Chapter 4 - Chemical Bonding

chemical bond

a link between atoms that results from the mutual electrical attraction between the nuclei and valence electrons of both atoms

Why do atoms bond together

to become more stable?

Ionic Bond

a chemical bond resulting from electrical attractions between large numbers of positive and negative ions, cations, and anions (transfer)

Fixed in place with negative ions

Covalent Bond

a chemical bond resulting from the sharing of electron pairs between 2 atoms

Impure Bonds

range in between ionic and covalent; the degree to which they go can be estimated by comparing their electronegativities (polar covalent is all about them)

Electronegativities Chart

Ionic: 4.0-1.7

Polar Covalent: 1.7-0.3

Non-polar Covalent: 0.3-0.0

Polar Covalent Bond

a covalent bond where the united atoms have an unequal attraction for the shared electrons

Non-polar Covalent Bond

a covalent bond in which the bonding electrons are shared equally by bonded atoms with a resulting balances distribution of charges

Metallic Bonds

bonds between metals where metals give up electrons and are free to move around

Ionic Compounds

compounds composed of positive and negative ions combined so that the positive and negative charges are equal

many of these are 3-D crystalline solids not composed of individual molecules

Formula Units

use them to describe the pattern in an ionic compound

Lattice Energy

the energy released when one mole of an ionic crystalline compound is formed from gaseous ions (the larger the magnitude, the more stable the bonding)

Metallic Bonds

a chemical bond resulting from the attraction between metals, ions, and surrounding mobile electrons; referred to as “electron sea” in the crystal lattice

Heat of Vaporization

the amount of heat needed to change a metal in solid state to a gas

the stronger the metallic bond….

the more heat needed

Molecule

a neutral group of atoms held together by covalent bonds

Diatomic molecules

molecules with only 2 atoms

Molecular Compound

a chemical compound whose simplest units are molecules

When a covalent bond forms…

a bond forms when the attraction between the nucleus of one atom and the electrons of another atom outweighs repulsion

the electrons that are shared travel between each of the bonded atoms orbitals that are overlapping. Most of the time, the electrons are spent in this space

Bond energy

the energy required to break a chemical bond and form neutral atoms

The larger the bond energy…

the shorter the bond length

Bond length

the average distance between two bonded atoms; established once the orbitals overlap and at minimum potential energy

Octet Rule

chemical compounds will form so that each atom by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level

Structural Formula

a modified version of a Lewis Structure; does not show the unshared pair of electrons of the atoms in a molecule

Isomers

compounds with the same formula but a different arrangement of atoms

Triple bond have the highest bond energy and….

have shortest bond lengths

Covalent Networking Bonds

a continuous, 3D network of atoms covalently bonded together

Coordinate Covalent Compound

a covalent bond where just one atom donates both electrons and the other atom accepts them

Polyatomic Ion

a charged group of covalently bonded atoms

Electrons are delocalized

distributed over more than two atoms

Resonance

refers to bonding in molecules that cannot be correctly represented by a single Lewis Structure

metal + metal

metallic bonds

nonmetal + nonmetal

covalent bonds

metal + nonmetal

ionic bonds

ionic bonds are ____ than covalent bonds

stronger

ionic compounds have ____ melting points and boiling points

higher

Ionic compounds are harder than molecular because of their

rigid, crystalline alignment

Ionic compounds are brittle because if structure is broken

the layer will separate completely

Ionic compounds are good conductors when

they are dissolved in water

Molecular geometry gives a

compound’s arrangement in space

VSEPR Theory

describes where bonded atoms are relative to one another and used for predicting molecular geometry; electrons arrange themselves to minimize repulsion because the pairs don’t like being close together

Hybridization Theory

describes how they are bonded

Hybridization

the mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energy

Hybrid Orbitals

orbitals of equal energy produced by the combination of two or more orbitals on the same atom

Intramolecular Forces

within molecule

Intermolecular Forces

between molecules

Intermolecular Forces

These forces are generally weaker than bonds within the molecules or compounds

The strongest intermolecular forces act between

polar molecules

Dipole

equal, but opposite charges separated by a short distance

If dipoles oppose one another (working in opposite directions and symmetrical),

dipoles will cancel and result in a nonpolar molecule

If they are working in the same direction or do not cancel symmetrically

it will result in polar molecule

Dipole-Dipole Forces

forces of attraction between polar molecules

A polar molecule can induce a dipole in a nonpolar molecule by

momentarily attracting its electrons

Induction

the process of inducing or causing something to occur

Induction results in

temporary intermolecular forces and are weaker than regular dipole-dipole forces

Hydrogen Bonding

an intermolecular attraction between a hydrogen and an unshared pair of electrons on a strongly electronegative atom in another molecule

London Dispersion Forces

intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles and induced dipoles