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 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 0.3 Nonpolar covalent H H (0.0) 0.31 2.0 Polar covalent δ+ δ H F (1.9) >2.0 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.
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. 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
Ionic Bonding A positively charged ion, or cation, is produced when an atom loses one or more valence electrons. For example, magnesium belongs to Group 2 of the periodic table, so magnesium atoms have two valence electrons. A magnesium atom loses two valence electrons and produces a magnesium cation with a charge of 2+. Mg 2+
Ionic Bonding A negatively charged ion, or anion, is produced when an atom gains one or more valence electrons. For example, nitrogen belongs to Group 15 of the periodic table, so nitrogen atoms have five valence electrons. A nitrogen atom gains three valence electrons and produces a nitrogen anion with a charge of 3-. N 3-
Ionic Bonding An ionic compound is a compound composed of cations and anions. Although they are composed of ions, ionic compounds are electrically neutral. The total positive charge of the cations equals the total negative charge of the anions.
Ionic Bonding The forces that hold anions and cations together in ionic compounds are called ionic bonds. The chemical formula of an ionic compound is 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
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.