Chemical Bonding Chapter 6. Chemical bond A mutual electrical attraction between the nuclei and...

Chemical Bonding

Chapter 6

Chemistry chapter 6 2

Chemical bond• A mutual electrical attraction between

the nuclei and valence electrons of different atoms that binds them together


Types of bonds• Ionic

• electrons are transferred from one atom to another, forming ions

• Electrical attraction of large numbers of cations and anions holds the compound together

• Covalent – electron pairs are shared between the atoms


Ionic Bonds• Ions get very close together and are attracted

to each other.• Example: NaCl

Na+

Cl-


Covalent Bonds• The atoms share electrons, and their

electron clouds overlap.• Example: carbon monoxide

CO


Electronegativity difference• Can be used to predict how two atoms

will react with each other.• Has no units, because it is a

comparison.• When it is high, the bond is ionic.• When it is low, the bond is covalent.


Ionic – covalent continuum• Real bonds are usually sort of

ionic and sort of covalent.• See figure 6-2 on page 162.• Difference of 1.7 or less is

considered covalent• 50% ionic character or less

• Difference of more than 1.7 is considered ionic• More than 50% ionic character


Nonpolar-covalent bond• Bonding electrons are shared equally• Balanced distribution of electrical

charge• 0 to 5% ionic character• 0 to 0.3 electronegativity difference


Polar• Having uneven charge distribution


Polar-covalent bond• The bonded atoms have unequal

attraction for the shared electrons• 5 to 50% ionic character• 0.3 to 1.7 electronegativity difference


Polar and nonpolar bonds

• Electron density is greater around Cl than H because it has greater electronegativity

• d+ and d- indicate partial charge• Electron cloud drawings

nonpolar polar

-d+d


Comparing bond types• Comparison between ionic and covalent

bonding


Discuss• Page 163

• Sample problem• Practice problem• Section review 1 and 2


Molecule• A neutral group of atoms that are held

together by covalent bonds• A single molecule is capable of existing

on its own• May consist of atoms of the same

element or different elements


Molecular compound• Compound whose simplest units are

molecules


Chemical formula• Uses atomic symbols and subscripts to

indicate the relative numbers of atoms of each kind in a chemical compound.


Molecular formula• Shows the types and numbers of atoms

in a single molecule of a molecular compound

• H2O

• O2

• C12H22O11


Diatomic molecule• Contains only two atoms


Bonding• Bonding makes atoms be at lower

potential energy levels, which is favored by nature.

• Attractive forces (nucleus-electron) are balanced by repulsive forces (nucleus-nucleus and electron-electron).


Bond length• Average distance between bonded

nuclei at their lowest potential energy


Bond energy• The amount of energy that must be

added to break a chemical bond and form neutral isolated atoms.

• The same amount of energy that the atoms had to release when they bonded.


Octet rule• Chemical compounds tend to form so

that each atom, by gaining, losing or sharing electrons, has an octet of electrons in its highest occupied energy level

• The outermost s and p sublevels want to be full


Exceptions to octet rule• Hydrogen only needs two electrons, not

eight• Boron tends to only be surrounded by

six electrons.• Some atoms can have more than eight

when combining with fluorine, oxygen, and chlorine. This expanded valence also includes the d orbitals.


Electron dot diagrams• Usually only the valence electrons are

involved in chemical reactions.• Electron dot diagrams allow us to draw

these electrons around the symbol for an element.


Drawing electron dot diagrams

1. The element’s symbol represents the nucleus and all the electrons in the inner energy levels.

2. Write the noble gas configuration of the element. Add the superscripts of the s and p orbitals to get the number of valence electrons.


3. Draw dots on the sides to represent electrons.


Examples• Examples


Molecules• Can be represented by dot diagrams• Example: HCl• Shared pairs can be replaced by long

dashes


Lone pair• Unshared pair of electrons• Not involved in bonding


Discuss• Read sample problem 6-2 on page 170


Lewis diagram• Formulas in which atomic symbols

represent the nuclei and inner-shell electrons• Dot diagrams with dots or dashes for

shared electrons


Structural diagrams• Leaves out the unshared pairs• Indicates the kind, number,

arrangement, and bonds of atoms


Single bond• Covalent bond produced when one pair

of electrons is shared


Discuss• Sample problem 6-3 on pages 171-172


Double bond• Covalent bond produced when two

pairs of electrons are shared• Example: CO2


Triple bond• Covalent bond formed when three pairs

of electrons are shared• Example: N2


Multiple bonds• Double or triple bonds• Have higher bond energies• Have shorter bond lengths


Resonance structures• Bonding in molecules or ions that

cannot be correctly represented by a Lewis structure

• A double-headed arrow is placed between the two possibilities

• Example: ozone


Discuss• Covalent Bonds• More Covalent Bonds• Sample problem 6-4 on page 174


Examples• Draw a Lewis Structure and the structural

formula for each of the following: • H2O

• CH4

• CH4O

• O2

• C2H2

• C2H4


Discuss• Section Review on page 175


Ionic compound• Composed of positive and negative ions that

are combined so that the numbers of positive and negative charges are equal

• Usually crystalline solids• Not composed of independent units like

molecules• Ionic bonding• More ionic bonding


Chemical formula• For ionic compounds

• Shows the simplest ratio of the ions that gives no net charge


Formula unit• The simplest collection of atoms from

which an ionic compound’s formula can be established.

