AP Chemistry

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8. Bonding: general concepts

Types of chemical bonds

system behaves in a way that achieves the lowest possible energy

bond energy: measure of energy required to break a bond

used to gain information about the strength of a bonding information

ionic bonding: electrostatic attraction between oppositely charged ions

ionic compound: results when a metal reacts with a nonmetal

Coulomb’s law:

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E: energy (J)

negative: attractive (opposite)

positive: repulsive (same)

r: distance between ion centers (nm)

Q₁, Q₂: numerical ion charges

ion pair has lower energy than the separated ions

bond length: distance where energy is minimal

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energy terms:

net potential energy that results from the attractions and repulsions among the charged particles

kinetic energy due to motions of electrons

zero point of energy: atoms at infinite separation

at very short distances: energy rises very steeply because of the importance of the repulsive forces when the atoms are very close together

bond length: distance at which system has minimum energy

bond is formed

molecule more stable than two separated atoms by quantity of energy (bond energy)

simultaneous attraction of electrons by protons

pulls protons towards each other

balances proton-proton and electron-electron repulsive forces at distance corresponding to bond length

covalent bonding: electrons are shared by nuclei

polar covalent bond: atoms are not so different that electrons are completely transferred (ionic) but are different enough that unequal sharing results

δ: fractional charge

electrons in bonds not shared equally

Electronegativity

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electronegativity: the ability of an atom in a molecule to attract shared electrons to itself

determining values: Pauling

hypothetical molecule HX

relative electronegativities determined by comparing measured H—X bond energy with expected H—X bond energy (average of H—H and X—X bond energies)

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if H and X have identical electronegativities

(H—X)_act and (H—X)_exp are the same

Δ is 0

if X has greater electronegativity

H—X is polar

H has charge δ⁺ and X has charge δ⁻

(H—X)_act > (H—X)_exp

electronegativity trends

across a period: generally increases

down a group: generally decreases

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Bond polarity and dipole moments

dipolar (has dipole moment): has center of positive charge and center of negative charge

arrow points to negative charge

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any diatomic (two-atom) molecule that has a polar bond will also show a molecular dipole moment

polyatomic molecules can exhibit dipolar behavior

e.g. water

oxygen atom has greater electronegativity than hydrogen atoms

water molecule behaves in electric field as if it has two centers of charge (positive and negative)

some molecules have polar bonds but do not have a dipole moment

individual bond polarities arranged in such a way that they cancel each other out

e.g. carbon dioxide

opposing bond polarities cancel out due to linear structure

no preferential way for the molecule to line up in an electric field

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Ions: electron configurations and sizes

electron configuration of compounds

two nonmetals: covalent bond

share electrons in a way that completes the valence electron configurations of both atoms

both nonmetals attain noble gas electron configurations

nonmetal and representative-group metal: binary ionic compound

valence electron configuration of nonmetal achieves electron configuration of next noble gas atom

valence orbitals of metal are emptied

both ions achieve noble gas electron configurations

Predicting formulas of ionic compounds

“ionic compound” usually refers to solid state

ions close together

minimizes same-sign repulsions and maximizes opposite-sign attractions

contrasts with gas phase of ionic substance

ions far apart

pair of ions may get close enough to interact but large collections of ions do not exist

predicting compound

example: calcium and oxygen

consider valence electron configurations

Ca: [Ar]4s²

O: [He]2s²2p⁴

electronegativity of oxygen (3.5) is much greater than that of calcium (1.0)

electrons will be transferred from calcium to oxygen to form oxygen anions and calcium cations

number of electrons transferred

noble gas configurations are generally most stable

oxygen needs two electrons to fill 2s and 2p orbitals and achieve configuration of neon

calcium can lose two electrons to achieve configuration of argon

two electrons transferred from calcium to oxygen

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chemical compounds always electrically neutral (same quantities of positive and negative charges)

equal numbers of Ca²⁺ and O²⁻ ions

empirical formula: CaO

example: aluminum and oxygen

aluminum

configuration [Ne]3s²3p¹

loses three electrons to form Al³⁺ and achieve neon configuration

Al³⁺ and O²⁻ ions form

compound must be electrically neutral so empirical formula is Al₂O₃

metals losing electrons to form cations

Group 1A: lose one electron

Group 2A: lose two electrons

Group 3A: lose three electrons

nonmetals gaining electrons to form anions

Group 7A (halogens): gain one electron

Group 6A: gain two electrons

hydrogen typically behaves as nonmetal and can gain one electron to form hydride ion (H⁻) with electron configuration of helium

