Chapter 15: Ionic Bonding and Ionic Compounds
Unit 6 ---
Ionic and Covalent Bonds
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• Electron Configuration in Ionic Bonding • Ionic Bonds • Bonding in Metals
Chapter 15: Ionic Bonding and Ionic Compounds
Chapter 15: Ionic Bonding and Ionic Compounds
-- Electron Configuration in Ionic Bonding -Valence Electrons
-- Electron Configuration in Ionic Bonding -Electron Dot Structures
Electrons in the highest occupied energy level of an element’s atoms – Examples • Mg: 1s22s22p63s2 2 valence e- in level 3 • Br: 1s22s22p63s23p64s23d104p5 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
Oxygen (6 valence e-)
Nitrogen (5 valence e-)
Sodium (1 valence e)
Calcium (2 valence e-)
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Chapter 15: Ionic Bonding and Ionic Compounds
Chapter 15: Ionic Bonding and Ionic Compounds
-- Electron Configuration in Ionic Bonding -Electron Configurations for Ions
-- Electron Configuration in Ionic Bonding -Electron Configurations for Cations
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
1s22s22p63s1 1s22s22p6 + e8 valence electrons in the highest energy level
Chapter 15: Ionic Bonding and Ionic Compounds
Chapter 15: Ionic Bonding and Ionic Compounds
-- Electron Configuration in Ionic Bonding -Electron Configurations for Cations
-- Electron Configuration in Ionic Bonding -Electron Configurations for Cations
1s22s22p63s2 1s22s22p6 + 2e-
1s22s22p63s23p1 1s22s22p6 + 3e-
8 valence electrons in the highest energy level
8 valence electrons in the highest energy level
Chapter 15: Ionic Bonding and Ionic Compounds -- Electron Configuration in Ionic Bonding -Electron Configurations for Cations (Transition Metals) • • •
Chapter 15: Ionic Bonding and Ionic Compounds -- Electron Configuration in Ionic Bonding -Electron Configurations for Anions
Ideally, transition metals would have to lose their d orbital electrons to achieve a noble gas configuration. Example: Cobalt (1s22s22p63s23p64s23d7) 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.
1s22s22p63s23p64s23d7 1s22s22p63s23p63d7 15 electrons in the outer energy level Still a pseudo-noble gas configuration because of the s and p orbitals being filled
1s22s22p3 + 3e- 1s22s22p6 8 valence electrons in the highest energy level
Chapter 15: Ionic Bonding and Ionic Compounds
Chapter 15: Ionic Bonding and Ionic Compounds
-- Electron Configuration in Ionic Bonding -Electron Configurations for Anions
-- Electron Configuration in Ionic Bonding -Electron Configurations for Anions
1s22s22p4 + 2e- 1s22s22p6
1s22s22p63s23p64s23d104p5 + 1e- 1s22s22p63s23p64s23d104p6
8 valence electrons in the highest energy level 8 valence electrons in the highest energy level
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Chapter 15: Ionic Bonding and Ionic Compounds
Chapter 15: Ionic Bonding and Ionic Compounds
-- Ionic Bonds --
-- Ionic Bonds -Formation of Ionic Bonds
Ionic Bonds Forces of attraction that bind oppositely charged ions Examples – Sodium chloride Na+ attracted to a Cl– Aluminum bromide Al3+ attracted to 3 BrElectromagnetic 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
1s22s22p63s1
1s22s22p63s23p5
1s22s22p6 1s22s22p63s23p6
Transfer of the 3s1 electron of the Na to the p orbital of the Cl Both ions now have octets
Chapter 15: Ionic Bonding and Ionic Compounds
Chapter 15: Ionic Bonding and Ionic Compounds
-- Ionic Bonds -Formation of Ionic Bonds
-- Ionic Bonds --
Aragonite (CaCO3)
1s22s22p63s2
1s22s22p4
1s22s22p6
Barite (BaSO4)
Calcite (CaCO3)
Crystalline solids result from structured arrangements of ionic bonds.
