CHEMICAL BONDING Unit 4

BONDING

Chemical Bond: mutual electrical attraction between nuclei and valence electrons of different atoms

Type of bond depends on electron configuration and electronegativity

WHY DO ATOMS BOND?

TO BECOME STABLE!!Key: 8 valence electrons

TYPES OF BONDINGIonic: an electrostatic force of attraction between positive and negative ions; ions for when atoms lose/gain electrons (electrons are transferred from metal to nonmetal)

Covalent: sharing of electron pairs between atomspolar unequal sharing of electronsnonpolar equal sharing of electrons

Metallic: an attraction between metal atoms and outer mobile electrons

COVALENT BONDING

COVALENT BONDING

Occurs between NONMETALS

Nonmetals have high electronegativities and want to gain more electrons. They cannot lose valence electrons. In order to bond, nonmetals must then SHARE valence electrons.

The atoms share enough electrons to obtain 8 valence electrons (including the shared electrons)

TYPES OF COVALENT BONDS

Equal sharing of electron between atoms in a compound, no positive or negative poles exist: NONPOLAR!

Unequal sharing of electrons, positive and negative ends exist: POLAR!

DETERMINE BOND FROM ELECTRONEGATIVITY DIFFERENCES

0 0.3 1.7 3.3

Covalent Ionic

Non-polar Polar

DETERMINE THE TYPE OF BONDUse the following electronegativity values to determine the bond character of each:

H – S (2.2, 2.6)

S – Cl (2.6,3.2)

Rb – Se (0.8,2.6)

C – I (2.6,2.7)

Cs – S (0.8, 2.6)

O – O (3.4,3.4)

F – I (4.0,2.7)

Ca – N (1.0,3.0)

OTHER COVALENT BONDSCovalent network bonding: 3-D network of covalently bonded atoms macromolecules

Very high melting points, very hard solids

Ex: diamond, graphite, quartz

COORDINATE COVALENT BOND

DEF: When both of the shared electrons in a covalent bond come from the same atom

Coordinate Covalent bond will form here

TERMS TO KNOW!Molecular compound: neutral compound consisting of nonmetals covalently bonded in which the electrons are shared

Molecule: smallest representative unit of a molecular compound, can exist independently

DIATOMIC MOLECULES (MEMORIZE THESE!)

Diatomic molecules: molecules consisting of two atoms of the same element. They are ALWAYS found in pairs unless in a compound with something else

Ex: H2 O2 F2 Br2 I2 N2 Cl2“HOFBrINCl”

**Form a “7” starting at atomic number (Z) 7 and include H

DEFINITIONS

Bond length: average distance between nuclei of two bonded atoms (sum of atomic radii)

Bond angle: angle between two bonds in a molecule

Bond energy: energy needed to break a bond and form neutral atoms

FACTS

As bond length increases, bond energy decreases.

As bond length increases, bond strength decreases.

**Think about it!!!**

What relationship do each of the above have?

APPLY THAT INFORMATION!

Question: What can you say about the bond strength and bond length of these 2 compounds?

POLYATOMIC IONS

DEF: covalently bonded group of atoms that have an overall charge.

Examples: Nitrate (NO3-), Hydroxide (OH-), Ammonium (NH4

+)

PROPERTIES OF MOLECULES

Low melting & boiling pointsBrittle, dull solids or gasesPoor conductors of heat and electricity

WARM UP – 2/25

You are given a covalent compound. Answer the following:

Will it conduct electricity

Is it malleable

Could you melt it in your kitchen?

What is it doing with its electrons?

Are there any metallic atoms in it?

OCTET RULEDefinition: atoms tend to gain, lose or share electrons so that they have eight valence electrons

exceptions Hydrogen (2e-), Beryllium (4e-), Boron (6e-)

STEPS DRAWING LEWIS STRUCTURES1. Total the number of valence electrons in a molecule.

a) If there is an overall positive charge, subtract it from the number of valence electrons.

b) If there is an overall negative charge, add it to the number of valence electrons.

2. Draw the atom in the molecule connected with a single bond. If carbon is present, it is the center atom. Otherwise the atom in the singlemost form is the center atom.

STEPS DRAWING LEWIS STRUCTURES3. Add electrons so that all of the atoms are happy. (Octet

Rule)4. Compare the number of electrons you have drawn to the

number of valence electrons.a) If you have drawn too many electrons, make double or

triple bonds. In order to make a multiple bond, you must take away a pair of electrons from the atom on each side of the bond.

b) If you have not drawn enough electrons, add extra electrons to the central atom.

LET’S TRY THESE EXAMPLES!

NH3

CH4

H2O

BF3

CO2

MOLECULAR GEOMETRY a.k.a. Molecular Shape

VSEPR THEORY

Valence Shell Electron Pair Repulsion electron pairs (clouds) spread as far apart as possible to minimize repulsive forces

IMPORTANT VSEPR TERMS

Lone Pair: A pair of electrons that do NOT create a chemical bond

Shared Pair: A pair of electrons that make a chemical bond between elements Lone Pair of

electrons in an electron cloud

Shared Pair of electrons

VSEPR Shape # Bonds to Central Atom

# Lone Pairs to Central

Atom

Examples

Linear 2 atoms together or 2

bonds to central atom

0 H-ClO = C = O

Bent 22

21

Trigonal Planar

3 0

Trigonal Pyramidal

3 1

Tetrahedral 4 0

LET’S TRY IT - PRACTICE PROBLEMS!

