Unit 2 Beta

Compound Structure and Properties: the one-page sheet

2.1 Types of Chemical Bonds

Atoms bond because nuclei attract the valence electrons of neighboring atoms. Equal sharing gives a nonpolar covalent bond, unequal sharing a polar covalent bond with partial charges, and transfer from a metal to a nonmetal an ionic bond between oppositely charged ions. Metal atoms pool their valence electrons into a sea of delocalized electrons around the cations. The electronegativity difference sets where a bond falls on this continuum.

A covalent bond is a pair of electrons shared between two atoms; both nuclei attract it.
Unequal sharing makes a polar covalent bond: the more electronegative atom gets a partial negative charge (δ−), the other δ+.
An ionic bond is the Coulombic attraction between a cation and an anion, usually a metal with a nonmetal.
In metallic bonding, cations sit in a sea of delocalized valence electrons.
Bond character is a continuum: the larger the electronegativity difference, the more polar and the more ionic the bond.

the electrons end up between the nuclei, holding the atoms together: a chemical bond the shared pair is attracted equally: a nonpolar covalent bond the shared pair spends more time near it, giving partial charges δ− and δ+: a polar covalent bond an electron moves from the metal to the nonmetal, and the cation and anion attract: an ionic bond the cations sit in a sea of delocalized electrons that holds them together: metallic bonding

chemical bond: An attraction that holds two or more atoms or ions together, caused by nuclei attracting electrons.
covalent bond: A bond in which two atoms share a pair of electrons that both nuclei attract.
polar covalent bond: A covalent bond in which the shared electrons are pulled closer to the more electronegative atom; if the atoms are identical the sharing is equal and the bond is nonpolar.
partial charge: A charge smaller than one full electron charge, written δ+ or δ−, on the atoms of a polar bond.
dipole: A separation of positive and negative charge, such as the δ+ and δ− ends of a polar bond; drawn as an arrow pointing toward the negative end.
ionic bond: The Coulombic attraction between oppositely charged ions, usually formed when a metal gives electrons to a nonmetal.
metallic bond: The attraction between metal cations and the sea of delocalized valence electrons around them.
delocalized electrons: Describes electrons that are not held between one pair of atoms but spread over several or many atoms, as in a metal.

2.2 Intramolecular Force and Potential Energy

As two atoms approach, attraction lowers their potential energy until repulsion between the nuclei makes it rise again. The minimum of the potential energy curve gives the bond length (its distance) and the bond energy (its depth): the energy needed to break the bond. Larger atoms form longer bonds, and more shared pairs between the same atoms make shorter, stronger bonds. For ions, Coulomb's law shows that larger charges and smaller distances give stronger attractions.

An intramolecular force acts within a particle: covalent, ionic or metallic bonding.
The minimum of the potential energy curve sits at the bond length; its depth below zero is the bond energy.
Breaking a bond requires energy; forming one releases energy.
Between the same two atoms, a triple bond is shorter and stronger than a double bond, which is shorter and stronger than a single bond.
For ions, Coulomb's law (F ∝ q₁q₂/r²): larger charges and a shorter distance give a stronger attraction.

their potential energy is taken as zero the potential energy falls below zero the potential energy rises steeply the curve has a minimum: its distance is the bond length, its depth the bond energy double and triple bonds are shorter and stronger than single bonds

intramolecular force: A force between the atoms inside one particle: covalent, ionic or metallic bonding.
potential energy curve: A graph of the potential energy of two atoms against the distance between their nuclei; its minimum gives the bond length and the bond energy.
bond length: The distance between the nuclei of two bonded atoms, at the minimum of their potential energy curve.
bond energy: The energy needed to break one mole of a bond, equal to the depth of the potential energy minimum below zero.
single bond: The number of electron pairs shared between two atoms: one in a single bond, two in a double bond, three in a triple bond.

2.3 Structure of Ionic Solids

An ionic solid is a lattice of alternating cations and anions in which every ion is attracted to its oppositely charged neighbors. Coulomb's law sets the lattice energy: larger charges and shorter distances between ion centers give stronger attraction and higher melting temperatures. Ionic solids are hard but brittle, because shifting a layer brings like charges together. They do not conduct as solids, but conduct when melted or dissolved because the ions can then move.

