ch. 6 chemical bonding chemistry – mrs. demott. key ideas: atoms bond and stick together!! ...

Ch. 6 Chemical BondingChemistry – Mrs. DeMott

Key Ideas: Atoms bond and stick together!!

Everyone wants to look like a Noble Gas (filled octet)

Look at your periodic table. Who really wants to “give up” electrons?

Who really wants to “take” electrons?



Look at your periodic table. Who really wants to “give up” electrons? Alkali metals

Who really wants to “take” electrons?



Look at your periodic table. Who really wants to “give up” electrons? Alkali metals

Who really wants to “take” electrons? Halogens

I. Intro. to chemical bonding: A. Defn: an attraction between the

nucleus and the valence electrons of two or more atoms.

B. Types of chemical bonds: 1. Ionic bond: the electrostatic

attraction between positive and negative ions.

(forms ions – ex. Na and Cl) Happens mostly between atoms with large differences in electronegativity.

a. Electrons have been transferred from one atom to another.

2. Covalent bond: the sharing of electrons between atoms.

(happens when electronegativities are similar)

a. Non-polar covalent- electrons are shared equally by the atoms.

b. Polar covalent- the electrons are shared unequally by the atoms. The electrons are found nearer the more electronegative atom.

3. Metallic bond: the metal atoms give up their valence electrons, which move freely about in the metal.

C. In general, nonmetal atoms form covalent bonds with each other. Metal atoms form metallic bond with each other. Metal atoms form ionic bonds with nonmetal atoms.

D. The type of bond formed is determined by the difference in electronegativity. The greater the difference, the more ionic the bond. See Fig. 6-2, pg 162.

The degree to which bonding between atoms of two elements is ionic or covalent can be estimated by the difference in the elements’ electronegativities. Ex. F and Cs (4.0 – 0.7 = 3.3) see pg. 151

>1.7 ionic < 1.7 covalent

0 – 0.3 = nonpolar covalent (electrons shared more equally)

0.3-1.7 = polar covalent (electrons found nearer the more electronegative atom)

Bonding between sulfur and hydrogen: Electronegativity Difference

Bond Type

More-negative atom

Bonding between sulfur and hydrogen: Electronegativity Difference 2.5-2.1 =

0.4

Bond Type

More-negative atom


0.4

Bond Type - polar covalent

More-negative atom


0.4

Bond Type - polar covalent

More-negative atom sulfur

Try sulfur and cesium: Electronegativity Difference:

Bond Type:

More-negative atom:

Try sulfur and cesium: Electronegativity Difference: 2.5-0.7 =

1.8

Bond Type: ionic

More-negative atom: sulfur

Try sulfur and chlorine: Electronegativity Difference:

Bond Type:

More-negative atom:

Try sulfur and chlorine: Electronegativity Difference: 3.0 – 2.5 = 0.5

Bond Type: polar - covalent

More-negative atom: chlorine

II. Covalent Bonding: A. Defn: When two or more atoms join by sharing

electrons

S

B. Terminology: 1. Molecule- 2 or more atoms joined by a

covalent bond a. Diatomic molecule- two atoms of the

same element joined by a covalent bond

b. Examples: H2 C12, F2 Br2, O2, At2, N2.

2. Molecular compound- a chemical compound whose smallest unit is a molecule.

1. Chemical formula- it shows the relative number of atoms of each element in the compound.

a. The formula is made up of symbols and subscripts. It is a shorthand way to show the make-up of a compound.

b. Examples: Water: H20, Hydrogen chloride: HCl, Calcium bromide: CaBr2

4. Molecular formula-shows the kind and number of atoms in a single molecule.

5. Bond length- - the minimum potential energy distance between two bonded atoms.

6. Bond energy-amount of energy required to break a chemical bond. Energy is usually released when a bond is formed.

E. The formation of a chemical bond usually releases energy. Therefore it requires energy to break that bond. This is called bond energy. Bond length is the average distance between the atoms in the bond. See Table 6-1, pg 168 for examples of bond energy.

