Chapter 6 Chemical Bonding
Section 6.1 Intro to Chemical Bonding
6.1 Objectives Define chemical bond. Explain why most atoms form chemical bonds. Describe ionic and covalent bonding. Explain why most chemical bonding is neither purely ionic nor purely covalent. Classify bonding type according to electronegativity differences.
Imagine getting onto a crowded elevator. As people squeeze into the confined space, they come in contact with each other. Most people will experience a sense of being too close together. When atoms get close enough, their outer electrons repel each other. At the same time, however, each atom s outer electrons are strongly attracted to the nuclei of the surrounding atoms. The degree to which these outer electrons are attracted to other atoms determines the kind of chemical bonding that occurs between the atoms.
6.1 Chemical Bonding Chemical bond Mutual electrical attraction between the nuclei and valence e - of different atoms that bind the atoms together Nature favors chemical bonding because most atoms have lower potential energy when they are bonded to other atoms than they have as independent particles 3 main types of bonds 1. Ionic 2. Covalent 3. Metallic Nanographene
6.1 Ionic Bonds Chemical bonding resulting from the electrical attraction between cations and anions Atoms completely give up electrons to other atoms Between a metal and nonmetal Strongest bond type
6.1 Covalent Bonds Chemical bonding resulting from the sharing of electron pairs between two atoms Typically between nonmetals
6.1 Metallic Bond Chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons Free flowing, mobile valence electrons form a sea around the metal atoms Weakest bond type
6.1 Bond Determination Bonding is rarely purely ionic or purely covalent Degree of ionic or covalent bonding determined by differences in electronegativity between the bonding atoms RECALL: Electronegativity is a measure of an atom s ability to attract bonding electrons Calculate the difference in electronegativity between the two atoms and compare the difference to the table or diagram below Electronegativity Difference Bond Type 1.7-3.3 Ionic 0.3-1.6 Polar covalent < 0.3 Nonpolar covalent
6.1 Bond Determination Polar covalent bond Covalent bond in which the bonded atoms have an unequal attraction for the shared electrons Nonpolar covalent bond Covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge
6.1 Sample Problem A Use electronegativity values listed below to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative? S: 2.5 H: 2.1 Cs: 0.7 Cl: 3.0
6.1 Sample Problem A The electronegativity of sulfur is 2.5. The electronegativities of hydrogen, cesium, and chlorine are 2.1, 0.7, and 3.0, respectively. In each pair, the atom with the larger electronegativity will be the more-negative atom. Bonding between S and: Electroneg diff Bond type More negative atom H 2.5 2.1 = 0.4 Polar-covalent S Cs 2.5 0.7 = 1.8 Ionic S Cl 3.0 2.5 = 0.5 Polar-covalent Cl
Section 6.2 Part I Covalent Bonding and Molecular Compounds
6.2 Objectives Define molecule and molecular formula. Explain the relationships among potential energy, distance between approaching atoms, bond length, and bond energy. State the octet rule. List the six basic steps used in writing Lewis structures. Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both.
6.2 Molecular Compounds Molecule A neutral group of atoms that are held together by covalent bonds Molecular compound A chemical compound whose simplest units are molecules Chemical formula Indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts Molecular formula Chemical formula of a molecule Oxygen molecular formula = Water molecular formula = Sucrose molecular formula = O 2 H 2 O C 12 H 22 O 11
6.2 Covalent Bond Formation Most atoms have lower potential energy when they are bonded to other atoms than they have as are independent particles. The electron of one atom and proton of the other atom attract one another. The two nuclei and two electrons repel each other. These two forces cancel out to form a covalent bond at a length where the potential energy is at a minimum.
6.2 H 2 Covalent Bond Formation a. 2 H atoms are far enough away that there is NO interaction PE=0 b. As the atoms begin to move towards each other, their charged particles begin to interact e - /p + attraction is greater than the p + /p + & e - /e - repulsions, so atoms are drawn together and PE decreases
6.2 H 2 Covalent Bond Formation c. The attractive forces dominate until a distance is reached where the forces are equal PE is at its minimum value A stable H 2 molecule has formed! d. If the nuclei were to move closer, the repulsive forces would become greater PE increases
6.2 Covalent Bond Characteristics Atoms release energy as they form bonds, equal to the decrease in PE so conversely, energy must be added to separate bonded atoms Bond energy The energy required to break a chemical bond and form neutral isolated atoms kj/mol (energy required to break one mole of bonds) Bond length The distance between two atoms at their minimum potential energy The average distance between two bonded atoms
6.2 Covalent Bonds Bond length and energy vary depending on the bonding nuclei What is the relationship between the two?
6.2 Covalent Bond Characteristics Shared electrons form overlapping orbitals The goal is to achieve a noble-gas configuration The bonding of two hydrogen atoms allows each atom to have the stable electron configuration of helium, 1s 2.