An ionic solid is a repeating 3-D array of cations and anions; each ion is surrounded by opposite charges. There are no separate molecules.
Lattice energy grows with ion charge and shrinks with distance (Coulomb's law). Compare charges first, then sizes, for both compounds.
Strong attractions throughout the lattice give high melting temperatures and hardness.
Ionic solids are brittle: a shifted layer puts like charges side by side.
Ions carry the current: solid, no electrical conductivity; melted or dissolved, good conductivity.

they settle in a repeating three-dimensional lattice, each ion surrounded by opposite charges a lot of energy is needed to free the ions, so the melting temperature is high the lattice energy and melting temperature are larger (MgO far above NaCl) like charges line up and repel, so the solid is brittle the solid does not conduct, but the melt and the solution do

ionic solid: A solid made of a repeating three-dimensional array (a crystal lattice) of cations and anions, each ion surrounded by ions of opposite charge.
lattice energy: The energy needed to separate one mole of an ionic solid into ions far apart; larger ion charges and smaller distances make it larger.
electrical conductivity: How well a substance carries an electric current; it needs charged particles that can move, such as ions in a melt or solution or electrons in a metal.
brittle: Breaking or shattering when struck instead of bending; ionic solids are brittle because shifting a layer brings like charges together.

2.4 Structure of Metals and Alloys

In a metallic solid, cations sit in a sea of delocalized electrons. The mobile electrons make metals good conductors of electricity and heat, and because the sea follows the cations when layers slide, metals are malleable and ductile. An alloy mixes a metal with other elements and keeps metallic bonding. Similar-sized atoms replace host atoms (substitutional, as in brass); much smaller atoms fill the gaps (interstitial, as in steel), pinning the layers so the alloy is harder and less malleable.

A metallic solid is cations in a sea of delocalized electrons, which conduct electricity and heat.
Metals are malleable and ductile: layers slide while the electron sea keeps surrounding each cation.
An alloy is a mixture of a metal with other elements that keeps metallic bonding.
In a substitutional alloy the added atoms are about the same size as the host atoms and take their places (brass).
In an interstitial alloy much smaller atoms fill the gaps, pinning the layers: harder, less malleable (steel).

a metallic solid forms: cations held by delocalized electrons metals conduct electricity and heat well the bonding is kept, so metals are malleable and ductile instead of brittle a substitutional alloy forms, usually a little harder than the pure metal an interstitial alloy forms, harder and less malleable

metallic solid: A solid held together by metallic bonding: metal cations in a sea of delocalized valence electrons.
malleable: The ability of a solid to be hammered or bent into a new shape without breaking (ductile: drawn into wire); metals have it because their layers slide while the electron sea holds them.
alloy: A mixture of a metal with one or more other elements that keeps metallic bonding and metallic properties.
interstitial alloy: An alloy in which much smaller atoms fill the gaps between the host metal's atoms, making it harder and less malleable (steel: carbon in iron).
substitutional alloy: An alloy in which atoms of similar size take the places of some host metal atoms (brass: zinc in copper).

2.5 Lewis Diagrams

A Lewis diagram shows a molecule's or ion's valence electrons as bonding pairs (lines) and lone pairs (dots). Count the valence electrons, adjusting for charge; connect the outer atoms to the central atom; complete the outer octets; put leftovers on the center; and form double or triple bonds if the center is short. Most period 2 atoms follow the octet rule and H has two electrons, but boron can have an incomplete octet, odd-electron species leave one electron unpaired, and atoms from period 3 on can have expanded octets.

A Lewis diagram shows every valence electron: lines for bonding pairs, dot pairs for lone pairs.
Count first: group valence electrons for every atom, plus one per negative charge, minus one per positive charge.
Octet rule: most period 2 atoms end with eight electrons; H with two.
Put the central atom in the middle (never H), complete the outer atoms, put leftovers on the center, then make multiple bonds if the center is short.
Exceptions: incomplete octet (B), odd-electron species (NO), expanded octet from period 3 on (SF₆).

a Lewis diagram shows only valence electrons, counted from each atom's group and the ion's charge each line counts as two electrons for both atoms it joins most period 2 atoms end with an octet; H ends with two outer lone pairs become double or triple bonds incomplete octets, free radicals and expanded octets are the exceptions

Lewis diagram: A drawing of a molecule or ion that shows every valence electron, with bonding pairs as lines and lone pairs as pairs of dots.
central atom: The atom in a molecule or ion that is bonded to two or more other atoms; it is drawn in the middle of a Lewis diagram.
lone pair: A pair of valence electrons on one atom that is not shared in a bond; also called a nonbonding pair.
bonding pair: A pair of electrons shared between two atoms in a covalent bond, drawn as a line in a Lewis diagram.
octet rule: The pattern that atoms of period 2 in stable molecules end up with eight valence electrons around them (hydrogen with two).
expanded octet: A species that does not follow the octet rule: an incomplete octet (B in BF₃), an odd-electron species or free radical (NO), or an expanded octet from period 3 on (S in SF₆).
polyatomic ion: A charged group of atoms held together by covalent bonds, such as NH₄⁺ or SO₄²⁻.

