#1: Structure Determines Properties Flashcards
Atomic Number
Number of protons in the nucleus of a specific atom.
Wave Functions
The solutions to arithmetic expressions that express the energy of an electron in an atom.
Orbital
Strictly speaking, a wave function. It is convenient however, to think of an orbital in terms of the probability of finding an electron at some point relative to the nucleus, as the volume inside the boundary surface of an atom, or the region in space where the probability of finding an electron is high.
Principal Quantum Number
The quantum number (n) of an electron that describes its energy level. An electron with n=1 must be an s electron; one with n=2 has s and p states available.
Shell
The group of orbitals that have the same principal quantum number n.
Boundary Surfaces
The surface that encloses the region where the probability of finding an electron is high (90-95%).
Spin
Synonymous with spin quantum number.
Spin Quantum Number
One of the four quantum numbers that describe an electron. Can have a value of either +0.5 or -0.5.
Pauli Exclusion Principle
No electrons can have the same set of four quantum numbers. An equivalent expression is that only two electrons can occupy the same orbital, and then only when they have opposite spins.
Nodal Surfaces
A plane drawn through an orbital where the algebraic sign of a wave function changes. The probability of finding an electron at a node is 0.
Hund’s Rule
When two orbitals are of equal energy, they are populated by electrons so that each is half-filled before either one is doubly occupied.
Valence Elctrons
Outermost electrons, the ones most likely to be involved in chemical bonding and reactions. Maximum number of electrons in a valance shell is 8.
Noble Gases
Or rare gases. Characterized by an extremely stable “closed-shell” electron configuration and are very unreactive. Full valance shell. Group 8A elements.
Compounds
An assembly of two or more atoms with properties different from the individual atoms.
Chemical Bond
Attractive force between atoms in a compound.
Ionic Bond
Force of attraction between oppositely charged species (ions).
Very common in inorganic compounds, but rare in organic compounds.
Cations
Positively charged ions.
Anions
Negatively charged ions.
Ionization Energy
Large amount of energy that must be transferred to any atom to dislodge an electron.
In general, ionization energy increases across a row in the periodic table.
Endothermic
Term describing a process or reaction that absorbs heat. Energy has a + sign.
Exothermic
Term describing a reaction or process that gives off heat. Energy change has a - sign.
Election Affinity
Energy change for addition of an electron to an atom.
Energy change associated with the capture of an electron by an atom.
Electrostatic Attractions
Or coulombic attractions. Attractive forces between two oppositely charged particles.
Covalent Bond
Chemical Bond between two atoms that results from their sharing of two electrons.
Lewis Structure
A chemical formula in which electrons are represented by dots. Two dots (or a line) between two atoms represent a covalent bond in a Lewis structure. Unshared electrons are explicitly shown, and stable Lewis structures are those in which the octet rule is satisfied.
Bond Dissociation Enthalpy
For a substance A:B, the energy required to break the bond between A and B so that each retains one of the electrons in the bond.
Unshared Pairs
In a Lewis structure, pairs of valance electrons not involved in bonding.
Octet Rule
In forming compounds they gain, lose, or share electrons to achieve a stable electron configuration characterized by 8 valence electrons.
Double Bond
Bond formed by the sharing of four electrons between two atoms.
Triple Bond
Bond formed by the sharing of six electrons between two atoms.
Polar Covalent Bond
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 polar covalent.
Electronegativity
The tendency of an atom to attract electrons in a covalent bond toward itself.
Flourine is the most electronegative atom. Oxygen is second.
Electrostatic Potential Map
The charge distribution in a comlecule represented by mapping the interaction energy of a point positive charge with the molecules electric field on the van der Waals surface.
Electronegativity Scale
Electronegativity increases from left to right across a row in the periodic table. Of the second-row elements, the most electronegative is flourine, the least electronegative is lithium.
Electronegativity decreases going down a column. Of the halogens, fluorine is the most electronegative, then chlorine, then bromine, then iodine. Indeed, fluorine is the most electronegative of all atoms, oxygen is second.
The greater the electronegativity difference between two elements, the more polar the bond between them.
Bond Dipole Moments
The dipole moment of a bond between two atoms.
Dipole Moment
A dipole exists whenever opposite charges are separated from each other. A dipole moment is the product of the amount of the charge e multiplied by the distance d between the centers of charge.
u = e x d
debye, D
Unit customarily used for measuring dipole moments.
1D = 1 x 10^-18 esu x cm
Formal Charges
The charge, either positive or negative, on an atom calculated this way. Used to express molecular stability.
Formal Charge = Group number in periodic table - Electron Count
Electron Count = 0.5(Number of shared electrons) + Number of unshared electrons
Systematic Approach to Writing Lewis Structures
1) The molecular formula is determined experimentally.