F. When compounds are formed, the octet rule usually applies. Most atoms involved attempt to get or share in 8 electrons to achieve an octet in their outer-most energy level.

7. Octet rule- chemical compounds tend to form

in such a way so that each atom receives an octet in its highest occupied energy level by gaining, losing, or sharing electrons.

8. Lewis structures- a formula that uses the symbols of the elements involved to represent the kernel ( nucleus and inner shell electrons)

and dot pairs to represent the shared and unshared electrons in the covalent bonds. Shared electrons may be represented by a dash ( one for each pair shared)

a. Structural formula- a Lewis-type formula that shows only the shared pairs of es,

b. Lewis structure for C02:

Rules for Lewis Structures: 1. Determine the type and number of atoms in the

molecule. 2. Write the electron-dot notation for each type of atom in

the molecule. 3. Determine the total number of valence electrons in the

atoms to be combined. 4. Arrange the atoms to form a skeleton structure for the

molecule. If carbon is present, it is the central atom. If not, the least-electronegative atom is central (except Hydrogen – it is never central).

5. Add unshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is surrounded by eight electrons.

6. Count the electrons in the structure to check that the number of valence electrons used equals the number available.

9. Electron Dot Symbol- The valence electrons of an atom are shown as dots around the symbol for the atom. This was shown last chapter.

a. Electron Dot symbol for oxygen:

C. Single covalent bond- share 1 pair of e - :

1. Like atoms: a. Hydrogen molecule:

Orbital H _____

H _____ Structural H-HElectron Dot H : HMolecular H2

b. Chlorine molecule: Orbital

Cl ___ ____ ____ ____ ____ ____ ____ ____ ____

Cl ___ ____ ____ ____ ____ ____ ____ ____ ____

1s 2s 2p 3s 3p

b. Chlorine molecule

Structural Cl – Cl

Electron Dot see mini movie

Molecular Cl2

Oxygen:

Structural: O=O

Electron Dot

Molecular: O2

E. Triple Covalent bond – sharing 3 pair of electrons1. Nitrogen molecule:2. Orbital:

3. Structural: N= N4. Electron Dot:5. Molecular: N2

6. Note: Double and triple covalent bonds are called multiple bonds.

F. Resonance structures (hybrids) 1. Resonance is bonding in molecules that

cannot be represented by a single Lewis structure. There may be two or more arrangements that work for that compound.

2. Example: Ozone (03) There are two resonance hybrids possible. (They “resonate” between the two.)

3. To indicate resonance, a double headed arrow is placed between the structures.

III. Ionic Bonding: A. Defn. : Electrons are transferred

from one atom to another. This gives each atom an octet.

B. Ionic compound- a compound composed of positive and negative ions so the charges balance.

1. Formula Unit- the simplest ratio of atoms to represent the formula of an ionic compound.

C. Ionic bond formation: 1. NaCI:

Na Cl

D. When ionic bonds are formed, energy is released.

This energy is called lattice energy. The lattice energy is the energy released when one mole of ionic solid is formed from gaseous ions.

Lattice energy is a measure of the strength of an ionic bond.

E. Comparison of ionic and molecular compounds: 1. Bonding: a. lonic- transfer of electrons creates

the attraction of the ions for each other.

i. The ions are organized into a crystal structure.

ii. Each cation is surrounded by anions.

b. Covalent- the electrons are shared.

2. The attraction between the molecules is less than the attraction between the ions so this creates different properties.

a. Higher melting points, boiling points, and greater hardness for ionic compounds.

b. Ionic compounds are electrolytes (their water solutions conduct an electric current) because they ionize in water.

c. Ionic compounds are usually brittle. A row of ions, when shifted, causes repulsive forces to build up and the layers separate.

