In nature, only the noble gas elements exist as uncombined atoms. They are monatomic - consist of single atoms. All other elements need to lose or gain electrons To form ionic compounds Some elements share electrons To form covalent compounds
A neutral group of atoms held together by covalent bonds is called a molecule Some molecules are made up of the same element Those molecules are called diatomic elements 7 naturally occurring diatomics are : H, N, O, F, Cl, Br, I
Molecules can also be made of atoms of different elements. A compound composed of molecules is called a molecular compound. A molecular formula gives the recipe for a molecular compound. A molecular formula shows how many atoms of each element a substance contains.
The molecular formula of water is H 2 O. Notice that the subscript written after an element s symbol indicates the number of atoms of each element in the molecule. If there is only one atom, the subscript 1 is omitted.
A molecular formula reflects the actual number of atoms in each molecule. The subscripts are not necessarily the lowest whole-number ratios. For example, the formula for peroxide is H 2 O 2 Each molecule of peroxide contains 2 hydrogen atoms and 2 oxygen atoms
In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases. The octet rule states that chemical compounds form so each atom (through gaining, losing, or sharing electrons) will have 8 valence electrons Exception: atoms trying to be like helium
We use Lewis Dot Diagrams to show covalent bonding However, we do not need to put the dots in the same order as before We need to put them in singles before we can pair them up
In the F 2 molecule, each fluorine atom contributes one electron to complete the octet. Notice that the two fluorine atoms share only one pair of valence electrons. That is a single covalent bond When we show the bonding, we use Lewis structures Structural formulas are a neater way to show bonding
A pair of valence electrons that is not shared between atoms is called an unshared pair In F 2, each fluorine atom has three unshared pairs of electrons.
A double covalent bond is a bond that involves two shared pairs of electrons.
Similarly, a bond formed by sharing three pairs of electrons is a triple covalent bond.
Even when electrons are being shared, the sharing is not equal The bonding pairs of electrons in covalent bonds are pulled between the nuclei of the atoms sharing the electrons.
When the atoms in the bond pull equally, the bonding electrons are shared equally, and each bond formed is a nonpolar covalent bond A nonpolar covalent bond always results when an element bonds with itself
A polar covalent bond, is a covalent bond between atoms in which the electrons are shared unequally. The more electronegative atom attracts more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge. δ+ δ H O
The electronegativity difference between two atoms tells you what kind of bond is likely to form. Electronegativity Differences and Bond Types Electronegativity difference range Most probable type of bond Example 0.0 0.30 Nonpolar covalent H H (0.0) 0.31 2.00 Polar covalent δ+ δ H F (1.9) >2.00 Ionic Na + Cl (2.1)
The polar nature of the bond may also be represented by an arrow pointing to the more electronegative atom. H O
Electron dot structures fail to reflect the three-dimensional shapes of molecules. The electron dot structure and structural formula of methane (CH 4 ) show the molecule as if it were flat and merely twodimensional. Methane (structural formula) Methane (Lewis Structure)
To determine the 3D shape of the molecule, we use VSEPR (valence shell electron pair repulsion) theory The theory states that repulsion between sets of valence level electrons surrounding an atom causes these sets to be oriented as far apart as possible Or simply, unshared pairs of electrons want to be as far apart as possible
The hydrogens in a methane molecule are at the four corners of a geometric solid called a regular tetrahedron. In this arrangement, all of the H C H angles are 109.5, the tetrahedral angle.
Unshared electron pair 107 The molecule ammonia (NH 3 ) is trigonal pyramidal shape. However, one of the valence-electron pairs is an unshared pair and it repels the bonding pairs, pushing them together. The measured H N H bond angle is only 107, rather than the tetrahedral angle of 109.5.
105 The water molecule is planar (flat) but bent. With two unshared pairs repelling the bonding pairs, the H O H bond angle is compressed to about 105.
CO 2 is a linear molecule The carbon in a carbon dioxide molecule has no unshared electron pairs. The double bonds joining the oxygens to the carbon are farthest apart when the O=C=O bond angle is 180 Carbon dioxide (CO 2 ) 180 No unshared electron pairs on carbon
Here are some common molecular shapes. Linear Trigonal planar Bent Pyramidal Tetrahedral Trigonal bipyramidal Octahedral Square planar
In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals. It gives information regarding both shape and bonding
Hybridization in a single bond Recall that the carbon atom s outer electron configuration is 2s 2 2p 2, but one of the 2s electrons is promoted to a 2p orbital to give one 2s electron and three 2p electrons. You might suspect that one bond in methane would be different from the other three. In fact, all the bonds are identical. The one 2s orbital and three 2p orbitals of a carbon atom mix to form four sp 3 hybrid orbitals at the tetrahedral angle of 109.5.
