Bonding: General Concepts

Bonding: General Concepts • Bonding • determines chemical and physical properties of substances • graphite vs. diamond • both consist of carbon atoms bonded together • graphite - slippery, soft • diamond - hard, crystalline • SiO2 vs. CO2 • SiO2 - hard, crystalline, found in sand and quartz • CO2 - gaseous

Types of Chemical Bonds • Bonds • forces that hold groups of atoms together and make them function as a unit

Types of Chemical Bonds • Ionic Bond • electrostatic attraction between closely packed oppositely charged ions • formed when a metal reacts with a nonmetal • a metal loses electrons easily to form a positive ion which is attracted to the negative ion formed when the nonmetal gains electrons

Types of Chemical Bonds • Coulomb’s law • tells us the energy of interaction between a pair of ions • E = 2.31 x 10-19J.nm . Q1Q2 r r = the distance between the ion centers (in nm) Q = charge on the ion

Types of Chemical Bonds • This energy of interaction, E • has a negative sign when ions with opposite signs are brought closer together • indicates energy is released as ions are brought closer together • I.e., there is an attractive force between oppositely charged ions

Types of Chemical Bonds • The energy of interaction • becomes more negative with larger charges • the attractive force is greater • becomes more negative as r decreases, I.e., as the ions get closer together • The more negative the energy of interaction, the stronger the ionic bond.

Types of Chemical Bonds • Covalent Bonding • electrons are shared between nuclei • electrons are shared equally between identical atoms • electrons may be shared unequally between two different atoms, resulting in a polar covalent bond

Types of Chemical Bonds • Bringing two hydrogen atoms together results in • proton-proton repulsion • electron-electron repulsion • proton-electron attraction • A bond will form between two hydrogen atoms if the energy of the aggregate (H2) is less than the energy of the two separate atoms

Types of Chemical Bonds • A system will achieve the lowest possible energy • In the case of the two hydrogen atoms • minimize proton-proton and electron-electron repulsive forces • maximize proton-electron attractive forces • Bond length • the distance between the hydrogen atoms where the energy is at a minimum

Types of Chemical Bonds • Interaction of two hydrogen atoms:

Types of Chemical Bonds • H2 molecule • electrons lie between the two nuclei • electrons are attracted by both nuclei resulting in stability

Electronegativity • Electronegativity • the ability of an atom in a molecule to attract shared electrons to itself • I.e., …how much an atom “wants” electrons within a bond • determined by Linus Pauling (1901 - 1995) • Nobel prize in Chemistry and Peace

Electronegativity • See table of E.N.’s • F is the most E.N. atom with an electronegativity value of 4.0 • Cs is the least E.N. atom with an electronegativity value of 0.7 • Periodic Trend • going across a period • E.N. increases • going down a column • E.N. decreases

Electronegativity • Given H-X • X has a greater E.N. than H • X will have a partial negative charge while H will have a partial positive charge • The H-X bond is a polar covalent bond • The H-X bond has a partial ionic component • The greater the difference in E.N. between H and X, the more polar the bond, or the greater the ionic component of the bond

Electronegativity • Arrange the following bonds in order of increasing polarity: • H-H, O-H, Cl-H, S-H, and F-H • (H-H < S-H < Cl-H < O-H < F-H) • (DEN = 0, 0.4, 0.9, 1.4, 1.9)

Bond Polarity and Dipole Moments • Place a polar molecule in an electric field • the molecule will line up so that its “negative” end will line up with the positive pole and the “positive end” will line up with the negative pole

Bond Polarity and Dipole Moments • Dipole Moment • a polar molecule has a dipole moment • a polar molecule has a center of positive charge and a center of negative charge • Ex:

Bond Polarity and Dipole Moments • Some molecules have polar bonds, but are nonpolar • these molecules have no dipole moment • due to the overall geometry of the molecule, the bond polarities cancel out, so the molecule has no net dipole moment • Ex:

Ions: Electron Configurations and Sizes • An atom in a stable compound has a noble gas electron configuration • In ionic compounds • nonmetals gain electrons to achieve a noble gas electron configuration • metals lose electrons to achieve a noble gas electron configuration • Na+ - • Cl- -

Ions: Electron Configurations and Sizes • Group I forms +1 ions • Group II forms +2 ions • Group III forms +3 ions • Group VI forms -2 ions • Group VII forms -1 ions • Exceptions: Pb+2 and Pb+4; Sn+2 and Sn+4; Bi+3 and Bi+5; Tl+ and Tl+3

Ions: Electron Configurations and Sizes • Size of Ions • determines structure and stability of ionic solids • determines properties of ions in solution • determines biological effects of ions • cannot define the precise size of ions • different sources may report different values for ionic sizes

Ions: Electron Configurations and Sizes • Size of Ions • Positive ion • formed by the loss of an electron from the parent atom • cation is smaller than the parent due to the loss of electron pair repulsion and the increased nuclear charge felt by each electron • Negative ion • formed by the parent atom gaining an electron • anion is larger than the parent atom due to increased electron pair repulsion

Ions: Electron Configurations and Sizes • Periodic trend for ionic size • going down a group • ionic radius increases • going across a period • not a regular increase or decrease due to metals losing electrons and nonmetals gaining electrons

Ions: Electron Configurations and Sizes • Isoelectronic Ions • ions with the same number of electrons • Na+, Mg+2, Al+3, F-, O-2, N-3 • all contain 10 electrons • all have the same electron configuration as Ne • in terms of size, N-3>O-2>F->Na+>Mg+2>Al+3 • the ion with the greater number of protons in an isoelectronic series will be the smallest due to the greater nuclear charge pulling the electrons in closer

