3 Types of Chemical Bonds

1. Ionic Bonds: refers to the electrostatic forces between oppositely charged particles (usually a metallic and a nonmetallic element).

Ex: NaCl ------ Na+ and Cl- Because Na is positively charged and Cl is negatively charged, when it is put together it is IONIC!

Formed when an atom that easily loses electrons reacts with one that has a high electron affinity.

2. Covalent Bonds: result from the sharing of electrons between two atoms (usually two nonmetallic elements).

Ex: P2O5 Basically, look at the periodic table and see if both elements are nonmetals, if so it is COVALENT!

In sharing, each atom attains a noble gas configuration.

3. Metallic Bonds: are found in solid metals where a metal atom is bonded to another metal. Ex: brass (copper + zinc) or steel (iron + carbon) Any metallic elements have metallic bonds.

Lewis Structures and the Octet RuleLewis Structure: shows how valence electrons are arranged among atoms in a molecule.What are valence electrons.....

determine an element's ability to form chemical bonds. These electrons are the ones residing in the outermost electron shell of the atom, the valence shell.

* To determine valence electrons look at the periodic table. All the elements in each column have the same number of electrons in their outer shells.

- All the elements in the first column (H, Li, Na, K, etc.) have a single valence electron. - The second column elements all of have 2 valence electrons (Be, Mg, Ca, Sr, etc.) - All the transition metals have 2 valence electrons.- The elements in the next column (B, Al, Ga, etc.) have 3 valence electrons.- The elements in the next column (N, P, As, etc.) have 4 valence electrons.- O, S, Se, and the others in this column have 5 valence electrons.- The halogen group, which is the second to last column (F, Cl, Br, etc.) have 7

valence electrons.- The noble gases in the last column (Ne, Ar, Kr, tc.) have all 8 valence

electrons, except for He which one has 2 valance electrons.* If an atom is an ion, you must include the charge also:

- For a positive ion, for each charge subtract one electron. For example, Na+ has 1 valence electron and a positive one charge, so it would be 1-1=0. BUT, it has 8 valence electrons because NOW it has the same electron configuration as Ne. Just as K+ has the same electron configuration as Ar. Therefore, all the Alkali metals will have 8 valence electrons.

- For a negative ion, add one electron for each charge. For example, O2- has 6 valence electrons and a negative 2 charge, so it would be 6+2=8 valence electrons.

Valence electrons

*To show the valence electrons in bonded atoms, we use what is called the electron dot structure, or commonly known as the Lewis dot structure.

- In order to do this, we must use the octet rule.

- What is the octet rule.....

-Octet rule states that the noble gas configurations will be achieved if the Lewis structure shows eight electrons around each atom. Exceptions to the octet rule: Hydrogen, Beryllium, and Boron, who are stable with less than an octet.

Total the valance electrons of all atoms. If it is an ion, add or subtract electrons appropriately. (To learn how to add or subtract valence electrons, refer to how todetermine valence electrons).

Determine the central atom and make a skeletal structure. Hint: Carbon, if present, are usually the central atom, if more than one carbon they are usually linked in chains; hydrogen is never the central atom; halogens are rarely a central atom, but maybe if oxygen is present; oxygen is rarely a central atom, but it might link 2 carbon atoms in a carbon chain. (The central atom is usually that one that appears once in the formula).

Draw the molecular skeleton by placing single bonds (single bonds=2 electrons) to connect all the atoms to the central atom. These electrons are called bonding pairs.

Add electron pairs to all outer atoms until they have an octet and/or you run out of electrons.

If there are remaining electrons put it on the central atom, but if there isn't any electrons left and all atoms have an octet then you are done. If the central atom does not have an octet, pull in electron pairs from outer atoms to for multiple bonds.

IF it is in the third row or lower, it can have up to 12 electrons and exceed the octet rule.

If the structure is an ion, add brackets[ ] and with the appropriate charge. For example, NO3

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H---- always surrounded by a maximum of 2 electrons.

Be---- usually has 4 electrons

B------ usually has 6 electrons

Ex: PO4

1) P-5 valence electrons

O4-6 valence electrons, but there is 4 F, so multiply 6x4=24

Total: 32 valence electrons

2) P would be the central atom, and it would be surrounded by 4 O’s.

3) This is the molecular skeleton for PO4:

Rules to the Lewis Structure

Exceptions to the octet rule:

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4) Distribute the valence electron until each atom has a 8 valence electrons:

5) Since all 32 electrons are used and the central atom, P has an octet the Lewis Dot Structure is complete.