2.6 Resonance and Formal Charge

When several valid Lewis diagrams differ only in where electrons are placed, they are resonance structures, and the real species is a single hybrid with delocalized electrons. Bonds shared in resonance are identical, with fractional bond orders and lengths between single and double bonds. Formal charge, valence electrons minus lone-pair electrons minus half the bonding electrons, compares diagrams: the best one has formal charges nearest zero, with any negative charge on the more electronegative atom.

Resonance structures have the same atom positions and differ only in where electrons are drawn.
The real species is one resonance hybrid with delocalized electrons; it does not flip between diagrams.
Bond order in a hybrid = bonding pairs ÷ number of bonds sharing them (nitrate 4/3, ozone 3/2). Equal, in-between lengths are the evidence.
Formal charge = valence electrons − lone-pair electrons − ½ bonding electrons; the formal charges add up to the overall charge.
Best diagram: formal charges closest to zero, and any negative one on the more electronegative atom.

they are resonance structures, linked by a double-headed arrow the real species is one resonance hybrid with delocalized electrons each bond has a fractional bond order and an in-between length formal charge (valence − lone-pair electrons − ½ bonding electrons) compares them the best diagram has formal charges nearest zero, with negative charge on the more electronegative atom

resonance: The description of a species by two or more Lewis diagrams that differ only in where electrons are placed; the real species is a single hybrid of them.
formal charge: A bookkeeping charge for an atom in a Lewis diagram: valence electrons minus lone-pair electrons minus half the bonding electrons.
bond order: The number of bonding pairs between two atoms (1, 2 or 3); in a resonance hybrid it can be a fraction, such as 4/3 in nitrate.

2.7 VSEPR and Hybridization

VSEPR predicts shape from the Lewis diagram: the electron domains on the central atom (each bond or lone pair counts once) spread as far apart as possible, giving linear, trigonal planar, tetrahedral, trigonal bipyramidal or octahedral arrangements. The molecular geometry names only the atoms, so lone pairs change it, and because lone pairs repel more strongly they shrink bond angles. A molecule is polar when its bond dipoles do not cancel in its shape. Two, three and four domains correspond to sp, sp² and sp³ hybridization; a single bond is one σ bond, and each extra bond in a double or triple bond is a π bond.

VSEPR: electron domains (each bond or lone pair counts once) spread as far apart as possible.
Domains set the arrangement; molecular geometry names the atom positions only.
Lone pairs repel more than bonding pairs and shrink bond angles: CH₄ 109.5°, NH₃ about 107°, H₂O about 104.5°.
Molecular polarity: polar bonds plus an unsymmetrical shape give a net dipole.
Hybridization matches domains: 2 sp, 3 sp², 4 sp³. Single bond 1 σ; double 1 σ + 1 π; triple 1 σ + 2 π.

they repel and spread as far apart as possible (VSEPR) the number of domains sets the arrangement: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral the molecular geometry differs from the arrangement (NH₃ pyramidal, H₂O bent) it pushes bonding pairs together, so bond angles shrink (109.5° to 107° to 104.5°) symmetrical shapes cancel them (nonpolar); lopsided shapes leave a net dipole (polar)

VSEPR: Valence shell electron pair repulsion: the model that electron domains around a central atom repel and spread as far apart as possible, which sets the molecule's shape.
electron domain: A region of electron density around a central atom: one lone pair, or one bond to another atom whether single, double or triple.
molecular geometry: The arrangement of the atoms (not the lone pairs) around a central atom, such as linear, bent, trigonal planar, trigonal pyramidal or tetrahedral.
bond angle: The angle between two bonds that meet at the same atom, such as about 104.5° for H–O–H in water.
molecular polarity: Whether a molecule has a net dipole: it is polar when its bond dipoles do not cancel because of its shape.
hybridization: The mixing of a central atom's valence orbitals to match its electron domains: sp for 2 domains, sp² for 3, sp³ for 4.
sigma bond: A sigma (σ) bond has its shared electrons along the line between the nuclei (every single bond, and the first bond of a multiple bond); each extra bond in a double or triple bond is a pi (π) bond.