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AP CHEMISTRY CHAPTER 8 BONDING. bond energy - energy required to break a chemical bond -We can measure bond energy to determine strength of interaction. ionic compound - a metal reacts with a nonmetal.
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bond energy- energy required to break a chemical bond-We can measure bond energy to determine strength of interaction
ionic compound- a metal reacts with a nonmetal • Ionic bonds form when an atom that loses electrons easily reacts with an atom that has a high affinity for electrons. The charged ions are held together by their mutual attraction. • Ionic bonds form because the ion pair has lower energy than the separated ions. All bonds form in order to reach a lower energy level.
Bond length- the distance where the energy is at a minimum. We have a balance among proton-proton repulsion, electron-electron repulsion, and proton-electron attraction.In H2, the two e- will usually be found between the two H atoms because they are spontaneously attracted to both protons. Therefore, electrons are shared by both nuclei. This is called covalent bonding.
Polar covalent bonds occur when electrons are not shared equally. One end of the molecule may have a partial charge. This is called a dipole. + H F H H + - O -
Electronegativity- the ability of an atom in a molecule to attract shared electrons to itself.-determined by comparing the measured bond energy and the expected bond energy. Expected H—X = H—H bond energy + X—X bond energybond energy 2 Electronegativity difference = (Actual H—X bond energy) – (expected H—X bond energy) If X has a greater electronegativity than H, the e-’s are closer to X and the molecule is polar. If the electronegativities are the same, the molecule is nonpolar.
Periodic Trends-Electronegativity generally increases across a period and decreases down a group. It ranges from 0.79 for cesium to 4.0 for fluorine.
+ - H—F polar H—H nonpolar has dipole moment O+ + H H O S O bent, polar O - has dipole moment planar no dipole momentCH4 tetrahedral NH3 trigonal pyramidal no dipole moment has dipole moment
Electron Configurations: Stable compounds usually have atoms with noble gas electron configurations. Two nonmetals react to form a covalent bond by sharing electrons to gain valence electron configurations.
When a nonmetal and a group A metal react to form a binary ionic compound, the ions form so that the valence electron configuration of the nonmetal is completed and the valence orbitals of the metal are emptied to give both noble gas configurations.
Ions form to get noble gas configurations. -exceptions in Group A metals: Sn2+ & Sn4+ Pb2+ &Pb4+ Bi3+ & Bi5+ Tl+ & Tl 3+ Metals with d electrons will lose their highest numerical energy level electrons before losing their inner d electrons.
Size of Ions Positive ions (cations) are smaller than their parent atoms since they are losing electrons. Negative ions (anions) are larger than their parent atoms since they are gaining electrons. Ion size increases going down a group.
Isoelectronic ions–ions containing the same number of electrons O2-, F-, Na+, Mg2+, Al3+ all have the Ne configuration. They are isoelectronic.*** For an isoelectronic series, size decreases as Z increases.
Lattice energy- the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. Na+(g) + Cl-(g) NaCl(s) If exothermic, the sign will be negative and the ionic solid will be the stable form. We can use a variety of steps to determine the heat of formation of an ionic solid from its elements. This is called the Born-Haber cycle. See example on page 355.
Lattice energy can be calculated using the following: where k is a proportionality constant that depends on the structure of the solid and the electron configuration of the ions. Q1 & Q2 are the charges on the ions. r is the distance between the center of the cation and the anion. Since the ions will have opposite charges, lattice energy will be negative (exothermic). The attractive force between a pair of oppositely charged ions increases with increased charge on the ions or with decreased ionic sizes.
The Structure of Lithium Fluoride
There are probably no totally ionic bonds. Percent ionic character in binary compounds can be calculated. Percent ionic character increases with electronegativity difference. See Figure 8.12, pg. 360. Compounds with more than 50% ionic character are considered to be ionic (electronegativity diff. of about 1.7).
The Relationship Between the Ionic Character of a Covalent Bond and the Electronegativity Difference of the Bonded Atoms
Polyatomic ions are held together by covalent bonds. We call Na2SO4 ionic even though it has 4 covalent bonds and 2 ionic bonds. Ionic compound- any solid that conducts an electrical current when melted or dissolved in water Salt- an ionic compound
A chemical bond is a model “invented” by scientists to explain stability of compounds. A bond really represents an amount of energy. The bonding model helps us understand and describe molecular structure. It is supported by much research data. However, some data suggests that electrons are delocalized. That is, they are not associated with a particular atom in a molecule.
Single bond- one pair of shared electrons • Double bond- two pair of shared electrons • Triple bond- three pair of shared electrons
Bond energies and bond lengths are given on page 374. We can use bond energies to calculate heats of reaction. H = D(bonds broken)- D(bonds formed) 2H2 + O2 2H2O Ex. H = [2(432) + 495] –[4(467)] = -509 kJ 2 H-H O=O 4H-O exothermic
Bonding Models: Molecular Orbital Model- Electrons occupy orbitals in a molecule in much the same way as they occupy orbitals in atoms. Electrons do not belong to any one atom. -very complex model
Localized electron model- • molecules are composed if atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms • traditional model
lone pair- pair of electrons localized on an atom (nonbonding)shared pair or bonding pair- electrons found in the space between atoms
Lewis structure -shows how the valence electrons are arranged among the atoms in the molecule
The most important requirement for the formation of a stable compound is that the atoms achieve noble gas configurationsionic [ Na ]+ [Cl]- only valence electrons are includedmolecular H2O H – O - H
duet rule- hydrogen forms stable molecules when it shares two electrons H:H-filled valence shellWhy does He not form bonds?Its valence orbitals are already filled.octet rule – most elements need 8 electrons to complete their valence shell Cl-Cl
Rules for writing Lewis structures1. Add up the number of valence electrons from all atoms.2. Use 2 electrons to form a bond between each pair of bound atoms. A dash represents a pair of shared electrons.3. Arrange the remaining electrons to satisfy the duet rule for H and the octet rule for most others.
Ex. CO2# of valence electrons = 4 + 6 + 6 = 16 O – C – O This uses 20 electrons! O = C = O
NH3 has 8 valence electrons H N- H H
HCN HCN has 10 valence electrons. H-C≡N
NO+ NO+ has 5 + 6 -1 = 10 electrons N≡O
CO32- Carbonate has 4 + 18 + 2 = 24 valence electrons. O 2- C O O
Exceptions:Boron and beryllium tend to form compounds where the B or Be atom have fewer than 8 electrons around them. BF3 = 24 valence electronsF B F FCommon AP equation: NH3 + BF3 H3NBF3
Some elements in Period 3 and beyond exceed the octet rule.Ex. SF6 S has 12 electrons around it48 valence electrons F F F F S F F
d orbitals are used to accommodate the extra electrons.Elements in the 1st or 2nd period of the table can’t exceed the octet rule because there is no d sublevel.If the octet rule can be exceeded, the extra electrons are placed on the central atom.
See examples of exceptions on pg 375.Ex. I3-, ClF3, RnCl2 I - I - I F F - Cl - F Cl - Rn - Cl
Resonance--occurs when more than one valid Lewis structure can be written for a particular moleculeactual structure is an average of all resonance structures-this concept is needed to fit the localized electron model (electrons are really delocalized)