2) Based on the molecular formula, count the number of valance electrons.
3) Given the connectivity, connect bonded atoms by a shared electron pair bond (:) represented by a dash (-).
4) Count the number of electrons in the bonds (twice the number of bonds), and subtract this from the total number of valence electrons to give the number of electrons that remain to be added.
5) Add electrons in pairs so that as many atoms as possible have eight electrons. It is usually best to begin with the most electronegative atom. (Hydrogen is limited to two electrons.)
6) If one or more atoms (excluding hydrogens) has fewer than eight electrons, use an unshared pair from an adjacent atom to form a double or triple bond to complete the octet. Use one double bond for each deficiency of two electrons to complete the octet for reach atom.
7) Calculate formal charges.
Molecular Formula
Chemical formula in which subscripts are used to indicate the number of atoms of each element present in one molecule. In organic compounds, carbon is cited first, hydrogen second, and the remaining elements in alphabetical order.
Connectivity
Order in which a molecule’s atoms are connected. Synonymous with constitution.
Isomers
Different compounds that have the same molecular formula.
Constitutional Isomers
Isomers that differ in respect to the order in which the atoms are connected.
Stereoisomers
Isomers with the same constitution but that differ in respect to the arrangement of their atoms in space. Stereoisomers may be either enantiomers or diastereomers.
Resonance
When two or more Lewis structures that differ only in the distribution of electrons can be written for a molecule, no single Lewis structure is sufficient to describe its true electron distribution.
Resonance Hybrid
The true structure of the various Lewis formulas, called contributing structures, that can be written for the molecule.
Localized
Electrons are either shared between two atoms in a covalent bond or are unshared electrons in a single atom.
Delocalized
Association of an electron with more than one atom. The simplest example is the shared electron pair (covalent) bond. Delocalization is important in conjugated pi electrion systems, where an electron may be associated with several carbon atoms.
Curved Arrows
Arrows that show the direction of electron flow in chemical reactions; also used to show differences in electron placement between resonance forms.
When Can Resonance Be Considered?
1) The connectivity must be the same in all contributing structures; only the electron positions may vary among the various contributing structures.
2) Each contributing structure must have the same number of electrons and the same net charge. The formal charges of individual atoms may vary among the various Lewis structures.
3) Each contributing structure must have the same number of unpaired electrons.
4) Contributing structures in which the octet rule is exceeded for second-row elements make no contribution. (The octet rule may be exceeded for elements beyond the second row.)
Which resonance form contributes more?
1) As long as the octet rule is not exceeded for second-row elements, the contributing structure with the greater number of covalent bonds contributes more to the resonance hybrid. Maximizing the number of bonds and satisfying the octet rule normally go hand in hand. This rule is more important than the other two.
2) When two or more structures satisfy the octet rule, the major contributor is the one with the smallest separation of oppositely charged atoms.
3) Among structural formulas that satisfy the octet rule and in which one or more atoms bears a formal charge, the major contributor is the one in which the negative charge resides on the most electronegative atom.
What is the effect of resonance?
Electron delocalization stabilizes a molecule. Resonance is a way of showing electron delocalization. Therefore, the true electron distribution is more stable than any of the contributing structures. The degree of stabilization is greatest when the contributing structures are of equal stability.
Condensed Formula
Structural formula in which subscripts are used to indicate replicated atoms of groups, as in (CH3)2CHCH2CH3
Bond-Line Formula or Carbon Skeletal Diagram
Formula in which connections between carbons are shown but individual carbons and hydrogens are not.
We assume that there is a carbon atom at every vertex and at the end of the line. Hydrogens attached to a vertex or to the end of a line are left out.
When hereroatoms are present, hydrogens attached to them are shown.
Heteroatoms
Atoms in an organic molecule that aren’t carbon or hydrogen.
VSEPR Model
Valence Shell Electron-Pair Repulsion Model. An electron pair, either a bonded pair or an unshared pair, associated with a particular atom will be as far away from the atom’s other electron pairs as possible.
Method for predicting the shape of a molecule based on the notion that electron pairs surrounding a central atom repel one another. Four electron pairs will arrange themselves in a tetrahedral geometry, three will assume a trigonal planar geometry, and two will adopt a linear arrangement.
Tetrahedral Angle
109.5 degrees
The angle between one line directed from the center of a tetrahedron to a vertex and a second line from the center to a different vertex. This angle is 109 degrees 28’.
Repulsive Force
Bonded Pair-Bonded Pair < Unshared Pair-Bonded Pair < Unshared Pair-Unshared Pair
Unshared Pairs have the most repulsion and thus will force bonded electrons to be closer to one another.
