Bonding Flashcards by Donalee Brandel T. Gonzales

an attractive force between two ions or between two atoms

Chemical Bond

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Why do chemical bonds form?

The compound resulting from this is more stable and lower in energy than the separate atoms

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How do chemical bonds form?

Octet Rule

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principle which states that atoms tend to gain, lose, or share electrons in order to achieve a full outer shell with eight electrons

except hydrogen and helium for the outer shell (1s) can occupy only 2

Octet Rule

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An atom is most stable if its outer shell is either filled or contains ___, and it has NO electrons of higher energy.

eight electrons

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Main-group elements in chemistry tend to adopt the electron configuration of the nearest ___.

noble gas

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a chemical bond may be ___ and ___.

ionic and covalent

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types of covalent bonds

polar and nonpolar

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bond formed as a result of the electrostatic attraction between ions of opposite charge
bond formed from the transfer of electrons
usually formed from the reaction of metals with nonmetals

Ionic Bond

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compounds formed by ionic bonds

Ionic Compound

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attractive forces between opposite charges

Electrostatic Attraction

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bond formed as a result of sharing electrons between two nuclei

Covalent Bond

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compounds formed by covalent bonds

Molecular Compounds

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measure of the ability of an atom to pull the bonding electrons toward itself

Electronegativity

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a covalent bond between atoms with the same electronegativity

Nonpolar Covalent Bond

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a covalent bond between atoms with different electronegativities

Polar Covalent Bond

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electronegativity difference of a pure covalent bond

< 0.4

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electronegativity difference of a polar covalent bond

between 0.4 and 1.8

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electronegativity difference of an ionic bond

> 1.8

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a pair of equal and oppositely charged poles separated by a distance

Dipole

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measure of dipole
magnitude of the charge on either atom x distance between the two charges

Dipole Moment

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what molecules have dipoles

Polar Molecules

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electron-dot structures

Lewis Structures

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line-bond structures

Kekulé Structures

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* *NOT* an actual charge * used for bookkeeping of electrons * the charge the atom would have if each bonding electron pair in the molecule were shared *equally* between atoms the difference between the number of valence electrons an atom has when it is not bonded to any other atoms and the number it “owns” when it is bonded

Formal Charge

formula for getting the formal charge

FC = (no. of valence e- in free atom) - (number of valence e- in bonded atom) FC = (no. of valence e- in free atom) - (no. of bonding e- / 2) - (no. of nonbonding e-)

omitting of the covalent bonds and listing atoms bonded to a particular carbon (or nitrogen or oxygen) next to it (with a subscript if there is more than 1 of a particular atom)

Condensed Structures

* Carbon atoms are *NOT* usually shown. * A carbon atom is assumed to be at each intersection of two lines (bonds) and at the end of each line. * Hydrogen atoms bonded to carbon are *NOT* shown. * Atoms other than carbon and hydrogen are shown.

Skeletal Structures

3D structures for drawing chemical structures

* Perspective Drawing * Ball-and-Stick Model * Space-Filling Model

shows 3D shape

Perspective Drawing

shows bond angles accurately

Ball-and-Stick Model

shows atoms in scale

Space-Filling Model

bond lies in the plane of the paper (or screen, when viewed electronically) | straight line ----------------

"normal" bond

bond extends backwards, away from the viewer, so effectively "into" the paper (or screen)

dashed bond | dashed line - - - - - - - - -

bond protrudes forward, towards the viewer, so effectively "out of" the paper (or screen)

wedged bond | horizontal triangle

* geometry based on the *arrangement of atoms in a molecule* * defined by bond angles

Molecular Geometry

* geometry based on *valence electron pairs* (bonding and non-bonding) around a central atom = bonding electron pairs = nonbonding/lone electron pairs * defined by bond angles

Electron Pair Geometry

model for the *prediction of molecular geometry* based on the *minimization of electron repulsion* between regions of electron density around an atom

Valence-Shell Electron-Pair Repulsion (VSEPR) Model

The best arrangement of a given number of electron pairs (bonding and nonbonding) is the one that ____ the repulsions among them.

***minimizes*** the repulsions

arrange the pair repulsions

lone-pair vs lone-pair repulsion > lone-pair vs bonding pair repulsion > bonding pair vs bonding pair repulsion

Without lone pairs: electron pair geometry =

molecular geometry

Molecular geometry can be described in terms of ___.

bond angles

assumes that the electrons in a molecule occupy *overlapping* atomic orbitals of the individual atoms

Valence Bond Theory

assumes the formation of molecular orbitals from the atomic orbitals

Molecular Orbital Theory

type of bond formed when atomic orbitals on neighboring atoms overlap one another

Covalent Bond

head-on overlap of atomic orbitals; stronger

Sigma (σ) Bonds

sideway overlap of atomic orbitals; weaker

Pi (π) Bonds

atomic orbitals obtained when two or more nonequivalent orbitals of the same atom combine for covalent bond formation

Hybrid Orbitals

produces the molecular geometry of the molecule

Hybridization

shape of any hybrid orbital is [same/different] from the shapes of the original atomic orbitals

different

as s character increases, bond length _ and bond strength _

bond length: decreases bond strength: increases | inversely proportional

the greater the electron density in the region of orbital overlap, the _ and _ the bond

stronger and shorter

as s character increases, bond angle _

bond angle: increases | directly proportional

a region of space in a molecule where electrons are most likely to be found

Molecular Orbital

2 ways to combine atomic orbitals into molecular orbitals

1. Additive 2. Subtractive

lower energy, bonding molecular orbital

Additive

higher energy, antibonding molecular orbital

Subtractive

occurs when an electron is shared by more than 2 atoms with π bonds (p orbitals)

π Electron Delocalization

π electron delocalization structures are represented by

resonance contributors

refers to two individual line-bond structures and and their special resonance relationship is indicated by the double-headed arrow between them

Resonance Forms

The only difference between resonance forms is the placement of the _ and _.

π bond and nonbonding valence electrons

A compound with resonance forms _ jump back and forth between two resonance forms.

does NOT

has a single unchanging structure of the two individual forms and has characteristics of both

Resonance Hybrid

slide 80

Bonding Flashcards

2- Structure and Bonding