6.2 Multiple Covalent Bonds Double covalent bond A covalent bond in which two pairs of electrons are shared between two atoms Often found in molecules containing C, N, and O Triple covalent bond A covalent bond in which three pairs of electrons are shared between two atoms Double and triple bonds are referred to as multiple bonds, or multiple covalent bonds. In general, double bonds have greater bond energies and are shorter than single bonds. Triple bonds are even stronger and shorter than double bonds.
Covalent Network Solids Some covalent compounds are not individual molecules Covalent network solid each atom is joined to all its neighbors in a covalently bonded, 3D network No distinct units, just like ionic compounds Usually contain at least one element from Group 14 Very hard with high melting points SiO 2
Section 6.2 Part II Octet Rule & Lewis Structures
6.2 Objectives Define molecule and molecular formula. Explain the relationships among potential energy, distance between approaching atoms, bond length, and bond energy. State the octet rule. List the six basic steps used in writing Lewis structures. Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both.
6.2 The Octet Rule Remember, in nature, things tend towards their lowest energy states Noble gases can exist in the monoatomic state because they have a minimum energy due to the special stability of their e - configurations Outer s and p orbitals full with a total of 8 electrons (an octet) Chemical compounds tend to form so that each atom, by gaining, losing or sharing electrons, has an octet of electrons in its highest occupied energy level
6.2 Octet Rule Exceptions Exceptions to the octet rule include those for atoms that cannot fit eight electrons, and for those that can fit more than eight electrons, into their outermost orbital. Hydrogen forms bonds in which it is surrounded by only two electrons. Boron has just three valence electrons, so it tends to form bonds in which it is surrounded by six electrons. Main-group elements in Periods 3 and up can form bonds with expanded valence, involving more than eight electrons.
6.2 Electron dot notation To keep track of valence electrons, it is helpful to use electron-dot notation Electron-dot notation an electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element s symbol. The innershell electrons are not shown.
6.2 Sample Problem B a. Write the electron-dot notation for hydrogen. b. Write the electron-dot notation for nitrogen. a. A hydrogen atom has only one occupied energy level, the n = 1 level, which contains a single electron. b. The group notation for nitrogen s family of elements is ns 2 p 3. Nitrogen has five valence electrons.
6.2 Lewis Structures Electron-dot notation can also be used to represent molecules. The pair of dots between two symbols represents the shared electron pair. Must also show unshared electrons Unshared pairs of electrons (aka lone electrons) are not involved in bonding and belong exclusively to one atom
6.2 Lewis Structures The pair of dots representing a shared pair of electrons in a covalent bond is often replaced by a long dash. Structural formula indicates the kind, number, arrangement, and bonds but NOT the unshared pairs of the atoms in a molecule. F-F H-Cl
6.2 Lewis Structure Rules
6.2 Lewis Structure Rules
6.2 Sample Problem C Draw the Lewis structure of iodomethane, CH 3 I. 1. Determine the type and number of atoms in the molecule. The formula shows one carbon atom, one iodine atom, and three hydrogen atoms. 2. Write the electron-dot notation for each type of atom in the molecule. Carbon: Group 14, 4 valence e - Iodine: Group 17, 7 valence e - Hydrogen: 1 valence e -
6.2 Sample Problem C 3. Determine the total number of valence electrons available in the atoms to be combined. CH 3 I Atom # Atoms Valence e- Total valence e- C 1 4 4 I 1 7 7 H 3 1 3 Total valence e - = 14
6.2 Sample Problem C 4. Arrange the atoms to form a skeleton structure for the molecule. If carbon is present, it is the central atom. Otherwise, the least-electronegative atom is central. Hydrogen, is never central. 5. Add unshared pairs of electrons to each nonmetal atom (except H) such that each is surrounded by eight electrons.
6.2 Sample Problem D Draw the Lewis structure for methanal, CH 2 O, which is also known as formaldehyde. 1. Determine the type and number of atoms in the molecule. The formula shows one carbon atom, two hydrogen atoms, and one oxygen atom. 2. Write the electron-dot notation for each type of atom in the molecule. Carbon: Group 14, 4 valence e - Oxygen: Group 16, 6 valence e - Hydrogen: 1 valence e -
6.2 Sample Problem D 3. Determine the total number of valence electrons available in the atoms to be combined. CH 2 O Atom # Atoms Valence e- Total valence e- C 1 4 4 O 1 6 6 H 2 1 2 Total valence e - = 12
6.2 Sample Problem D 4. Arrange the atoms to form a skeleton structure for the molecule. Connect the atoms by electron-pair bonds. 5. Add unshared pairs of electrons to each nonmetal atom (except H) such that each is surrounded by eight electrons.
