Unit Six --- Ionic and Covalent Bonds Electron Configuration in Ionic Bonding Ionic Bonds Bonding in Metals Valence Electrons Electrons in the highest occupied energy level of an element s atoms Examples Mg: 1s 2 2s 2 2p 6 3s 2 2 valence e - in level 3 Br: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 7 valence e - in level 4 Identification of group number gives valence electrons for the representative elements Group 1A elements (i.e., hydrogen, lithium, etc.) have 1 valence electron Group 6A elements (i.e., oxygen, sulfur, etc.) have 6 valence electrons Usually the only electrons used in chemical bonds Only electrons shown in the electron dot structures Electron Dot Structures Oxygen Nitrogen Sodium Calcium (6 valence e - ) (5 valence e - ) (1 valence e) (2 valence e - )
Electron Configurations for Ions Ions strive to become like noble gases Octet rule Atoms tend to achieve the electron configuration of a noble gas Ions strive to have 8 valence electrons Metals will lose electrons to go back to a noble gas configuration in the greatest energy level of their electron configurations Nonmetals will gain electrons to go to a noble gas configuration in the greatest energy level of their electron configurations Transition metals will go to a pseudo-noble gas configuration in their electron configurations Electron Configurations for Cations 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 + e - 8 valence electrons in the highest energy level Electron Configurations for Cations Electron Configurations for Cations 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 + 2e - 8 valence electrons in the highest energy level 1s 2 2s 2 2p 6 3s 2 3p 1 1s 2 2s 2 2p 6 + 3e - 8 valence electrons in the highest energy level
Electron Configurations for Cations (Transition Metals) Ideally, transition metals would have to lose their d orbital electrons to achieve a noble gas configuration. Example: Cobalt (1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 ) would have to lose nine electrons to get back to a noble gas configuration. Transition metals can have pseudo-noble gas electron configurations by typically losing the s orbital electrons. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 1s 2 2s 2 2p 6 3s 2 3p 6 3d 7 15 electrons in the outer energy level Still a pseudo-noble gas configuration because of the s and p orbitals being filled Electron Configurations for Anions 1s 2 2s 2 2p 3 + 3e - 1s 2 2s 2 2p 6 8 valence electrons in the highest energy level Electron Configurations for Anions Electron Configurations for Anions 1s 2 2s 2 2p 4 + 2e - 1s 2 2s 2 2p 6 8 valence electrons in the highest energy level 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 + 1e - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 8 valence electrons in the highest energy level
-- Ionic Bonds -- Ionic Bonds Forces of attraction that bind oppositely charged ions Examples Sodium chloride Na + attracted to a Cl - Aluminum bromide Al 3+ attracted to 3 Br - Electromagnetic attraction Transfer (NOT sharing) of electron(s) from one neutral atom to another neutral atom to create ions Each ion will have an octet in the outer shell Exceptions: transition metals -- Ionic Bonds -- Formation of Ionic Bonds 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 6 Transfer of the 3s 1 electron of the Na to the p orbital of the Cl Both ions now have octets -- Ionic Bonds -- Formation of Ionic Bonds -- Ionic Bonds -- Aragonite (CaCO 3 ) Barite (BaSO 4 ) Calcite (CaCO 3 ) 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 4 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 Transfer of the 3s 2 electrons of the Mg to the p orbital of the O Both ions now have octets Hematite (Fe 2 O 3 ) Pyrite (FeS 2 ) Crystalline solids result from structured arrangements of ionic bonds.