Ex: NH3

1) N-5 valence electrons

H-3x1 valence electron=3

Total: 8 valence electrons

2) N would be the central atom, and it would be surrounded by 3 H’s.

3) This is the molecular skeleton for NH3:

4) Since the H atoms can only have 2

electrons and each H already has a

complete outer shell, we must add the remaining 2 electrons to the central atom, N.

5) Because of the lone pair(non-bonding pair) that was added to the central atom, N it creates an octet, which means it’s now DONE!

Ex: CO2

1) C-4

O-2x6=12

Total: 16 valence electrons

2) C would be the central atom, and it would be surrounded by 2 O’s.

3) This is the molecular skeleton for CO2:

4) Distribute the valence electrons around the atoms until each atom has a eight valence electrons.

5) Since all 16 electrons have been used, we double-check it the central atom has a complete octet. However, C only has four surrounding electrons with the two single bonds on both sides. This means that it doesn’t have an octet, and we must remove one pair of the nonbonding electrons from one of the O atoms and form a double bond between the C and that O atom.

6) Now, we check again to see if it forms an octet. The C atom now has 6 surrounding electrons, which means it still doesn’t have an octet. We have to remove another pair of nonbonding electrons from the other O atom and form another double bond between the C and the other O atom.

7) Now the central atom, C has eight surrounding electrons and both the O atoms have eight surrounding electrons, the Lewis Dot Structure is complete.

Ex: NH4+

1) N-5

H-4x1=4

Total: 9 valence electrons, but have to add one because of the positive one charge. The real total is 8 valence electrons.

2) The central atom is N, and it is surrounded by 4 H’s.

3) The molecular skeleton for NH4+:

4) Since each H atom only have 2 valence electrons, it already has a complete outer shell. All of the valence electron have been used, and the central atom, N is surrounded by 8 electrons, the Lewis Dot Structure is complete. However, the structure is an ion, you must add brackets with the appropriate charge, as shown:

Resonance structures: more than one Lewis structure can be drawn for a substance.

Ex: SO2

1) S-6

O-2x6=12

Total: 18 valence electrons

2) The central atom would be S, and it would be surrounded by 2 O’s.

3) The molecular skeleton for SO2:

4) Distribute the valence electrons until each atom has a eight valence electrons.

5) Add to the remaining electrons to the central atom, S.

6) Make sure that the central atom has an octet. In this case, the S atom only have 6 electrons and we need 8 electrons. That means we don’t have an octet. We must remove one of the nonbonding electrons around one of the O atoms and create a double bond between that O atom and the central atom, S.

However, when removing the nonbonding electron around one of the atoms, the questions remains: which O atom do we remove the nonbonding electrons to form the double bond. The answer is you can remove either or.

Both the structures are basically the same except for the choice of the O atom with which the S atom forms the double bond, so you would write it as a resonance.

The correct way to indicate it is a resonance structure you use a forward/backward arrow that is put in between the two Lewis Structures:

Draw the Lewis Dot Structures for the following. If it has a resonance structure provide all the different Lewis structures for the substance.

1) ClO4-

2) XeF4

3) BeCl2

4) TI4

5) NO31-

(Valence Shell Electron Pair Repulsion Model)

The best arrangement of a given number of electron pairs is the one that minimizes repulsions between them.

Electron Pair Geometry: shows the arrangement of electron pairs.

Ex: NH3

To solve for the electron pair geometry, you must draw the Lewis Dot Structure. Depending on how many electron pairs that determines the electron pair geometry.

Molecular Geometry: shows the arrangement of bonded atoms in space. Note: When describing the shape of molecules, we always give the molecular geometry rather than the electron-pair geometry.

Ex: NH3

Because it has 4 electron pairs, the electron pair geometry is tetrahedral.

Because it has 3 bonded pairs and 1 lone pair(non-bonding pairs), the molecular geometry is trigonal pyramidal.

Practice Exercise #1:

VESPR Theory

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To solve for the molecular geometry, you must draw the Lewis Dot Structure. Depending on how much bonded pairs and lone pairs that determines the molecular geometry.

Ex: PO4

For each of the following provide the Lewis Dot Structure, the number of bonding/non-bonding pairs, the electron pair geometry, and the molecular geometry.

1) SiBr4

2) SbBr3

3) O3

4) H2S

5) KrF2

PolarityPolarity is used to describe the sharing of electrons between atoms.

