Structure Determines Properties

Structure Determines Properties

Chapter 1

Overview

• A review the relationship between structure and properties is what chemistry is all about. – Lewis Structures– Arrhenius, Bronsted-Lowry, and Lewis pictures of

acids and bases– Concludes with the effects of structure on acidity

and basicity.

Atoms, Electrons, and Orbitals

• Each atom has– Protons, +– Electrons, -– Neutrons, no charge

• A neutral atom– # protons = # electrons

• Atomic number (Z)– Number of protons


• Wave functions are the mathematical descriptions giving the probability of locating the electron.

• Wave functions are called orbitals.• Orbitals are descried by specifying their size,

shape, and directional properties.


• Rules of electron configuration– Aufbau Principle

• The rule that electrons occupy the orbitals of lowest energy first.

– Pauli Exclusion Principle• An atomic orbital may describe at most two electrons, each

with opposite spin direction.– Hund’s Rule

• Electrons occupy orbitals of the same energy in a way that makes the number or electrons with the same spin direction as large as possible.


• The letter s is preceded by the principal quantum number which specifies the shell and is related to the energy of the orbital.– Hydrogen: 1s1 – Helium: 1s2

• Electrons possess the property of spin. – The spin quantum number of an electron can

have a value of either +1/2 or -1/2.


• Valence electrons of an atom are the outermost electrons, the ones most likely to be involved in chemical bonding and reactions

• Four main group elements, the number of valence electrons is equal to its group number in the periodic table.

• Structure determines properties and the properties of atoms depend on atomic structure.


• Neutral atoms have as many electrons as the number of protons in the nucleus.

• Those electrons occupy orbitals in order of increasing energy, with no more than two electrons in any one orbital.

• most frequently encountered atomic orbital in this text are s orbitals and p orbital.

Ionic Bonds

• The attractive forces between atoms in a compound is a chemical bond

• Ionic bond is the fore of attraction between oppositely charged ions.

• Ionic bonds are very common in inorganic compounds, but rare in organic ones.

Ionic Bonds

• Positively charged ions are called cations• Na(g) → Na+(g) + e-

• Negatively charged ions are called anions.• Cl(g) + e- → Cl-(g)

Ionic Bonds

• An ionic bond is the force of electrostatic attraction between two oppositely charged ions.

• Ionic bonds in which carbon is the cation or anion are rare.

Covalent Bonds, Lewis Structures and Octet Rule

• Covalent or shared electron pair is the bond formed when electrons are shared.

• Lewis structures are structural formulas in which electrons are represented as dots.


• The valence electrons that are not involved in the bonding are called unshared pairs.

• Octet rule– In forming compounds, they gain, lose, or share

electrons to achieve a stable electron configuration characterized by eight valence electrons.


• The most common kind of bonding involving carbon is covalent bonding.

• A covalent bond is the sharing of a pair of electrons between two atoms.

• Lewis structures are written on the basis of the octet rule, which limits second-row elements to no more eight electrons in their valence shells.

• In most of its compounds, carbon has four bonds.

Double and Triple Bonds

• By pairing the unshared electrons of one carbon with its counterpart of the other carbon, a double bond results and the octet rule is satisfied for both carbons.


• The ten valence electrons of acetylene (C2H2) can be arranged in a structural formula that satisfies the octet rule when six of them are shared in a triple bond between the carbons.


• Many organic compounds have double or triple bonds to carbon. Four electrons are involved in a double bond: six in a triple.

Polar Covalent Bonds, Electronegativity and Bond Dipoles

• Electrons in covalent bonds are not necessarily shared equally by the two atoms that they connect.


• If one atom has a greater tendency to attract electrons toward itself than the other, the electron distribution is polarized, and the bond is described as polar covalent.

• The tendency of an atom to attract the electrons in a covalent bond toward itself defines its electronegativity.


• A dipole exists whenever opposite charges are separated from each other, and a dipole moment m is the product of the amount of the charge e multiplied by the distance d between the centers of charge.


• An important difference between a C-H bond and a C-O bond, and that is the direction of the dipole moment. In a C-H bond the electrons are drawn away from H, toward C. In a C-O bond, electrons are drawn from C toward O.


• When two atoms that differ in electronegativity are covalently bonded, the electrons in the bond are drawn toward the more electronegative element.

Formal Charge

• Lewis structures frequently contain atoms that bear a positive or negative charge.

• Electrons in covalent bonds are counted as if they are shared equally by the atoms they connect, but unshared electrons belong to a single atom.

Formal Charge

• Formal charge = Group number in periodic table – Electron count

• Electron count = ½(Number of shared electrons) + Number of unshared electrons

Formal Charges

• It will always be true that a covalently bonded hydrogen has no formal charge (formal charge = 0).

• It will always be true that a nitrogen with four covalent bonds has a formal charge of +1. (A nitrogen with four covalent bonds cannot have unshared pairs, because of the octet rule.)

• It will always be true that an oxygen with two covalent bonds and two unshared pairs has no formal charge.

• It will always be true that an oxygen with one covalent bond and three unshared pairs has a formal charge of -1.

Formal Charges

• Counting electrons and assessing charge distribution in molecules is essential to understanding how structure affects properties. A particular atom in a Lewis structure may be neutral, positive, or negatively charged. The formal charge of an atom in the Lewis structure of a molecule can be calculated by comparing its electron count with that of the neutral atom itself.

Structural Formulas of Organic Molecules

• Different compounds that have the same molecular formula are called isomers. If they are different because their atoms are connected in a different order, they are called constitutional isomers.

Structural Formulas of Organic Molecules

• Isomers can be either constitutional isomers (differ in connectivity) or stereoisomers (differ in arrangement of atoms in space).

Resonance

• Sometimes more than one Lewis formula can be written for a molecule, especially if the molecule contains a double or triple bond.

Resonance

• Lewis formulas show electrons as being localized; they either are shared between two atoms in a covalent bond or are unshared electrons belonging to a single atom. In reality, electrons distribute themselves in the way that leads to their most stable arrangement. This sometimes means that a pair of electrons is delocalized, or shared by several nuclei.

The Shapes of Simple Molecules

• The shapes of molecules can often be predicted on the basis of valence shell electron-pair repulsion.

Molecular Dipole Moments

• Knowing the shape of a molecule and the polarity of its various bonds allows the presence or absence of a molecular dipole moment and its direction to be predicted.

Curved Arrows and Chemical Reactions

• Curved arrows increase the amount of information provided by a chemical equation by showing the flow of electrons associated with bond making and bond breaking.

Acids and Bases: Arrenhius View

• According to the Arrhenius definitions, an acid ionizes in water to produce protons (H+) and a base produces hydroxide ions (OH-). The strength of an acid is given by its equilibrium constant Ka for ionization in aqueous solution.

• Or more conveniently by its pKa:pKa = - log10Ka

Acids and Bases: Bronsted-Lowry

• According to the Bronsted-Lowry definitions, an acid is a proton donor and a base is a proton acceptor.

B: + H-A → B-H + :A-

• The Bronsted-Lowry approach to acids and bases is more generally useful than the Arrhenius approach.

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How Structure Affect Acid Strength

• The strength of an acid depends on the atom to which the proton is bonded.

• The two main factors are the strength of the H-X bond and the electronegativity of X.

• Bond strength is more important for atoms in the same group of the periodic table; electronegativity is more important for atoms in the same row.

• Electron delocalization in the conjugate base, usually expressed via resonance between Lewis structures, increase acidity by stabilizing the conjugate base.

Acid-Base Equilibria

• The position of equilibrium in an acid-base reaction lies to the side of the weaker acid.Stronger acid + stronger base Weaker acid +

weaker base• The equilibrium will be to the side of the acid

that holds the proton more tightly.Keq =

Lewis Acids and Lewis Bases

• A Lewis acid is an electron-pair acceptor. A Lewis-base is an electron-pair donor.

• The Lewis approach incorporates the Bronsted-Lowry approach as a subcategory in which the atom that accepts the electron pair in the Lewis acid is a hydrogen.

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Structure Determines Properties