-- Ionic Bonds -- Properties of Ionic Most ionic compounds are crystalline solids. Ions are arranged in repeating, three-dimensional patterns. Fourteen kinds of arrangements Coordination number Number of ions of opposite charge that surround the ion in a crystal -- Ionic Bonds -- Properties of Ionic Sodium chloride (NaCl) -- Ionic Bonds -- Properties of Ionic Cesium chloride (CsCl) Face-centered cubic structure Coordination number of 6 (6 Cl - ions around each Na + ion) Simple cubic structure Coordination number of 8 (8 Cl - ions around each Cs + ion)
-- Bonding in Metals -- Metallic Bonds and Metallic Properties Metallic Bonds Consist of the attraction of free-floating valence electrons for positivelycharged metal ions What hold the metal together Allow for malleability (ability to reshape and bend) Allow for conductivity Alloys Mixtures of two or more elements, at least one of which is a metal Generally made by melting a mixture of the elements and then cooling the mixture Properties are often superior to lone metals Steels The Nature of Covalent Bonds Bonding Theories Polar Bonds and Molecules Unit Six: Covalent Bonding Single Covalent Bonds Two atoms share one pair of electrons Each atom ideally achieves an octet in a covalent bond so that they resemble the electron configuration of a noble gas The Nature of Covalent Bonds Structural formula is a chemical formula showing the arrangment of atoms in a molecule
Array of sodium ions and chloride ions: Collection of water molecules: Formula unit of sodium chloride: Molecule of water: Na + Cl - H O H Chemical formula: NaCl Chemical formula: H 2 O Single Covalent Bonds Covalent bonds result from combinations of nonmetals (I.e., group 4A, 5A, 6A, and 7A elements) Single Covalent Bonds - Halogens Unshared pairs Also known as lone pairs Pairs of valence electrons that are not shared between atoms of a molecule Unshared pairs do not change form in a structural formula
Single Covalent Bonds Larger Molecules Single Covalent Bonds Larger Molecules Single Covalent Bonds Larger Molecules Spreading out the electrons More stability Less energy required Preferred arrangements Double Covalent Bonds Bonds that involve two shared pairs of electrons Used to attain stable noble-gas configurations
Double Covalent Bonds Double Covalent Bonds -- Exceptions Oxygen gas (O 2 ) Expectation: formation of a double-bond to achieve octets Evidence: formation of a single-bond with two electrons in the gas being unpaired Triple Covalent Bonds Bonds that involve three shared pairs of electrons Used to attain stable noble-gas configurations Coordinate Covalent Bonds Covalent bond in which an atom contributes both bonding electrons Structural formulas of coordinate covalent bonds show the bonds as arrows that point from the atom donating the pair of electrons to the atom receiving them Examples Carbon monoxide (CO) Ammonium ion (NH 4+ ) Sulfur dioxide (SO 2 )
Coordinate Covalent Bonds Carbon Monoxide (CO) Coordinate Covalent Bonds Ammonium Ion (NH 4+ ) An octet has been achieved for each molecule, but nitrogen contributes the electrons needed. An octet has been achieved for each molecule, but oxygen contributes the electrons needed. Bond Dissociation Energies Total energy required to sever the bond between two covalently bonded atoms High in carbon compounds, resulting in high stability of carbon compounds Table 16.3, page 448 Resonance Example: H H + 435 kj H + H This means that it would require 435 kj of energy to break the bond between the two atoms in a hydrogen gas molecule (H 2 ). Structures that occur when it is possible to write two or more valid Lewis dot structures that have the same number of electron pairs for a molecule or ion Structures are in constant resonance NOTE: Single bonds are longer than double bonds; double bonds are longer than triple bonds
Exceptions to the Octet Rule Impossibilities occur where using the octet rule does not work. Examples: Nitrogen dioxide (NO 2 ) Oxygen gas (O 2 ) Phosphorus pentachloride (PCl 5 ) Sulfur hexafluoride (SF 6 ) Exceptions to the Octet Rule Nitrogen Dioxide (NO 2 ) Exceptions to the Octet Rule Phosphorus Pentachloride (PCl 5 ) Exceptions to the Octet Rule Sulfur Hexafluoride (SF 6 )
Exceptions to the Octet Rule Cases for exceptions More than 8 valence electrons Less than 8 valence electrons How to draw Typically, the central atom will be the first one listed in the formula. Hydrogens and halogens will typically surround the central atom. Diamagnetic Substance weakly repelled by a magnetic field Paramagnetic Substance strongly attracted to a magnetic field These substances have molecules containing two or more unpaired electrons. Not to be confused with ferromagnetism (as with magnets) Mass if offset in a magnetic field Unit Six: Covalent Bonding Bonding Theories Molecular Orbitals Covalent bonding occurs as a result of an imbalance between the attractions and the repulsions of the nuclei and the electrons of the atoms. If two atoms don t bond, the repulsion between nuclei of two atoms and the atoms electrons is greater than the attractions of the electrons to the opposing nuclei. If two atoms do bond, the attractions of the electrons to the opposing nuclei is greater than the repulsion between nuclei of the two atoms and the atoms electrons. Pi bonds ( ) and sigma bonds ( ) are responsible for covalent bonding. Overlapping of orbitals cause bonds. Sharing of electrons from overlapping Symmetrical bonding VSEPR Theory Valence Shell Electron Pair Repulsion Theory Electron pairs around atoms tend to be as far apart as possible. Similar charges (I.e., negative charges from electrons) tend to repel each other and want to be spaced apart at maximum angles. Used to predict molecular geometries Bond angles Angles between bonds Spacing apart as far as possible Lone pairs of electrons will repel bonded atoms a bit more than expected toward each other around the central atom