NH3

CH4

H2O

BF3

CO2

POLARITY

Molecular polarity is the distribution of molecular charge (even or uneven)

Molecular polarity depends on:SymmetryMolecular shapeLone pairs

Molecular polarity influences intermolecular forces.

POLAR VS NON POLAR STRUCTURES

Water is polar

Methane is non polarHydrogen Fluoride is polar

MOLECULAR POLARITY

Nonpolar molecules are usually symmetrical. (about their bonds)

Polar molecules are usually nonsymmetrical. (about their bonds)

Polar molecules are called dipoles. (positive and negative ends)

**As soon as you have a lone pair on your central atom, your molecule is polar.**

Two shapes are always polar: bent and trigonal pyramidal.

LET’S TRY IT - PRACTICE PROBLEMS!

NH3

CH4

H2O

BF3

CO2

INTERMOLECULAR FORCES

INTERMOLECULAR FORCES

Definition: weak forces of attraction between molecules

TypesDipole-dipole forcesHydrogen bondingLondon Dispersion forces

INTERMOLECULAR FORCES

DIPOLE-DIPOLE FORCES

Exists between polar molecules

Results in higher melting point and boiling point than expected

Substances exist mostly as solids or liquids

HYDROGEN BONDING

DEF: Special dipole-dipole force occurring when hydrogen on one molecule is attracted to N, O, F of another molecule

Ex: gives water its unusual properties, ice floats in liquid water, higher melting/boiling point, surface tension

HYDROGEN BONDING

http://www.chemguide.co.uk/atoms/bonding/hbond.html

LONDON DISPERSION

DEF: instantaneous dipole due to a shift in electrons in the molecule

Occurs between Nonpolar-Nonpolar

Strength increases with an increase in number of electrons

Substances with LDF only have low melting points/boiling points and are mostly gases

Ex: diatomic molecules, noble gases

DETERMINE THE IMF(S) PRESENT

1) H2O2) SCl23) PF3

4) CF4

Helpful Hint: Draw the molecular structure and determine if the bond is polar or non polar.

DispersionForces

Dipole-Dipole

Hydrogen Bonding

HOW DO THESE PROPERTIES RELATE TO BOND TYPE OR INTERMOLECULAR FORCES?Stronger bonds/imf’s higher melting point/boiling point

Weaker bonds/imf’s lower melting point/boiling point

Solubility “likes dissolves like”Polar dissolves polar, nonpolar dissolves nonpolar

METALLIC BONDING

METALLIC BONDING

Definition: an attraction between metal atoms and outer mobile electrons (delocalized electrons)

METALLIC BONDINGDEF: an attraction of metallic atoms for “delocalized electrons;” electrons roam traveling freely from one atom to another

positive ions in a sea of electrons

http://www.revisioncentre.co.uk/gcse/chemistry/bonding.html

http://www.ndt-ed.org/EducationResources/CommunityCollege/Materials/Structure/metallic.htm

FUN FACTSThe strength of a metallic bond is determined by number of outer electrons softer metals can be combined with harder metals. (forming alloys)The strongest metals are the transition metals.Some “d” electrons are delocalized and involved in bonding

PROPERTIES OF METAL CRYSTALS

Good conductors of heat and electricityShiny (luster)Hard, metallic crystalsMalleableDuctileHigh melting point

WARM UP – 2/28

Tell me if the following compounds are covalent, ionic, or metallic?

CH4 SO3

NaCl FeO2

K2SO4 Ag

Fe NH3

IONIC BONDING

IONIC BONDING

Definition: an electrostatic force of attraction between positive and negative ions; ions form when atoms lose/gain electrons

IONIC BONDINGIonic compounds are formed by the force of attraction between positive and negative ions. It involves a transfer of electrons from metal to nonmetal

# (+) = # (-) (metals – nonmetals)

Formula unit: simplest combining ratio of ions in a compound does not exist independently

http://www.visionlearning.com/library/module_viewer.php?mid=55

DETERMINING THE FORMULA UNIT

Key point: overall charge must be neutral!Potassium and fluorine

Calcium and iodine

Magnesium and nitrogen

LET’S TRY IT!

Calcium and nitrate

Sodium and phosphate

STRUCTURES - DOT DIAGRAM FOR NaCl

STRUCTURES - DOT DIAGRAM FOR CaBr2

IONIC COMPOUNDSCrystal lattice: 3-D crystal structure (arrangement of ions)

Lattice energy: energy released when one mole of ionic crystalline compound is formed from gaseous ions

http://www.visionlearning.com/library/module_viewer.php?mid=55

PROPERTIES OF IONIC COMPOUNDS

High melting point Most soluble in waterCrystalline solidGood conductors in liquid or aqueous stateHard solids but will fracture

REVIEW ACTIVITY Ionic or covalent?

LITHIUM BONDING WITH NITROGEN

TWO NONMETALS BONDING

FLUORINE BONDING WITH MAGNESIUM

METAL BONDING WITH NONMETAL

Cs2S

PH3

FBr3

CALCIUM AND PHOSPHATE

LOW MELTING POINT

HF

METALS ONLY

SHARING ELECTRONS

DELOCALIZED ELECTRONS

CCl4

H2S

Na3N

WARM UP – 3/1

Draw the following structures. If applicable, determine the shape and the strongest IMF present.

SiS2

Mg2N3

H2S

WARM UP – 3/4

Tell me the three MOST DIFFICULT things from this unit. (I said MOST so you can’t say “everything”)