VSEPR Model and Multiple Bonds
Multiple bonds are treated as a single unit in the VSPER model. Formaldehyde is a trigonal planar molecule in which the electrons of the double bond and those of the two single bonds are maximally separated. A linear arrangement of atoms in carbon dioxide allows the electrons in one double bond to be as far away as possible from the electrons in the other double bond.
Molecular Dipole Momenet
Resultant of all of the individual bond dipole moments of a substance. Some molecules, such as carbon dioxide, have polar bonds, but lack a dipole because their geometry causes the individual CO bond dipoles to cancel.
We can use electronegativity to tell us about the polarity of bonds and combine that with VSEPR to predict whether the molecule has a dipole moment.
Double-Barbed Arrow
Shows the movement of a pair of electrons, either a bonded or lone pair.
Single-Barbed Arrow
Shows the movement of one electron.
Acid
Substance that ionizes to give protons when dissolved in water. Strong acids ionize completely, weak do not.
Base
Substance that ionizes to gibe hydroxide ions when dissolved in water. Strong bases ionize completely, weak do not.
Acidity Constant
Measures strength of a weak acid. It’s the equlibrium constant Ka for its ionization in aqueous solution.
pKa
Expresses the strength of an acid. pKa = -logKa. pKa avoids exponential unlike Ka.
Bronsted-Lowry View
An acid is a proton donor, and a base is a proton acceptor. The reaction that occurs between an acid and a base is proton transfer.
B: + H-A B+-H + :A-
In this equation, the base uses an unshaired pair of electrons to remove a proton from an acid. The base is converted to its conjugate acid, and the acid is converted to its conjugate base. A base and its conjugate acid always differ by a single proton. Likewise, an acid and its conjugate base always differ by a single proton.
In the Bronsted-Lowry view, an acid doesn’t dissociate in water; it transfers a proton to water. Water acts as a base.
H2O + HA H3O+ + A-
The systematic name for the conjugate acid of water is oxonium ion. Its common name is hydronium ion.
Conjugate Acid
The species formed from a Bronsted base after it has accepted a proton.
Conjugate Base
The species formed for a Bronsted acid after it has donated a proton.
Strength of Conjugates
The stronger the acid, the weaker its conjugate base.
The stronger the base, the weaker its conjugate acid.
Structures Affecting Acidity
Depends on:
1) The strength of the bond to the atom from which the proton is lost.
2) The electronegativity of the atom from which the proton is lost.
3) Electron delocalization in the conjugate base.
Bond Strength
Bond strength decreases going down a group in the periodic table. As the halogen X becomes larger, the H-X bond becomes longer and weaker and acid strength increases.
Inductive Effects
Structural effects that are transmitted through the bonds.
Inductive effects depend on the electronegativity of the substituent and the number of bonds between it and the affected site. The more bonds between the substituent and the affected site, the less effective the inductive effect is.
Acid-Base Equilibria
In any proton-transfer reaction:
Acid + Base Conjugate Acid + Conjugate Base
we are concerned with the question of whether the position of equilibrium lies to the side of products or reactants. There’s an easy way to determine this. The reaction proceeds in the direction that converts the stronger acid and the stronger base to the weaker acid and the weaker base.
Stronger Acid + Stronger Base –> Weaker Acid + Weaker Base
K > 1
The reaction will be favorable when the stronger acid is on the left and the weaker acid is on the right. The equilibrium lies to the side of the acid that holds the proton more tightly.
Equilibrium Constant for an Acid-Base Reaction
Keq = Ka of reactant acid/Ka of product acid
= 10^-pKa of reactant acid/10^-pKa of product acid
Using pKa to Analyze Acid-Base Equilibria
1) They permit clear-cut distinctions between strong and weak acids and bases. A strong acid is one that is stronger than H3O+. A weak acid is one that’s weaker.
2) The strongest acid present in significant amounts at equilibrium after a strong acid is dissolved in water is H3O+. The strongest acid present in significant amounts when weak acid is dissolved in water is the weak acid itself.
Lewis Acids and Lewis Bases
An acid is an electron-pair acceptor, and a base is an electron-pair donor.
An unshared pair of electrons from the Lewis base is used to form a covalent bond between the Lewis Acid and the Lewis Base.
Lewis Acid/Lewis Base Complex
The species that results by covalent bond formation between a Lewis acid and a Lewis base.
Substitution
Included in Lewis Acids/Bases idea where one atom of group replaces another.
Nucleophiles
An atom or ion that has an unshared electron pair which can be used to form a bond to carbon. Lewis Bases.
Trophiles
An atom or ion that seek an electron pair to form a bond. Lewis Acids.