• For sodium chloride, NaCl• The simplest way to show a one to one

ratio


Ratio of ions• Depends on the charges of the ions

combined.• Na+ and Cl- form NaCl• Mg2+ and F- form MgF2


Forming ionic compounds• Examples:• NaCl• MgF2


Crystal lattice• Ions minimize their potential energy by

forming an orderly arrangement


Lattice energy• The energy released when one mole of

an ionic crystalline compound is formed from gaseous ions.

• Used to compare bond strengths


Molecular vs. ionic compounds

• Forces between molecules are weaker than ionic bonding forces

• Molecular compounds have lower melting points

• Ionic compounds are hard but brittle.• Ionic compounds conduct electricity

when melted or dissolved in solution


Polyatomic ions• A charged group of covalently bonded

atoms.• Form ionic bonds just like other ions.• Examples

• Ammonium ion NH4+

• Sulfate ion SO42-


Discuss• Use electron-dot notation to

demonstrate the formation of ionic compounds involving• Li and Cl• Ca and I

• What basic unit are molecular compounds composed of? Ionic compounds?


Metals• High electrical conductivity.• Metals form crystals and their orbitals

overlap.• All the atoms have the same attraction

for their electrons, so the electrons can move easily from one atom to another.


Valence electrons• Metals have very few• Their p sublevels are empty• Transition metals also have many

vacant d sublevels


Metallic bonding• Delocalized electrons – don’t belong to

one specific atom• Sea of electrons formed around the

metal atoms• Metallic bonding results from the

attraction between metal atoms and the surrounding sea of electrons


Metallic properties• High electrical and thermal conductivity• Ability to absorb a wide range of light

frequencies• Electrons enter excited states and then

return to ground states• This makes metals look shiny


Metallic properties• Metals bond the same way in all

directions• Their layers can slide past each other

without breaking• Malleabiltiy - the ability to be hammered or

beaten into thin sheets• Ductility – the ability to be drawn, pulled or

extruded into a thin wire


Metallic bond strength• When there is a bigger nucleus and

more electrons, the bond is stronger.• Stronger bonds have higher heats of

vaporization


Discuss• Describe the electron-sea model of

metallic bonding.• What is the relationship between

metallic bond strength and heat of vaporization?

• Explain why most metals are malleable and ductile but ionic crystals are not.


Molecular polarity• Uneven distribution of molecular charge• Determined by the polarity of each bond

and the geometry of the molecule


Molecular geometry• There are two equally successful

theories• One accounts for molecular bond

angles• The other describes the orbitals that

contain the valence electrons


VSEPR theory• Valence-shell, electron-pair repulsion• The repulsion between the sets of

valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible

• To predict geometry, one must consider the location of all electron pairs around the bonded atoms.


Linear molecules• A is the central atom• B is an atom bonded to it• Formulas like AB2 form linear

molecules, with the B atoms on opposite sides of the A atom


Trigonal Planar

• AB3

• A is in the center.• The 3 Bs make a triangle around it.• The molecule is flat (planar)


Tetrahedral

• AB4

• A pyramid with A at the center and a B at each corner


When there are lone pairs• The repel the other electrons just like

shared pairs do• Use an E to represent the unshared

pairs


Bent or angular

• AB2E

• Like trigonal planar, but without the top of the triangle


Trigonal pyramidal• AB3E

• Like tetrahedral, but without the top of the pyramid

• A is at the center, the Bs are the base, and E is at the top


Bent or angular

• AB2E2

• Like a tetrahedral, but only the atoms (not the unshared pairs) determine the shape


Trigonal bipyramidal

• AB5

• Like two pyramids on top of each other


Octahedral

• AB6

• 8 sided shape


VSEPR• Double and triple bonds are treated the

same way as single bonds.• Use Lewis structures and table 6-5 on

page 186 to predict molecule shapes


Hybridization• Explains how orbitals get rearranged

when atoms form covalent bonds• Describes how atoms are bonded


Hybrid orbitals• Orbitals of equal energy produced by

the combining of two or more orbitals on the same atom

• The number of hybrid orbitals equals the number of orbitals they were made from


Example: Methane

• CH4

• To get 4 equal orbitals, the s and p orbitals combine.

• They form 4 sp3 orbitals


Other hybrid orbitals


Intermolecular forces• The forces of attraction between

molecules• Usually weaker than covalent, ionic,

and metallic bonds• Can be measured by boiling point

• Higher boiling point means stronger intermolecular forces


Dipole• Equal but opposite charges separated

by a short distance• Direction is from the positive pole to the

negative pole

ClH


Dipole-dipole forces• Forces of attraction between polar

molecules• Strongest intermolecular forces• Only act over short distances


More than two atoms• Polarity of the molecule depends on the

polarity and orientation of each bond


Induced dipoles• A polar molecule can induce a dipole in

a nonpolar molecule by attracting its electrons.

• It is temporary


Hydrogen bonding• A hydrogen atom in a compound with a

very electronegative atom is attracted to an unshared pair of electrons in a nearby molecule.

• Examples:• Water (H2O)• Hydrogen fluoride (HF)• Ammonia (NH3)


Hydrogen bonding• Causes higher boiling points


Polarity and Hydrogen Bonding

• Polarity and Hydrogen Bonding


London dispersion forces• The intermolecular attractions caused

by the constant motion of electrons and the creation of instantaneous dipoles

• Act between all atoms and molecules• The only intermolecular forces among

noble gas atoms and nonpolar molecules• Low boiling points


discuss• What is the difference between

intramolecular forces and intermolecular forces?

• What is a dipole? In what kind of molecule are dipoles common?

• What is hydrogen bonding?• What are London dispersion forces?• Molecular Geometry review

Chemical Bonding Chapter 6. Chemical bond A mutual electrical attraction between the nuclei and...

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