Bonding

Energy and Chemical Bonds

1) Chemical Bond: A force of attraction between atoms in a compound

2) All elements bond for one reason

To acquire an electron configuration of a noble gas (8 valence electrons =stable octet(or 2 valence for He)

3) Bonds and Energy changes

· As bonds are formed, energy is released. This is a(n) exothermic process.

· As bonds are broken, energy is absorbed. This is a(n) endothermic process.

4) Bonds and Stability

· Bonds form to make elements/compounds stable.

· The greater the release of energy, the stronger the bond.

5) Electronegativity (review)

· The measure of attraction an atom has for electrons

· Metals have low electronegativity

· Non metals have high electronegativity

Types of chemical bonds

1) Ionic bonding

· Results from the complete transfer of electrons between metals (that lose electrons) and nonmetals(that gain electrons)

· Metals form positive ions called cations

· (memory hint: has a “t” in the name—looks like a plus sign for positive)

· Non metals form negative ions called anions

Example of ionic bonding:

Na Cl

2-8-1 2-8-7

Result is:

Na ⁺¹ Cl ^-1

2-8 2-8-8

How do you know?

Electronegativity Difference (ED)

· If the difference between the electronegativities (higher minus lower) is 1.7 or more, the bond is ionic.

Example:

Na 0.9 Cl 3.2 3.2 - 0.9 = 2.3

· Exception: the bond is always ionic for a metal hydride:

(group 1 and 2 + hydrogen is always ionic)

Example:

Na 0.9 H 2.1 2.1 - 0.9 = 1.2

2) Covalent Bond: Results from the sharing of electrons between two atoms

Polarity:

· Unequal sharing of electrons

· Each atom attracts electrons by different amounts (like a tug of war)

A. Polar Covalent Bond:

· one atom has a slightly higher affinity for electrons

· electronegativity difference (ED) is between 0.1 and 1.6

example of polar covalent bonds

NH₃ H₂O

N = 3.0 O = 3.4

H = 2.1 H = 2.1

0.9 1.3

B. Non polar Covalent Bond:

· Atoms share the electrons equally

· Electronegativity difference (ED) is 0.0

· Most common in diatomics or ‘triatomics’

· Br₂ H₂ O₂ N₂ Cl₂ I₂ F₂ and O₃

example:

O = 3.4

0.0

C. Coordinate Covalent Bond:

· Atoms share the electrons, but one atom donates both electrons

· **usually the bond is shared with a proton (H+)

example NH₃ + H⁺ → NH₄⁺

3) Metallic Bond:

· These bonds are only found in metals

· Metals do not have a strong attraction for electrons. The electrons are loosely held, so therefore, their electronegativity is very low.

· **These are often described as “positive ions surrounded by a sea of mobile electrons ”

· the positive ions form a strong attraction for the electrons surrounding them causing strong bonds

SUMMARY

E. D. = 0 --------------------------------------------1.7 -------------------------------3.3

Nonpolar Covalent Ionic

Polar Covalent

TYPES OF SOLIDS

1) Ionic solids

· contains ionic bonds

Properties

- regular geometric pattern - lattice

- relatively high melting and boiling points

- in the solid form, a poor conductor of electricity

- in the liquid or aqueous form, a good conductor of electricity

- ionic solids dissolve in water

2) Solids with Covalent Bonds:

Molecular Solids Network Solids

covalently bonded covalent in a 3-D network

1. soft 1. hard, brittle

2. poor conductors 2. poor conductors

3. low melting points 3. high melting points

examples: examples:

NH₃, HCl, H₂O, CH₄ * C , SiC, and SiO₂

(diamond, silicon carbide, silicon dioxide)

3) Metallic Solids

· Contains metallic bonds

· Good conductors of heat and electricity in due to

“sea of mobile electrons”

examples: Cu, Ag, Au

BONDING SUMMARY

The four types of solids:

a) Ionic Solids:

· bonding between metals and nonmetals

· results from the transfer of electrons

· high melting and boiling points

· do NOT conduct electricity in the SOLID state

· conducts only in the liquid (molten) or aqueous phases

· ions are held together by electrostatic attraction

· examples: NaCl (salt), MgO, MgCl₂

b) Metallic Solids:

· bonding between metal atoms

· "sea of mobile electrons"

· good conductors of heat and electricity in all phases

· malleable

· ductile

· shiny luster

· examples: Cu, Fe, Na, Ag

c) Molecular (Covalent) Solids:

· bonding between two or more nonmetals

· results from the sharing of electrons

· poor conductors ( good insulators)

· soft *

· low melting points *

· examples: water (H₂O), methane (CH₄),

ammonia (NH₄), hydrogen (H₂)

d) Network Solids

· Atoms are covalently bonded in a network (3-D)

· hard

· high melting points

· examples: diamond (C), silicon carbide (SiC),

silicon dioxide (SiO₂)

Types of Molecules, Symmetry and Polarity

1) Polar Molecules (dipoles)

· This represents unbalanced charge distribution along the bond

Examples:

H₂O +H-O-H+ NH₃+H-N-H+

H+

NaCl +NaCl-

2. Nonpolar Molecules

· This represents balanced charge distribution

Examples:

CO₂ CH₄

· All diatomics contain nonpolar molecules and nonpolar covalent bonds

Examples:

O₂ F₂

Molecular Attraction

· These are forces between molecules, not to be confused with attractive forces between atoms which are bonds.

1. Dipole – Dipole attraction

· Neighboring polar molecules orient themselves so that oppositely charge regions line up.

Example:

H - Cl H-Cl H-Cl

H-Cl

H H H H H

O O O O O

H H H H H

2. Hydrogen “bonds”

· This happens when hydrogen is bonded to a small highly electronegative element.

· Only happens with F, O and N

· 3 substances are HF, H₂O, and NH₃

· these attractive forces are SO strong they have been called bonds

· hydrogen bonds are the reason that water has such a relatively high boiling point; this also gives insects the ability to “walk on water”

3. van der Waal’s Forces:

· weak intermolecular forces of attraction between individual molecules

a. as molecular mass increases, van der Waal’s forces increase

b. as the distance between molecules increases, van der Waal’s forces decrease.

c. The stronger the van der Waal’s forces, the higher the melting and boiling point.

b.p C₅H₁₂ > b.p. C₄H₁₀ > b.p. C₃H₈

b.p. I₂ > b.p. B₂ > b.p. Cl₂ > b.p. F₂

4. Molecule-Ion Attraction:

· Ions are attracted to the negative and positive ends of water molecules or other polar solvents

· Example: NaCl in water—sodium have a positive charge and is attracted to the negative or oxygen end of the water molecules

Multiple Covalent Bonds

1. Double covalent bond

· 2 pairs of electrons are shared

· Stronger than a single bond

· Shorter than a single bond

· More stable than a single bond

2. Triple Covalent Bond

· 3 pairs of electrons are shared

· stronger than a single or double bond

· shorter than a single or double bond

· more stable than a single or double bond

SUMMARY: Bond strength

Single < double < triple

Write Your Questions Below

is methyl alcohol water soluble and is amyl alcohol water soluble?

February 21, 2012

The hydroxyl group -OH at the end of an alcohol will not form ions in solution (therefore will not conduct electricity) but they are polar and since water is also polar alcohols will be soluable in water.

February 21, 2012 - Replied By Expert