Ionic and Covalent Bonds

Transcription

1 Ionic and Covalent Bonds Ionic Bonds Transfer of Electrons When metals bond with nonmetals, electrons are from the metal to the nonmetal The becomes a cation and the becomes an anion. The between the cation and the anion results in an ionic compound. In Lewis Theory, we can show this by from the metal to the nonmetal Example: Potassium and Chlorine When Potassium and Chlorine bond: Potassium transfers its electron to chlorine This transfer gives chlorine an and leaves potassium with an octet in the previous principal shell (now its valence shell) The dot structure of the anion is usually written with with the in the upper right corner (outside the brackets) The positive and negative charges attract one another, resulting in the compound Magnesium and Oxygen Magnesium its 2 valence electrons, forming a charge. Oxygen these two electrons forming a charge and acquiring an octet Lewis theory can help predict the correct for ionic compounds. Ex. Lewis theory predicts K atom for every atom, forming. Sodium and Sulfur Sodium its 1 valence electron to get an octet in the previous principal shell Sulfur must gain electrons to get an octet The compound that forms between sodium and sulfur requires 2 sodium atoms for every one sulfur atom Correct Lewis Structure: Formula:

2 Covalent Bonds: Electrons shared When nonmetals bond with other nonmetals, a is formed. Molecular compounds contain covalent bonds, in which electrons are between atoms rather than transferred In Lewis theory, we represent covalent bonding by allowing atoms to share some of the valence electrons in order to attain (or duets for hydrogen) Hydrogen and Oxygen In water, hydrogen and oxygen share their electrons so that each hydrogen atoms gets a and each oxygen atom gets an. The shared electrons (those appearing in the between the two atoms) count toward the octets (or duets) of both atoms Electrons that are shared between two atoms are called electrons Electrons that are not shared but belong to only one atom are called electrons or. In water, there are bonding pairs of electrons (one between each Hydrogen and the oxygen atom) and lone pair electrons (both on the oxygen atom) Bonding pair electrons are often represented by to emphasize that they form a chemical bond. Single covalent bond When we represent compounds this way, it is known as a. Structural formula for water: Write the structural formulas for: Ammonia (NH 3 ) # of Bonding Pairs: # of Lone Pairs: Methane (CH 4 ) # of Bonding Pairs # of Lone Pairs

3 Molecular Hydrogen Lewis Structure for H: Each hydrogen has electron to share with the other hydrogen atom to form a duet. Bonding Pairs? Lone Pairs? bonding pair, lone pairs This shows why hydrogen exists in nature as a molecule instead of as a single atom Molecular Chlorine Lewis Structure of Cl: If 2 chlorine atoms are paired together, they each get an Bonding Pairs? Lone Pairs? bonding pair, lone pairs This shows why exist in nature as diatomic molecules Double and Triple Bonds In Lewis theory, two atoms may share more than one to get octets. Ex. Oxygen We know oxygen exists as a molecule If we pair two oxygen atoms up, we don t have electrons to give each O atom an octet We can take lone pair electrons and turn them into electrons Each oxygen atom now has an octet because the additional bonding pair counts toward the octet of both oxygen atoms Structural formula: : A bond in which two electron pairs are shared between two atoms In general, double bonds are and than single-bonds Ex. The distance between oxygen nuclei in an oxygen-oxygen double-bond (O=O) is pm In a single bond it is pm

4 A bond in which three electron pairs are shared between two atoms Ex. N 2 In order to have enough electrons to satisfy the octet rule for both N atoms, we need to convert additional lone pairs of electrons into bonding pairs. Triple bonds are even and than double bonds Because these bonds are so strong, Diatomic nitrogen is a fairly molecule in nature. Steps for writing Lewis Structures for Covalent Compounds 1. Write the correct for the molecule Atoms should be in the right Hydrogen atoms are always on the ends ( atoms) -never in the middle of a molecule Molecules tend to be if it contains several atoms of the same type 2. Calculate the for the Lewis structure by adding up the valence electrons for each atom in the molecule If writing the Lewis structure for a ion, the charge of the ion must be considered when calculating the total number of electrons 3. Distribute the electrons among the atoms, giving (or duets for hydrogen) to as many atoms as possible Start by placing electrons between each pair of atoms Then the remaining electrons, first to the terminal atoms, then to the central atom, giving octets to as many atoms as possible 4. If any atoms lack an octet, form or bonds as necessary to give them octets Do this by moving pairs from terminal atoms to bonding regions between atoms. Write the Lewis Structure for CO 2 1. Correct Skeletal Structure 2. Total number of electrons 3. Distribute electrons 4. Form Double/Triple Bonds

5 Lewis Structures for Polyatomic Ions Follow the same 4 steps, but pay special attention to the of the ion when calculating the number of electrons in the Lewis structure Add 1 electron for each charge Subtract 1 electron for each charge We normally show the Lewis structure for a polyatomic ion within and write the charge of the ion in the upper corner Ex. Cyanide Ions 1. Skeleton Structure 2. # of valence electrons 3. Distribute Electrons 4. Double/Triple Bonds 5. Enclose in brackets with charge in the upper right corner Exceptions to the Octet Rule Lewis theory is a simple theory, not sophisticated enough to be correct Ex. NO electrons Can exist as: In cases when we have an number of valence electrons, we write the best Lewis structure we can Boron tends to form compounds with only electrons around the B instead of 8 Ex. BH 3 SF 6 and PCl 5 have electrons around the central atom in their Lewis structure We call these