Pre AP Chemistry

Chapter 6

“Chemical Bonding”

Introduction to Chemical Bonding

• Chemical bond – a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

• Why are most atoms bonded together?

Types of Bonds: IONIC bonds• Ionic bonding – the electrical attraction between anions

and cations.

• Electrons are transferred between atoms.

• Ionic bonds often form between metals on the left of the periodic table and nonmetals on the right side of the table

Types of Bonds: METALLIC BONDS

• Metallic bonding – the attraction between metal ions and the surrounding sea of electrons

• Freedom of electrons to move in a network of metal atoms results in metallic properties, ability to conduct heat and electricity, and to reflect light.

Metallic bonds, cont.

• Two important properties related to metallic bonding are:

• Malleability- metals can be hammered into sheets

• Ductility- metals can be pulled to form wires.

Types of Bonds: COVALENT BONDS

• Covalent bonding – the result of sharing electron pairs

• Covalent bonds with uneven electron sharing are polar.

• Covalent bonds with even electron sharing are nonpolar.

Bond Polarity• electronegativity - the tendency for an atom to

attract electrons to itself when bonding with other atoms

• The difference between electronegativity can be used to indicate bond type.

• Similar atoms tend to from nonpolar covalent bonds, while atoms with greater differences in their electronegativities form polar covalent or ionic bonds.

Bond Type• To calculate bond type

– 1. Look up the electronegativity of the atoms in question using the periodic table or chart in the textbook.

– 2. Find the difference between the electronegativity values of the atoms. (subtract)

– 3. Use the chart to classify the bond as• nonpolar covalent (0.0 - 0.49)• polar covalent (0.5 – 1.9)• ionic (>2)

Determine the bond type of each of the following.

• H2

• H2O

• NaCl

• Assignment: Page 177, 1 - 6

Why do atoms share electrons?• Hydrogen exists as a molecule because the proton in one

hydrogen atom attracts the electron in the other and vice versa.

• The electrons repel each other but this repulsive force is far less than the attractive forces.

• Several elements from diatomic molecules: H2, F2, Cl2, Br2, I2, O2, and N2.

DIATOMIC ELEMENTS• NAME DOT FORMULA MOLECULAR FORMULA

• hydrogen H2

• nitrogen N2

• oxygen O2

• fluorine F2

• chlorine Cl2

• bromine Br2

• iodine I2

• astatine At2

Covalent Bonding and Molecular Compounds

•The particle which results from the covalent bonding of two or more atoms is a molecule.•There are eight elements in which two atoms bond forming a diatomic molecule. They do not normally exist as single atoms. They are referred to as diatomic elements.•diatomic elements - two identical atoms bonded covalently•molecular compound - a chemical compound whose simplest formula units are molecules

Formulas Represent Compounds

• chemical formula - a shorthand method of using atomic symbols and subscripts to represent the composition of a substance

• molecular formula - a formula indicating the composition of a molecule Ex) C6H12O6

represents glucose

Formulas, cont.

• dot formula - a formula using dot notation to indicate valence electrons

• formula unit – represents the composition of an ionic compound, empirical formula

Ex) NaCl, CaBr2

Lewis Structures• Lewis structures are

formulas in which symbols represent nuclei and inner shell electrons and dots represent valence electrons.

• The Lewis structure of water would be similar to the dot formula.

Lewis dot structures for compounds

Steps:

1)Find the total number of valence e- for each element in the compound

2)Put the least electronegative atom in the center (H+ is always on the outside!)

Lewis structures, cont.

3) Complete octets on the outside atoms

4) If the central atom does not have an octet, move e- from outer atoms to form a double or triple bond Ex) CO

Polyatomic Ions

• A polyatomic ion consists of two or more atoms bonded covalently which has a net charge.

• **See the back of your periodic table for a list of common polyatomic ions

• Show the Lewis structure for SO4-2

• Show the Lewis structure for H2 SO4.

Examples

• Show the Lewis structures for– MgBr2

– Na2O

– H3PO4

– CH3Cl

– Al2S3

Resonance

• Resonance is an attempt to describe bond structure based on data collected about bond length.

• Example: ozone, O3

• Assignment: Page 189, 1-5

Ionic Bonding and Ionic Compounds

• Ionic bonding results from electron transfer.

• Ion – an atom or group of atoms that has an unbalanced electrostatic charge

• Crystal – the particle resulting from ionic bonding

• Most ionic compounds are solids.

• Salts are examples of ionic compounds.

Assignment: Pg. 194 #1-5

VSEPR Theory

• VSEPR - valence shell electron pair repulsion – the valence electron pairs repel each other which moves bonded atoms to an equilibrium position

• VSEPR accounts for the bent shape of the water molecule.

Water

Molecular Shape

• A quick indicator of molecular shape is the number of atoms in a molecule.

• The un-bonded electrons must be taken into account to get the exact shape.

• What do you think is the shape of:• H2

• H2O• NH3

• CH4

Molecular Type• Molecular type is either polar or nonpolar. These

are not to be confused with bond type.

• Polar molecule – a molecule which lacks symmetry

• Nonpolar molecule – a molecule which has symmetry

• Which shapes do you expect to be polar and which do you expect to be nonpolar?

• Page 207, 1-6