Ions: Electron Configurations and Sizes • Arrange the ions Se-2, Br-, Rb+, and Sr+2 in order of decreasing size • Choose the largest ion in each of the following groups • Li+, Na+, K+, Rb+, Cs+ • Ba+2, Cs+, I- , Te-2 • (Se-2 > Br- > Rb+ > Sr+2) • (Cs+) • (Te-2)

Formation of Binary Ionic Compounds • Just how stable is an ionic compound? • How strong is a particular ionic bond? • Look at the lattice energy to answer these questions • Lattice energy • the change in energy when separated gaseous ions are packed together to form an ionic solid • M+(g) + X-(g) --> MX(s)

Formation of Binary Ionic Compounds • See handout of lattice energy calculations

Partial Ionic Character of Covalent Bonds • Ionic character increases as the electronegativity difference increases • None of the bonds, even in typical ionic compounds, actually reach 100% ionic character • Actually, the typical ionic compounds have bonds that are 50% or greater in ionic character…in the gaseous state • In the solid state, the existence of ions is favored, I.e., there is 100% ionic character

Partial Ionic Character of Covalent Bonds • An ionic compound • conducts electricity when melted

The Covalent Chemical Bond • Chemical bond • a force that cause a group of atoms to behave as a unit • Why do chemical bonds form? • So a system can achieve the lowest possible energy • so a system can become more stable

The Covalent Chemical Bond • CH4 + 1652 kJ ---> C + 4 H • it takes 1652 kJ to break one mole of methane molecules to one mole of carbon atoms and 4 moles of hydrogen atoms • C + 4H ---> CH4 + 1652 kJ • 1652 kJ are released when one mole of methane molecules are made from one mole of carbon atoms and 4 moles of hydrogen atoms

The Covalent Chemical Bond • Energy Diagram for formation of CH4:

The Covalent Chemical Bond • Since CH4 contains 4 C-H bonds, we can divide the 1652 kJ/mol CH4 by 4 C-H bonds to determine the energy released from forming one C-H bond • 1652 kJ = 413 kJ released/mole C-H bond 4

Covalent Bond Energies & Chemical Reactions • Bond Energies • average energy required to break a particular bond • breaking a C-H bond will require different amounts of energy depending on the “molecular environment” (depending on what else is bonded to the C, for example)

Covalent Bond Energies & Chemical Reactions • Process Energy Required(kJ/mol) • CH4(g) --> CH3(g) + H(g) 435 • CH3(g) --> CH2(g) + H(g) 453 • CH2(g) --> CH(g) + H(g) 425 • CH(g) --> C(g) + H(g) 339 • Total 1652 • Average 1652/4 = 413

Covalent Bond Energies & Chemical Reactions • See tables of Average Bond Energies and Average Bond Length • What is the relationship between the number of shared electron pairs and the bond length? • What is the relationship between the number of shared electron pairs and the bond energy (bond strength)?

Covalent Bond Energies & Chemical Reactions • Bond length/energy for selected bonds Bond Bond Type Bond Length (pm) Bond Energy (kJ/mol) C-C single 154 347 C=C double 134 614 C=C triple 120 839 C-O single 143 358 C=O double 123 745 C-N single 143 305 C=N double 138 615 C=N triple 116 891

Bond Energy and Enthalpy • Use bond energies to calculate approximate energies for reactions (DHrxn) • Breaking bonds requires energy, I.e., energy must be added to the system, an endothermic process • Making bonds releases energy, an exothermic process

Bond Energy and Enthalpy • DHrxn = SD (bonds broken) - SD (bonds formed) (energy required) (energy released) • Ex: Using the bond energies, calculate the DHrxn for the reaction: • CH4(g) + 2 Cl2(g) + 2 F2(g) --> CF2Cl2(g) + 2HF(g) + 2 HCl(g) • (DHrxn = -1194 kJ)

Localized Electron Bonding Model • Localized Electron Bonding Model • applies to all kinds of molecules (simple and complicated) • a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms • electron pairs are localized on a particular atom or in the space between two atoms

Localized Electron Bonding Model • Lone pairs • electrons localized on a particular atom • Bonding pairs (or shared pairs) • electrons localized in the space between two atoms • On to Lewis structures!

Exceptions to the Octet Rule • Boron can form compounds so that boron has fewer than eight electrons…I.e., it can “get by” with six electrons • BF3

Exceptions to the Octet Rule • Beryllium can “get by” with four electrons • BeF2

Resonance • Sometimes more than one valid Lewis structure can be drawn for a given molecule • NO3-

Resonance • Physical evidence shows that NO3- has three equivalent bonds • The correct description of NO3- is not one of the three Lewis structures, but an “average” of the three Lewis structures

Resonance • Resonance occurs when more than one valid Lewis structure can be written for a particular molecule • The actual electron structure would be the average of the resonance structures • In resonance structures, the arrangement of the nuclei is the same, but the placement of the electrons differ

Resonance • Resonance structures • electrons are actually delocalized then…the bonding electrons are not necessarily found in one place…but can be found throughout the molecule

Odd Electron Molecules • Some nitrogen compounds form molecules with odd number of electrons (another exception to the octet rule) • NO, NO2

Formal Charge • A way to handle odd electron molecules • A way to handle molecules or polyatomic ions that have atoms that can exceed the octet rule • Determining the formal charge on atoms allows us to determine which Lewis structure is most valid

Formal Charge • Calculating oxidation numbers was one way to calculate a “charge” on an atom or ion…but does not really work for determining validity of a Lewis structure • Formal Charge • the difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule

Bonding: General Concepts