A non-polar covalent bond is one in which the electrons are shared equally between two atoms. A polar covalent bond is when one atom has a greater attraction for the electrons than the other atom. If the attraction is big enough, then the bond is ionic.In general, to know if it is polar or non-polar just look at their Lewis Dot Structure. If it has a asymmetric shape then it is polar and if it has a symmetric shape it is non-polar.Ex: NH3

Ex: PO4

Because it has 4 bonded pairs and 0 lone pair(non-bonding pairs), the molecular geometry is tetrahedral.

This is polar because of the lone pair that is attached to the central atom. Because of the lone pair NH3 has an asymmetric shape.

This is non-polar because there is no lone pair, so it makes a symmetrical shape.

More Practice #2:

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Hybridization

Sp hybrid orbitals are for covalent bonds(electron shared) not for ionic bonds(electrons transferred).To determine the hybridization just look at the electron pair geometry.

Linear 2 spTrigonal planar 3 sp2

Tetrahedral 4 sp3

Trigonal Bipyramidal 5 sp3dOctahedral 6 sp3d2

Based on their Lewis Structures predict whether each of the following compound is polar or non-polar. Also state their hybridization.

1) SCl42) POCl33) H2S

Intermolecular Forces3 Main Types of Intermolecular forces: London Dispersion Forces, Dipole-Dipole Forces, and Hydrogen Bonding.London Dispersion Forces: is the weakest intermolecular force that causes non-polar substances to condense to liquids and to freeze into solids when the temperature is lowered sufficiently. Ex: PO4

Dipole-Dipole Forces: is the attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. To determine if it is a dipole-dipole force, you must see if the molecules are polar.

Ex: OF2

Hydrogen Bonding: is another dipole force, but this has an attractive force between hydrogen and an electronegative atom. Usually for hydrogen bonding to occur, hydro gen is usually attached to O, N, F. This is the strongest intermolecular force between the three.Ex: NH3

Since PO4 is non-polar, then its intermolecular force would be London Dispersion.

2 is polar, then its intermolecular force would be Dipole-Dipole.

Since NH3 is polar, you might think that its intermolecular force is Dipole-Dipole. However, there is hydrogen in it and it is attached to N, so this is actually Hydrogen Bond

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Electron Pair Geometry # of orbitals needed Hybridization

Practice #3:

Since OF

ElectronegativityElectronegativity: is the ability of an atom in a molecule to attract bonding electrons to itself. This is used to determine whether a given bond will be non-polar covalent, polar covalent, or ionic.Fluorine is the most electronegative element and the least electronegative is Francium. In general electronegativity increases from left to right along a period. It decreases as you go down a group. We use the difference in electronegativity to determine the bond polarity. To determine their electronegativity difference, you will be given a chart.

Electronegativity used to determine bond type: 0 - 0.4 ----- non-polar covalent0.5 – 1.67 ------ polar covalent1.7 or greater ------- ionic bond

Ex: H-H---- 2.1-2.1=0, so it is non-polar covalent. H-F----- 4.0-2.1=1.9, so it is ionic H-Br---- 2.8-2.1=0.7, so it is polar covalent.However, if you don’t have the chart and you were just determining which one has a

higher electronegativity, you just look at how far apart they are from fluorine.List the following in order from increasing electronegativity without looking at the

chart.1) C, N, O2) Si, Ge, Sn3) S, Se, Cl

Lattice EnergyLattice Energy: the energy that is required to separate a mole of solid into ions.

Answers1)

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Practice #4:

Practice #1:

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Bonded Pairs: 4Non-bonded Pairs: 0Electron Pair Geometry: TetrahedralMolecular Geometry: Tetrahedral

Bonded Pairs: 3Non-bonded Pairs: 1Electron Pair Geometry: TetrahedralMolecular Geometry: Trigonal Pyramidal

Bonded pairs: 2Non-bonded Pairs: 1Electron Pair Geometry: Trigonal PlanarMolecular Geometry: Bent 120

Bonded Pairs: 2Non-bonded Pairs: 2Electron Pair Geometry: TetrahedralMolecular Geometry: Bent 109

Bonded Pairs: 2Non-bonded Pairs: 3Electron Pair Geometry: Trigonal BipyramidalMolecular Geometry: T-shaped

Practice #2:

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1) O, N, C 2) Si, Ge, Sn

3) Cl, S, Se

For more help you can try out this link. Think link provides information of how to draw Lewis Dot Structures and at the end it provides you with three different quizzes.

http://chemistry.alanearhart.org/Tutorials/Lewis/lewis-part9.html

Polar or Non-polar: Polar

Hybridization: sp3d

Polar or Non-polar: Non-polarHybridization: sp3

Polar or Non-polar: PolarHybridization: sp3

Practice #3:

Practice #4:

Link: