Ionic and Covalent Bonds

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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

 Lewis Structures:

 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)

 Lewis Structure:

 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: __________

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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

 Lewis Structure:

 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 (NH3 )

 # of Bonding Pairs:  # of Lone Pairs:

 Methane (CH4 )

 # of Bonding Pairs  # of Lone Pairs

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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

 Structural Formula:

 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

 Structural Formula:

 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

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 _________________________________ – A bond in which three electron pairs are shared between two atoms

 Ex. N2

 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.

 Structural Formula:

 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 CO2

1. Correct Skeletal Structure

2. Total number of electrons

3. Distribute electrons

4. Form Double/Triple Bonds

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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. BH3

 SF6 and PCl5 have ____________________________ electrons around the central atom in their Lewis structure

 We call these _______________________________________________