1s22s22p6
Transfer of the 3s2 electrons of the Mg to the p orbital of the O
Both ions now have octets Hematite (Fe2O3)
Pyrite (FeS2)
Chapter 15: Ionic Bonding and Ionic Compounds -- Ionic Bonds -Properties of Ionic Compounds • • • •
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
Chapter 15: Ionic Bonding and Ionic Compounds
Chapter 15: Ionic Bonding and Ionic Compounds
-- Ionic Bonds -Properties of Ionic Compounds
-- Ionic Bonds -Properties of Ionic Compounds
Sodium chloride (NaCl)
Cesium chloride (CsCl)
Face-centered cubic structure
Simple cubic structure
Coordination number of 6 (6 Cl- ions around each Na+ ion)
Coordination number of 8 (8 Cl- ions around each Cs+ ion)
Chapter 15: Ionic Bonding and Ionic Compounds -- Bonding in Metals -Metallic Bonds and Metallic Properties •
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Chapter 16: Covalent Bonding
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
Chapter 16: Covalent Bonding -- The Nature of Covalent Bonding -Single Covalent Bonds
Chapter 16: Covalent Bonding
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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
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Structural formula is a chemical formula showing the arrangment of atoms in a molecule
The Nature of Covalent Bonds
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- The Nature of Covalent Bonding --
-- The Nature of Covalent Bonding -Array of sodium ions and chloride ions:
Collection of water molecules:
Formula unit of sodium chloride:
Molecule of water: Na+
Chemical formula: NaCl
Cl-
H Chemical formula: H2O
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- The Nature of Covalent Bonding --
-- The Nature of Covalent Bonding --
Single Covalent Bonds
Single Covalent Bonds - Halogens
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Covalent bonds result from combinations of nonmetals (I.e., group 4A, 5A, 6A, and 7A elements)
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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
O H
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- The Nature of Covalent Bonding --
-- The Nature of Covalent Bonding --
Single Covalent Bonds – Larger Molecules
Single Covalent Bonds – Larger Molecules
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- The Nature of Covalent Bonding --
-- The Nature of Covalent Bonding --
Single Covalent Bonds – Larger Molecules •
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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
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- The Nature of Covalent Bonding --
-- The Nature of Covalent Bonding --
Double Covalent Bonds
Double Covalent Bonds -- Exceptions •
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- The Nature of Covalent Bonding --
-- The Nature of Covalent Bonding --
Triple Covalent Bonds • •
Oxygen gas (O2) – Expectation: formation of a double-bond to achieve octets – Evidence: formation of a single-bond with two electrons in the gas being unpaired
Bonds that involve three shared pairs of electrons Used to attain stable noble-gas configurations
Coordinate Covalent Bonds • •
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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 (NH4+) – Sulfur dioxide (SO2)
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- The Nature of Covalent Bonding --
-- The Nature of Covalent Bonding --
Coordinate Covalent Bonds – Carbon Monoxide (CO)
Coordinate Covalent Bonds – Ammonium Ion (NH4+)
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.
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Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- The Nature of Covalent Bonding --
-- The Nature of Covalent Bonding --
Bond Dissociation Energies
Resonance
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
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 (H2).