6.2 Sample Problem D 6. Count the electrons in the Lewis structure to be sure that the number of valence electrons used equals the number available. The structure has 14 electrons. The structure has two valence electrons too many. Subtract one or more lone pairs until the total number of valence electrons is correct. Move one or more lone electron pairs to existing bonds until the outer shells of all atoms are completely filled.
Section 6.3 Ionic Bonding and Ionic Compounds
6.3 Objectives Compare a chemical formula for a molecular compounds with one for an ionic compound. Discuss the arrangements of ions in crystals. Define lattice energy and explain its significance. List and compare the distinctive properties of ionic and molecular compounds. Write the Lewis structure for a polyatomic ion given the identity of the atoms combined and other appropriate information.
6.3 Ionic Bonds Most of the rocks and minerals that make up Earth s crust consist of positive and negative ions held together by ionic bonding. Ionic compound Composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal Most exist as crystalline solids 3D network of positive and negative ions mutually attracted to each other MgO NaCl
6.3 Ionic Bonds Unlike a molecular compound, an ionic compound is not composed of independent, neutral units that can be isolated The chemical formula of an ionic compound represents not molecules, but the simplest ratio of the compound s ions Formula unit (ionic formula) Simplest collection of atoms from which an ionic compound s formula can be established Ratio of ions depends on the ionic charges being combined, must achieve electrical neutrality
6.3 Ionic Bond Formation Group 1 and 2 elements easily LOSE electrons Group 17 elements (and sometimes O) easily GAIN electrons The ions are electrically attracted to one another RECALL: Look at electronegativity differences!
6.3 Ionic Bond Characteristics In an ionic crystal, ions minimize their potential energy by forming a crystal lattice Maximizes attractive forces between oppositely charged ions and minimizes repulsive forces between like ions Lattice energy Energy released when one mole of an ionic crystalline compound is formed from gaseous ions
6.3 Ionic vs Covalent Bonding The force that holds ions together in an ionic compound is a very strong electrostatic attraction. In molecular compounds, the covalent bonds are also strong, BUT, the forces of attraction BETWEEN molecules of a covalent compound are much weaker Difference in the strength of attraction between the basic units of molecular and ionic compounds gives them very different properties
6.3 Ionic vs Covalent Compounds Physical properties of materials depend on intermolecular forces, the strength of attraction between compounds Physical Property Ionic Compound Molecular Compound Hardness All solid at room temperature, hard but brittle Many are gas and liquid at room temp Melting point High > 300 C Low < 300 C Boiling point High, usually solid @ room temp Solubility in H 2 O High, ions separate Low High to low, depends on polarity Conductivity Only when ions can move as in molten or Nonconductive dissolved states Types of elements Metal & nonmetal Nonmetals
6.3 Polyatomic Ions Certain atoms can bond covalently to form a group of atoms with both molecular and ionic characteristics Polyatomic Ion A charged group of covalently bonded atoms Have a charge from either a shortage or excess of electrons
Section 6.4 Metallic Bonding
6.4 Objectives Describe the electron-sea model of metallic bonding, and explain why metals are good electrical conductors. Explain why metal surfaces are shiny. Explain why metals are malleable and ductile but ionic-crystalline compound are not.
6.4 Metallic Bonding In metals, vacant orbitals in the atoms outer energy levels overlap allows the outer electrons of the atoms to roam freely throughout the entire metal The electrons are delocalized; they do not belong to any one atom but move freely about the network of empty atomic orbitals These mobile electrons form a sea of electrons around the metal atoms, which are packed together in a crystal lattice. The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons is called metallic bonding Weakest of the 3 bond types
6.4 Metallic Properties The unique characteristics of metallic bonding gives metals their characteristic properties, listed below. Electrical conductivity in the solid state due to highly mobile valence electrons Thermal conductivity - mobile electrons also transfer heat faster Malleability - the ability of a substance to be hammered or beaten into thin sheets Ductility - the ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire
6.4 Metallic Properties Metals contain many orbitals separated by very small energy differences They can absorb a wide range of light frequencies, exciting the electrons When excited electrons fall back to the ground state, they emit light at a frequency similar to the absorbed frequency Luster shiny appearance of metals due to radiation or reflection of light
Comparison of Properties
Property Comparison Metallic Ionic Compound Molecular Compound Covalent Network Solid Bonding Melting point / phase Strong attraction between ions and delocalised electrons High melting points, most are solids at room temp Strong electrostatic attraction between positive and negative ions High melting points, always solid at room temp Strong bonds between atoms, electrons are shared Low melting points, gases and liquids at room temp Strong bonds between atoms, electrons are shared High melting points, always solid at room temp Conductivity Conduct electricity and heat easily Conduct electricity when molten and in solution Nonconductive Most are nonconductive (except graphite) Crystalline arrangement Regular crystal arrangement Regular crystal arrangement Weak forces may form molecular solids Regular crystal arrangement Solubility in H 2 O Poor, but some metals react with water Generally good Generally poor, but some substances react with water Poor Example Sodium Sodium chloride Iodine Diamond
Section 6.5 Molecular Geometry
6.5 Objectives Explain VSEPR theory. Predict the shapes of molecules or polyatomic ions using VSEPR theory. Describe dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces and their effects on properties such as boiling and melting points. Explain the shapes of molecules or polyatomic ions using VSEPR theory.