exceptions

tin: Sn²⁺ and Sn⁴⁺

lead: Pb²⁺ and Pb⁴⁺

bismuth: Bi³⁺ and Bi⁵⁺

thallium: Tl⁺ Tl³⁺

Sizes of ions

plays important role in determining:

structure and stability of ionic solids

properties of ions in aqueous solution

biologic effects of ions

impossible to define precisely sizes of ions

radii most often determined from measured distances between ion centers in ionic compounds

assumption about how distance should be divided up between the two ions

considerable disagreement among ionic sizes in various sources

factors that influence ionic size

relative sizes of an ion and its parent atom

positive ion

formed by removing electron(s) from neutral atom

resulting cation is smaller than parent atom

negative atom

formed by adding electron(s) to neutral atom

anion is larger than parent atom

sizes vary depending on positions of parent elements in periodic table

down a group: ion size increases

across a period: complicated

change from predominantly metals (left) to nonmetals (right)

period contains both elements that give up electrons (cations) and accept electrons (anions)

isoelectronic ions: ions containing the same number of electrons

number of protons changes

electrons experience greater attraction so ions become smaller

Energy effects in binary ionic compounds

resulting ionic solid forms because aggregated oppositely charged ions have lower energy than original elements

lattice energy: change in energy that takes place when separated gaseous ions are packed together to form an ionic solid

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example:

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reaction broken into steps

sublimation (solid → gas) of solid lithium

ionization of lithium atoms to form Li⁺ ions in gas phase

dissociation of fluorine molecules

formation of F⁻ ions from fluorine atoms in gas phase

formation of solid lithium fluoride from gaseous Li⁺ and F⁻ ions

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main motivation for formation of ionic compound rather than covalent compound: strong mutual attractions among ions

Lattice energy calculations

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k: proportionality constant (depends on structure of solid and electron configurations of ions)

Q₁ and Q₂: charges on the ions

r: shortest distance between the centers of the cations and anions

lattice energy has negative sign when Q₁ and Q₂ have opposite signs

bringing cations and anions together is an exothermic process

process becomes more exothermic as:

ionic charges increase

distances between the ions in the solid decrease

Partial ionic character of covalent bonds

probably no totally ionic bonds between discrete pairs of atoms

percent ionic character of a bond

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ionic character increases with electronegativity difference

no bonds reach 100% ionic character (no individual bonds are completely ionic)

compounds with more than 50% ionic character are normally considered to be ionic solids

many substances contain polyatomic ions, which contain covalent bonds

any compound that conducts an electric current when melted is classified as ionic

Covalent chemical bond: model

concept summaries

what is a chemical bond: forces that cause a group of atoms to behave as a unit

why chemical bonds occur: bonds result from the tendency of a system to seek its lowest possible energy

when collections of atoms are more stable (lower in energy) than the separate ions

useful to interpret molecular stability in terms of chemical bonds (human invention)

method for dividing up energy evolved when a stable molecule is formed from its component atoms

arbitrarily represents a quantity of energy obtained from the overall molecular energy of stabilization

Models: an overview

model: attempt to explain how nature operates on the microscopic level based on experiences in the macroscopic world

originate from observations of properties of nature

example: concept of bonds arose from observations that most chemical processes involve collections of atoms and that chemical reactions involve rearrangements of the ways atoms are grouped

tendency towards lower energy; collections of atoms occur because the aggregated state has lower energy than the separated atoms

molecules can be thought of as collections of common fundamental components

can help chemists systemize reactions of the millions of existing compounds

physically sensible

makes sense that atoms can form stable groups by sharing electrons

shared electrons give a lower energy state because they are simultaneously attracted by two nuclei

supports the existence of discrete bonds that are relatively independent of the molecular environment

delocalization of electrons: electrons assumed to be free to move throughout the entire molecule

fundamental properties of models

a model does not equal reality: models are human inventions, always based on an incomplete understanding of how nature works

models are often wrong: based on speculation and always oversimplifications

models tend to become more complicated as they age: as flaws are discovered, they are “patched” and more detail is added

understand the assumptions inherent in a particular model before using it to interpret observations or make predictions

simple models usually involve very restrictive assumptions and can be expected to yield only qualitative information

must understand its strengths and weaknesses and ask only appropriate questions

when a model is wrong, we often learn much more than when it is write: usually means we do not understand some fundamental characteristics of nature

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