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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
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Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- The Nature of Covalent Bonding --
-- The Nature of Covalent Bonding --
Exceptions to the Octet Rule
Exceptions to the Octet Rule – Nitrogen Dioxide (NO2)
Impossibilities occur where using the octet rule does not work. Examples: – Nitrogen dioxide (NO2) – Oxygen gas (O2) – Phosphorus pentachloride (PCl5) – Sulfur hexafluoride (SF6)
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- The Nature of Covalent Bonding --
-- The Nature of Covalent Bonding --
Exceptions to the Octet Rule – Phosphorus Pentachloride (PCl5)
Exceptions to the Octet Rule – Sulfur Hexafluoride (SF6)
Chapter 16: Covalent Bonding -- The Nature of Covalent Bonding -Exceptions to the Octet Rule •
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Chapter 16: Covalent Bonding
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
Bonding Theories
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- The Nature of Covalent Bonding --
-- Bonding Theories --
Molecular Orbitals •
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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 • • • • •
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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
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- Bonding Theories --
-- Bonding Theories --
Molecular Geometries • Species type: AX3 • Geometry: Trigonal planar • Predicted bond angle(s): 120
Molecular Geometries • Species type: AX4 • Geometry: Tetrahedral • Predicted bond angle(s): 109.5
Example of geometry:
Example of geometry:
CO32-:
CH4:
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- Bonding Theories --
-- Bonding Theories --
Molecular Geometries • Species type: AX5 • Geometry: Trigonal bipyramidal • Predicted bond angle(s): 90, 120, 180
Molecular Geometries • Species type: AX6 • Geometry: Octahedral • Predicted bond angle(s): 90, 180
Example of geometry:
Example of geometry:
PCl5:
SF6:
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- Bonding Theories --
-- Bonding Theories --
Molecular Geometries • Species type: AX2E2 (E: lone electron pair around the central atom) • Geometry: Bent • Predicted bond angle(s): 104.5
Molecular Geometries • Species type: AX2E3 (E: lone electron pair around the central atom) • Geometry: Linear • Predicted bond angle(s): 180 Example of geometry:
H2O:
XeF2:
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- Bonding Theories --
-- Bonding Theories --
Molecular Geometries • Species type: AX3E2 (E: lone electron pair around the central atom) • Geometry: T-shaped • Predicted bond angle(s): 90, 180
Molecular Geometries • Species type: AX4E (E: lone electron pair around the central atom) • Geometry: See-saw • Predicted bond angle(s): 90, 120, 180
Example of
Example of geometry:
geometry:
F3:
SeCl4:
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- Bonding Theories --
-- Bonding Theories --
Molecular Geometries • Species type: AX4E2 (E: lone electron pair around the central atom) • Geometry: Square planar • Predicted bond angle(s): 90, 180
Molecular Geometries • Species type: AX5E (E: lone electron pair around the central atom) • Geometry: Square pyramidal • Predicted bond angle(s): 90, 180
Example of geometry:
Example of geometry:
XeF4:
ClF5:
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- Bonding Theories --
-- Bonding Theories --
Hybrid Orbitals •
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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 – sp2: One s orbital is mixed with two p orbitals – sp3: 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
sp3 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.
Chapter 16: Covalent Bonding -- Bonding Theories --
Chapter 16: Covalent Bonding
Hybrid Orbitals and Their Geometries Number of electron pairs
Atomic Orbitals
Hybrid Orbitals
Geometry
Examples
2
s, one p
sp
Linear
BeF2, CO2
3
s, two p
sp2
Trigonal planar
BF3, CO32-, SO3
4
s, three p
sp3
Tetrahedron
CH4, NH3, H2O
Polar Bonds and Molecules
Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- Polar Bonds and Molecules --
-- Polar Bonds and Molecules --
Bond Polarity •
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“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: H2, O2, N2 – Polar covalent • When atoms pull the bond unequally • Happens with two atoms of different electronegativities • Example: HCl, HF, NH
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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
Moderately covalent bond
H – Cl (0.9 difference)
1.0 – 2.0 difference
Very polar covalent bond
H – F (1.9 difference)
2.0 + difference
Ionic bond
Na+Cl- (2.1 difference)
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Chapter 16: Covalent Bonding
Chapter 16: Covalent Bonding
-- Polar Bonds and Molecules --
-- 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.
Chapter 16: Covalent Bonding
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Attractions Between Molecules 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
Characteristics of Ionic and Covalent Compounds
-- 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
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 (> 300C)
Low (< 300C)
Solubility in water
Usually high
High to low
Electrical conductivity of aqueous solution
Good conductor
Poor conductor or doesn’t conduct at all
