Chapter 6.2 and 6.5 Covalent Compounds. Covalent Bonds Sharing Electrons –Covalent bonds form when atoms share one or more pairs of electrons nucleus

Chapter 6.2 and 6.5Chapter 6.2 and 6.5Chapter 6.2 and 6.5Chapter 6.2 and 6.5

Covalent CompoundsCovalent Compounds

Covalent Bonds• Sharing Electrons

– Covalent bonds form when atoms share one or more pairs of electrons• nucleus of each atom is attracted to

electron cloud of other atom• neither atom removes an electron from

the other

Covalent Bonding

Covalent Bonds• Sharing Electrons

– Covalent bonds• space where electrons move is called

molecular orbital– made when atomic orbitals overlap

Molecules

Covalent Bonds

• Energy and Stability– Noble gases are stable (full octet)

(low P.E.)– Other elements are not stable (high

P.E.)• covalent bonding decreases potential

energy because each atom achieves electron configuration like noble gas

Covalent Bonds

• Energy and Stability•because P.E. decreases when atoms bond, energy is released –i.e., atoms lose P.E. when they bond

–loss of P.E. implies higher stability

Covalent Bonds• Energy and Stability

•potential energy determines bond length

–at minimum P.E., distance between two bonded atoms is called bond length

»bonded atoms vibrate»therefore, bond length is an

average length

Covalent Bonds• Energy and Stability

•bonds vary in strength–bond energy is the amount of

energy required to break the bonds in 1 mol of a chemical compound

–bond energy predicts reactivity–bond energy is equal to loss of

P.E. during formation

Bond Energies and Lengths

Bonds and Energy• Single bonds are the longest bonds

with the least bond energy• Double bonds are shorter, stronger

and have intermediate bond energy• Triple bonds are the shortest,

strongest, and have the highest bond energy

Covalent Bonds• Electronegativity

– Atoms share electrons equally or unequally•nonpolar covalent bond:

bonding electrons shared equally•polar covalent bond: shared

electrons more likely to be found around more electronegative atom


– Atoms share electrons equally or unequally•difference in electronegativity can be used to predict type of bond (but boundaries are arbitrary)

Bond Types

0.3 1.7

Practice: Calculate the bond type

• N and H• F and F• Ca and Cl• C and O

• Polar • Non-polar• Ionic• Polar


– Polar molecules have positive and negative ends• such molecules called dipoles• δ means partial in math and science• positive end—δ+

• negative end—δ-

• example: Hδ+Fδ-

Electronegativity Difference for Hydrogen

Halides


– Polarity is related to bond strength•greater electronegativity

means •greater polarity

means •greater bond strength


– Bond type determines properties of substances•metallic bonds: electrons can move

from one atom to another—good conductors

•ionic bonds: hard and difficult to break apart

•covalent bonds: low melting/boiling points

Properties of Substances with Different Types of Bonds

Polarity of Molecules

• Two atoms: bond polarity is the molecular polarity

• More than 2 atoms: the geometry of the molecule must be considered

• If the bonds are non-polar, the molecular is non-polar

• Some molecules with polar bonds • can be non-polar

More • Sometimes the partial charges

cancel each other out because they are directly opposite each other

• Consider CO2 and CCl4• The symmetrical distribution of the

bonds leads to cancellation of the charges

Polar and Non-polar Molecules

Drawing and Naming• Lewis Electron-Dot Structures

– Lewis structures represent valence electrons with dots•position of electrons is symbolic

(not literal)•shows only the valence electrons

of an atom•dots around atomic symbol

represent electrons

Lewis Structures of Second-Period Elements

Drawing and Naming

• Lewis Electron-Dot Structures– Cl2

– HCl


– Drawing

–1. Gather information»draw Lewis structure for each

atom in compound; place one electron on each side before pairing

»determine total number of valence electrons


– Drawing•2. Arrange atoms

–arrange structure to show bonding

–halogens and hydrogen usually make one bond at end of molecule

–carbon usually in center


– Drawing•3. Distribute the dots so that each

atom satisfies octet rule (except H, Be, B)

•4. Draw the bonds as long dashes•5. Verify the structure by counting

number of valence electrons


– Polyatomic Ions• use brackets [] to show overall charge• example:


– Multiple Bonds• sharing two pairs of electrons is a

double bond• sharing three pairs of electrons makes

triple bonds• example:


– Resonance Structures•sometimes, multiple structures are possible

•show all possibilities•example:

Drawing and Naming• Naming Covalent Compounds

– First name: name of first element in formula• usually least electronegative• requires a prefix if more than one of them

– Second name: ends in –ide• requires a prefix if more than one of them

Naming Covalent Compounds

Practice Naming1. antimony tribromide

2. hexaboron silicide

3. chlorine dioxide

4. iodine pentafluoride

5. dinitrogen trioxide

6. ammonia

More Naming Practice1. P4S5

2. PI33. NO

4. N2F4

5. CO2

6. H2O

Molecular Shapes

• Determining Molecular Shapes– Three-dimensional shape helps

determine physical and chemical properties

– valence shell electron pair repulsion (VSEPR) theory predicts molecular shapes• based on idea that electrons repel one

another

Molecular Shapes

• Lewis structures show which atoms areconnected where, and by how many

bonds, but they don't properly show 3-D shapes of molecules.

To find the actual shape of a molecule, first draw the Lewis structure, and then use VSEPR Theory.

MOLECULAR MOLECULAR GEOMETRYGEOMETRY

VSEPRVSEPR • VValence alence SShell hell EElectron lectron PPair air

RRepulsion theory.epulsion theory.• Most important factor in determining Most important factor in determining

geometry is relative repulsion between geometry is relative repulsion between electron pairs.electron pairs.

MOLECULAR MOLECULAR GEOMETRYGEOMETRY

Molecule Molecule adopts the adopts the shape that shape that minimizes minimizes the electron the electron pair pair repulsions.repulsions.

VSEPR Rules• To apply VSEPR theory:• 1: Draw the Lewis structure of the

molecule and identify the central atom

• 2: Count the number of electron charge clouds (lone and bonding pairs) surrounding the central atom.

• 3: Predict molecular shape by assuming that clouds orient so they are as far away from one another as possible

VSEPR Shape Predictor Table -

117°

VSEPR Shape Predictor Table -

Bond Angles

• Lone-pairs of electrons behave as if they are slightly bigger than bonded electron pairs and act to distort the geometry about the atomic center so that bond angles are slightly smaller than expected:

Bond Angles

• Methane, CH4, has a perfect tetrahedral bond angle of 109.5°, while the H-N-H bond angle of ammonia, NH3, is slightly less at 107°, trigonal pyramidal

•

Bond Angles• The oxygen of water

has two bonded electron pairs and two nonbonded "lone" electron pairs giving a total VSEPR coordination number of 4.

• But the geometry is defined by the relationship between the H-O-H atoms and water is said to be "bent" or "angular" shape of 105°.

105°

Molecular Shapes• Determining Molecular Shapes and

angles.– Let’s try some.

• CO

• CO2

• BF3

• CH4

• CBr4

• PCl3

Simple Shapes

Bond angle for linear is 180°

Trigonal Planar

Bond angle of 120°

TetrahedralBond angle of 109.5°

Bent

Bond angle of 105°

Documents

Chapter 6.2 and 6.5 Covalent Compounds. Covalent Bonds Sharing Electrons –Covalent bonds form when atoms share one or more pairs of electrons nucleus