10.Polyatomic ions- a. Def": a group of covalently

bonded atoms that act as a single unit forming an ion.

b. Examples: Ammonium ion: NH4 + Sulfate ion: S04-2

Phosphate ion: P04-3

Draw a couple of polyatomic ions in their Lewis structure:

C2O42-


C2O42-

2 C atoms with 4 e-4 O atoms with 6 e-8 + 24 = 32 electrons2- charge - 2 extra electrons32+2 = 34 electrons


C2O42- 34 electrons

Try NH4+

How many total electrons should we have?

Try NH4+

How many total electrons should we have?

8 Because N has 5 e-, 4 H’s have 4 e- and

we have a positive 1 charge, so one electron has been lost.

Try NH4+

IV. Metallic Bonding: A. Def": A bond formed by the

attraction between positive metal ions and the 'sea' of mobile electrons they release.

1. The bond strength varies with nuclear charge and the number of valence electrons available.

2. The strength of the bond can be measured by how much heat is required to vaporize the metal (heat of vaporization).

B. Metallic properties: 1. Conductor- metals conduct heat and

electricity very well because of the free- moving electrons.

2. Malleability- the material can be beaten into different shapes by hammering or rolling. One plane of atoms is able to slide past another with resistance or breaking the bonds.

3. Ductility- the material can be drawn, pulled, or extruded through a small opening to form a wire.

4. Metals have “luster” because as light strikes the surface, the electrons in the “electron sea” absorb and re-emit the light.

5. Electrons are “shared” by all surrounding atoms.

6. In general, the strength of the metallic bond increases moving from left to right on any row of the periodic table.

V. Molecular Geometry: A. VSEPR Theory- Valence Shell

Electron Pair Repulsion 1. The electrostatic repulsion

between valence electron pairs cause these pairs to be spaced as far apart as possible about an atom's nucleus.

2. The VSEPR theory is used to predict molecular geometry.

3. Types of molecular shapes: a. Linear- AB2 CO2

b. Bent- AB2E SnCI2

c. Triangular planer- AB3 BF3

d. Tetrahedral- AB4 CH4

e. Triangular pyramidal- AB3E NH3

f. Bent- AB2E2 H20

g. Triangular bi pyramidal- AB5 PCL5

h. Octahedral- AB6 SF6

B. Hybridization- the mixing of 2 or more atomic orbitals of similar energy, on the same atom, to give new orbitals of equal energy.

1. Methane: CH4

a. Normally carbon has 2 unpaired electrons in its outer shell.

C:

The 1s and 3p orbitals hybridize to form 4 sp3 orbitals. These orbitals are 109.5° apart. This forms a tetrahedral molecule.

b. Hybridization can occur in other ways as well:

- sp; 2 hybrid orbitals. Linear. -sp2; 3 hybrid orbitals. Triangular planer. -sp3; 4 hybrid orbitals. Tetrahedral.

Linear Triangular

C. Intermolecular Forces: 1. Def": The forces of attraction that exist

between molecules. 2. Compared with molecular bonds, the

strength of intermolecular forces is weaker.

3. A polar molecule contains regions of positive charge and a region of negative charge.

1. A molecule of hydrogen chloride is polar because the chlorine attracts the shared electrons more strongly than the hydrogen atom does.

2. Types of intermolecular forces: a. Dipole-dipole forces-

-a dipole is a molecule which has a positive and a negative region. Like a magnet.

-The attraction of the (+) region in one molecule for the (-) region in another molecule creates the bond.

-A dipole can sometimes cause a nonpolar molecule to act as a dipole but this creates a weaker bond.

b. Hydrogen bonding- - The attraction between the hydrogen atom

of one molecule and an unshared pair of electrons on the oxygen atom of another molecule. This occurs in water.

b. London Dispersion forces- - These are intermolecular attractions resulting from the movement of electrons which creates momentary dipoles. They increase in strength as the atomic mass of the atom increases.

ch. 6 chemical bonding chemistry – mrs. demott. key ideas: atoms bond and stick together!! ...

Documents

atoms bond

negative atom sulfur

bond type polar covalent

chlorine slide

electronegative atom

covalent bonding

metallic bond

demott slide