Bonding Many factors contribute to the properties of compounds Bond type Shape Interparticle Forces
Ionic Bonding An ionic compound is a compound composed of cations and anions. Although they are composed of ions, ionic compounds are electrically neutral. Anions and cations have opposite charges and attract one another by means of electrostatic forces. The electrostatic forces that hold ions together in ionic compounds are called ionic bonds.
Ionic Bonding A chemical formula shows the numbers of atoms of each element in the smallest representative unit of a substance. The chemical formula of an ionic compound refers to a ratio known as a formula unit. A formula unit is the lowest whole-number ratio of ions in an ionic compound.
Ionic Bonding Most ionic compounds are crystalline solids at room temperature. The component ions in such crystals are arranged in repeating threedimensional patterns.
Ionic Bonding Ionic compounds have the strongest of all interparticle forces The attraction between the oppositely charged particles creates a strong bond As a result, ionic compounds: Have high boiling and melting points Tend to be solids Have a high solubility in water Good conductor of dissolved or melted solids
Covalent compounds have weaker interparticles forces than ionic compounds Although the sharing of electrons does create a bond, it requires less energy to break that bond As a result, covalent compounds: Have low melting and boiling points Tend to be gases, liquids, or solids (depending on the other molecular interactions) High to low solubility in water Poor to non conducting of dissolved solids
Molecules can be attracted to each other by a variety of different forces. Intermolecular attractions are weaker than either ionic or covalent bonds. Among other things, these attractions are responsible for determining whether a molecular compound is a gas, a liquid, or a solid at a given temperature.
Dipole interactions occur when polar molecules are attracted to one another. The electrical attraction occurs between the oppositely charged regions of polar molecules. Dipole interactions are similar to, but much weaker than, ionic bonds.
Hydrogen bonds are attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom. The other atom may be in the same molecule or in a nearby molecule Hydrogen bonds are the strongest of the intermolecular forces.
This relatively strong attraction, which is also found in hydrogencontaining molecules is called a hydrogen bond. Hydrogen bond
Dispersion forces, the weakest of all molecular interactions, are caused by the motion of electrons. They occur even between nonpolar molecules. When the moving electrons happen to be momentarily more on the side of a molecule closest to a neighboring molecule, their electric force influences the neighboring molecule s electrons to be momentarily more on the opposite side. The strength of dispersion forces generally increases as the number of electrons in a molecule increases.
Fluorine and chlorine have relatively few electrons and are gases at ordinary room temperature and pressure because of their especially weak dispersion forces. Bromine molecules therefore attract each other sufficiently to make bromine a liquid under ordinary room temperature and pressure. Iodine, with a still larger number of electrons, is a solid at ordinary room temperature and pressure.
Bonding This table summarizes some of the characteristic differences between ionic and covalent (molecular) substances. Characteristics of Ionic and Molecular Compounds Characteristic Ionic Compound Molecular Compound Representative unit Formula unit Molecule Bond formation Transfer of one or more electrons between atoms Sharing of electron parts between atoms Type of elements Metallic and nonmetallic Nonmetallic Physical state Solid Solid, liquid, or gas Melting point High (usually above 300 C) High (usually below 300 C) Solubility in water Usually high High to low Electrical conductivity of aqueous solution Good conductor Poor to nonconducting
Bonding Collection of water molecules Array of sodium ions and chloride ions Molecule of water Formula unit of sodium chloride Chemical Formula H 2 O NaCl Chemical Formula
Metallic Bonding To model the valence electrons in a metal, it would consist of closely packed cations and loosely held valence electrons rather than neutral atoms. The valence electrons of atoms in a pure metal can be modeled as a sea of electrons. The valence electrons are mobile and can drift freely from one part of the metal to another.
Metallic Bonding Metallic bonds are the forces of attraction between the free-floating valence electrons and the positively charged metal ions. These bonds hold metals together Metals are good conductors of electric current because electrons can flow freely in the metal. As electrons enter one end of a bar of metal, an equal number of electrons leave the other end.
Metallic Bonding Both metals and ionic compounds form crystal structures. However, they have different configurations of electrons. The sea of electrons surrounding cations in a metal allows metals to be ductile and malleable. Ionic crystals will fracture under pressure. Sea of electrons Metal cation Metal Force Force Ionic crystal Nonmetal anion Metal cation Strong repulsions