent bond. Three bonding pairs of electrons in a triple covalent bond. The overall shape of the molecule is determined by its BOND ANGLES (the angles made by the lines joining the nuclei of the atoms in the molecule). The central idea is that bonding and non-bonding pairs around a given atom will be positioned as far apart as possible. Mutual repulsions among electron groups lead to their electron group geometry. It s all about repulsion, baby! The electron group geometries are: Two e - groups: Three e - groups: Four e - groups: Five e - groups: Six e - groups: Seven e - groups: Let s reason our way through a few CO 2 BCl 3 CH 4 PCl 5 SF 6 IF 7 8
It does get a little more complicated How will we do this? Step 1: Draw a Lewis structure of the molecule. Step 2: Determine the number of electron groups around the central atom, and identify each as either a bonding group or a lone pair. Step 3: Identify the electron group geometry. Step 4: Identify the molecular geometry. Ok here goes nothing!! Homework: BeCl 2 Do: VSEPR W.S. BF 3 SO 2 NH 3 XeF 2 Covalent Bonding and Orbital Overlap VSEPR provides a simplistic model for predicting molecular shape, but does not explain why bonds exist. Lewis Theory + Quantum Mechanics = Atoms sharing electrons concentrates electron density between the two nuclei. These orbitals are then said to share a region of space, or to OVERLAP. Take H 2 for example: Electron overlap creates the covalent bond. Another involves HCl: Chlorine has an electron configuration of [Ne]3s 2 3p 5. All its orbitals are full except for one 3p orbital with a single electron. This electron pairs up with the single electron of H (1s 1 ) to form the covalent bond. Yet another Cl 2. 9
There is always an optimum distance between the two bonded nuclei in any covalent bond. Overlapping of orbitals helps us to understand the formation of covalent bonds, but its not easy to extend these ideas to polyatomic molecules. In these cases, we must explain both the formation of electron-pair bonds and the observed geometries of the molecules. To explain geometries, we assume that the atomic orbitals on an atom mix to form new orbitals called HYBRID ORBITALS. During the process of HYBRIDIZATION, the total number of atomic orbitals on an atom remains constant, so, the number of hybrid orbitals on an atom equals the number of atomic orbitals mixed. The Hybrid Orbital Model sp Hybrid Orbitals Developed by Linus Pauling in 1931. There are several types of hybridization to know:,,, and Each one of these types of hybridization are connected to the five basic electron geometries: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral. Linus Pauling (1901-1994) Consider BeF 2 : VSEPR predicts that the molecule is linear with two identical Be-F bonds. Each F has an electron configuration of 1s 2 2s 2 2p 5. This 2p e - can be paired with an unpaired electron Unpaired e - in from the Be atom to form a polar covalent bond. a 2p orbital. The orbital diagram for a ground state Be atom is: Notice that there are no unpaired electrons available, therefore, it is incapable of forming bonds with Fluorine atoms. With this e - configuration, the Be atom now has two unpaired e - it can form two polar covalent bonds with the F atoms. But these concepts still do not explain the structure of BeF 2 since: BUT it does, by promoting one of the 2s electrons to a 2p orbital: 10
This is accomplished by mixing the 2s and 2p orbitals to generate two new, equal orbitals as follows: The e - in the sp hybrid orbital can form two e - bonds with the two Fluorine atoms. These two orbitals are identical in shape, but point in opposite directions. Since these orbitals are equivalent, but in opposite directions, BeF 2 has two identical bonds and a linear geometry. This is what we call a sp HYBRID ORBITAL. A linear arrangement of electron domains implies sp hybridization. sp 2 and sp 3 Hybrid Orbitals This results in three equivalent sp 2 hybrid orbitals: Mixing a certain number of atomic orbitals always results in the same number of hybrid orbitals. For example: Mixing one 2s and one 2p orbital results in two equivalent sp hybrid orbitals pointing in opposite directions. There are other possible combinations. Using BF 3 as an example, a 2s e - on the B atom can be promoted to a vacant 2p orbital. These three sp 2 hybrid orbitals lie in the same plane, 120 o apart from one another. This leads to the TRIGONAL PLANAR geometry of BF 3. Note that an unfilled 2p orbital remains unhybridized... more to come later (hint: Carbon)! In sp 3 hybridization, such as in the C in CH 4, an s orbital can mix with all three p orbitals in the same subshell. The result is four equivalent sp 3 hybrid orbitals. Each one of these hybrid orbitals has a large lobe that points towards a vertex of a tetrahedron. These hybrid orbitals can form two electron bonds by overlapping with the atomic orbitals of H to form four equivalent sp 3 hybrid orbitals on the C. 11
Hybridization Involving d Orbitals Atoms in the third period and beyond can also use d orbitals to form hybridized orbitals. For example: Mixing one s orbital, three p orbitals, and one d orbital leads to FIVE sp 3 d hybrid orbitals directed towards the vertices of a trigonal bipyramid (eg: PF 5 ). Mixing one s orbital, three p orbitals, and two d orbitals leads to SIX sp 3 d 2 hybrid orbitals directed towards the vertices of an octahedron (eg: SF 6 ). To Summarize: Hybridization offers a convenient model for using valence bond theory to describe covalent bonds whose geometries are predicted by the VSEPR. To predict hybrid orbitals used by an atom in bonding: 1. Draw the Lewis structure for the molecule or ion. 2. Determine the electron-domain geometry using the VSEPR model. 3. Specify the hybrid orbitals needed to accommodate the electron pairs based on their geometric arrangement. An example: Indicate the hybridization of orbitals employed by the central atom in: NH 2 - SF 4 Multiple Bonds In all the hybridization examples so far, e - density is concentrated symmetrically about the line connecting the nuclei (the intermolecular axis). The line joining the two nuclei passes through the middle of the overlap region. These bonds are called SIGMA (σ) BONDS. 12
Multiple bonds involve a second type of bond resulting from perpendicular overlap between two p orbitals. This produces a PI (π)bond. In pi bonds, the overlap regions are above and below the intermolecular axis. The sideways orientation of p orbitals in a π bond makes for weaker overlap. Single, double, and triple bonds are all different: In almost all cases, single bonds are σ bonds. A double bond consists of one σ bond and one π bond. A triple bond consists of one σ bond and two π bonds. Consider Ethylene (C 2 H 4 ) which possesses a C=C double bond: As a result, π bonds are generally weaker than σ bonds. A bond angle of 120 o suggests that each C atom uses sp 2 hybrid orbitals to form σ bonds with the other Carbon and two Hydrogens. Since Carbon has four valence e -, after sp2 hybridization, one e - in each of the Carbon atoms remains in the unhybridized 2p orbital: I m so lonely This unhybridized 2p orbital is directed perpendicular to the plane that contains the three sp 2 hybrid orbitals. The same principle can be used to explain triple bonds. Consider Acetylene (C 2 H 2 ): Acetylene is also linear and uses sp hybrid orbitals to form σ bonds with the other Carbon and one Hydrogen. The two remaining unhybridized 2p orbitals overlap to form a pair of π bonds. 13
Homework: What is on the exam? Do: Bonding W.S. Study for your quiz! VSEPR and Hybridization Periodic Trends Atomic and Ionic Radii Ionization Energy Electron Affinity and Electronegativity Lewis Theory Elements of the Theory Ionic Bonds Covalent Bonds Multiple Bonds Polarity Intermolecular Bonding Hydrogen Bonding London Dispersion Forces Drawing Lewis structures Simple Ionic Compounds Structures that Obey the Octet Rule Structures that Violate the Octet Rule Formal Charge Resonance VSEPR Hybridization 14