6.5 Intermolecular Forces Intermolecular forces The forces of attraction between molecules The boiling point of a liquid is a good measure of the intermolecular forces between its molecules: the higher the boiling point, the stronger the forces between the molecules. Strength varies, but generally weaker than intramolecular forces (bonds) between atoms or ions 3 types 1. Dipole-dipole (including hydrogen bonding) 2. Dipole-induced dipole 3. Induced dipole-induced dipole (London dispersion forces)
6.5 Intermolecular Forces The strongest intermolecular forces exist between polar molecules. Because of their uneven charge distribution, polar molecules have dipoles. Dipole A molecule or a part of a molecule that contains both positively and negatively charged regions, separated by a short distance. Dipole direction is from the positive pole to the negative pole.
6.5 Dipole-dipole The negative region in one polar molecule attracts the positive region in adjacent molecules Dipole-dipole forces act at short range, only between nearby molecules. Dipole-dipole forces explain, for example the difference between the boiling points of iodine chloride, I Cl (97 C), and bromine, Br Br (59 C).
6.5 Hydrogen Bonding Some H-containing compounds have unusually high boiling points because of a particularly strong type of dipole-dipole force. In compounds containing H F, H O, or H N bonds, the large electronegativity differences between make the bonds highly polar This gives the hydrogen atom a positive charge that is almost half as large as that of a bare proton The small size of the hydrogen atom allows the atom to come very close to an unshared pair of electrons in an adjacent molecule. The strong hydrogen bonding between H 2 0 molecules accounts for many of its characteristic properties.
6.5 Dipole-Induced Dipole A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons Creates a short-range intermolecular force that is weaker than the dipole-dipole force. Induced dipoles account for the fact that a nonpolar molecule, oxygen, O 2, is able to dissolve in H 2 0, a polar molecule.
6.5 London Dispersion Forces Even noble gas atoms and nonpolar molecules can experience weak intermolecular attraction. In any atom or molecule polar or nonpolar the electrons are in continuous motion. As a result, at any instant the electron distribution may be uneven. A momentary uneven charge can create a positive pole at one end of an atom of molecule and a negative pole at the other. This temporary dipole can then induce a dipole in an adjacent atom or molecule. The two are held together for an instant by the weak attraction between temporary dipoles. The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles are called London dispersion forces
6.5 Molecular Geometry Molecular geometry The three-dimensional arrangement of a molecule s atoms, influences properties of molecules Chemical formula does not indicate much about the geometry of molecules Molecular polarity Uneven distribution of molecular charge due to bond polarity and molecular geometry Strongly influences intermolecular forces in liquids and solids
6.5 VSEPR Theory VSEPR Valence-shell, electron pair repulsion Repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible Diatomic molecules are always linear A is the central atom and B the atoms bonded to A Size of B can distort the bond angles, so some variation in molecular shape
6.5 VSEPR Theory Lone pairs occupy space just like bonding pairs do, that is what gives water a bent shape Unshared electron pairs actual repel other electron pairs more strongly than bonding pairs do. That is why water s B atoms are actually closer together than ammonia's http://phet.colorado.edu/en/simulation/molecule-shapes
6.5 Molecular Geometry
6.5 Molecular Geometry
6.5 Sample Problem E Use VSEPR theory to predict the molecular geometry of boron trichloride, BCl 3. Write the Lewis structure for BCl 3 *NOTE* Boron only wants 6 valence electrons in its outer energy level Total valence electrons = 24e - AB 3 molecule = trigonal-planar
6.5 Sample Problem F a. Use VSEPR theory to predict the shape of a molecule of carbon dioxide, CO 2. Draw the Lewis structure. AB 2 molecule = linear
6.5 Sample Problem F b. Use VSEPR theory to predict the shape of a chlorate ion, ClO 3 - Draw the Lewis structure. AB 3 E molecule = Trigonal-pyramidal