Species type: AX 3 Geometry: Trigonal planar Predicted bond angle(s): 120 Species type: AX 4 Geometry: Tetrahedral Predicted bond angle(s): 109.5 Example of geometry: CO 3 2- : Example of geometry: CH 4 : Species type: AX 5 Geometry: Trigonal bipyramidal Predicted bond angle(s): 90, 120, 180 Species type: AX 6 Geometry: Octahedral Predicted bond angle(s): 90, 180 Example of geometry: PCl 5 : Example of geometry: SF 6 :
Species type: AX 2 E 2 (E: lone electron pair around the central atom) Geometry: Bent Predicted bond angle(s): 104.5 H 2 O: Species type: AX 2 E 3 (E: lone electron pair around the central atom) Geometry: Linear Predicted bond angle(s): 180 Example of geometry: XeF 2 : Species type: AX 3 E 2 (E: lone electron pair around the central atom) Geometry: T-shaped Predicted bond angle(s): 90, 180 Example of geometry: F 3 : Species type: AX 4 E (E: lone electron pair around the central atom) Geometry: See-saw Predicted bond angle(s): 90, 120, 180 Example of geometry: SeCl 4 :
Species type: AX 4 E 2 (E: lone electron pair around the central atom) Geometry: Square planar Predicted bond angle(s): 90, 180 Example of geometry: XeF 4 : Species type: AX 5 E (E: lone electron pair around the central atom) Geometry: Square pyramidal Predicted bond angle(s): 90, 180 Example of geometry: ClF 5 : Hybrid Orbitals Hybridization Atomic orbitals mix to form the same total number of equivalent hybrid orbitals Number of hybrid orbitals is equal to the number of atomic orbitals that are mixed Classifications sp: One s orbital is mixed with one p orbital sp 2 : One s orbital is mixed with two p orbitals sp 3 : One s orbital is mixed with three p orbitals Explains why atoms that should not be able to bond covalently can bond Based on the number of electrons pairs Unshared as well as shared electron pairs can be located in hybrid orbitals sp 3 hybridization : hydrogen electrons bonding with carbon electrons Carbon should only be able to bond with two other electron orbitals normally, but it can bond with four when its orbitals are hybridized.
Number of electron pairs Hybrid Orbitals and Their Geometries Atomic Orbitals Hybrid Orbitals Geometry Examples 2 s, one p sp Linear BeF 2, CO 2 Unit Six: Covalent Bonding Polar Bonds and Molecules 3 s, two p sp 2 Trigonal planar 4 s, three p sp 3 Tetrahedron BF 3, CO 3 2-, SO 3 CH 4, NH 3, H 2 O -- Polar Bonds and Molecules -- Bond Polarity The Tug of War The pairs of electrons that are bonds between atoms are pulled between the nuclei of the atoms in a bond. The electronegativities of the atoms determine the winner. Classifications for Bonds Nonpolar covalent When atoms pull the bond equally Happens with two atoms of equal electronegativity, most often using the same atoms Examples: H 2, O 2, N 2 Polar covalent When atoms pull the bond unequally Happens with two atoms of different electronegativities Example: HCl, HF, NH -- Polar Bonds and Molecules -- Bond Polarity Ionic When atoms pull the bond equally in an ionic compound Happens with two atoms of different electronegativities, most often using a metal and a nonmetal Examples: NaCl, KBr In a polar molecule, one end of the molecule is slightly more electronegative than the other atom, resulting in one atom being slightly negative ( -) because of higher electronegativitiy, and the other atom being slightly positive ( +) because of lower electronegativity. is known as a partial charge since it is much less than 1+ or 1- charge.
-- Polar Bonds and Molecules -- Bond Polarity Electronegativities and Bond Types See page 405, Table 14.2 for electronegativities. H: 2.1 Cl: 3.0 Since hydrogen is less, it will have the positive partial charge while chlorine has the negative partial charge. 3.0 2.1 = 0.9 HCl is polar covalent. 0.0 0.4 difference Nonpolar covalent bond H H (0.0 difference) 0.4 1.0 difference 1.0 2.0 difference Moderately covalent bond Very polar covalent bond H Cl (0.9 difference) H F (1.9 difference) -- Polar Bonds and Molecules -- Polar Molecules Dipole Molecule that has two poles Example: HCl from the previous page Polar vs. Nonpolar Water will be polar (charge goes from bottom to top even though the two cancel out sideways) Carbon dioxide will be nonpolar because the charges cancel out in all directions. 2.0 + difference Ionic bond Na + Cl - (2.1 difference) -- Polar Bonds and Molecules -- Intermolecular Attractions van der Waals forces Two types: dispersion forces and dipole interactions Dispersion forces Weakest of all molecular interactions Caused by movement of electrons Strength increases as number of electrons in the molecule increases Examples: Br-Br, F-F, etc. Dipole interactions Occurs when polar molecules are attracted to one another Partial charge ( +) of one polar molecule is attracted to the opposite partial charge ( -) of another molecule -- Polar Bonds and Molecules -- Attractions Between Molecules Hydrogen bonding Hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom Example: water
Characteristics of Ionic and Covalent Characteristic Ionic Compound Covalent Compound Representative unit Formula unit Molecule Bond formation Transfer of electrons Sharing of electrons Types of elements Metals and nonmetals Nonmetals Physical state at room temperature Solid Solid, liquid, gas Melting point High (> 300 C) Low (< 300 C) Solubility in water Usually high High to low Electrical conductivity of aqueous solution Good conductor Poor conductor or doesn t conduct at all