CHAPTER 6 IONIC AND COVALENT BONDS, NAMING, AND MOLECULAR GEOMETRY

CHAPTER 6CHAPTER 6IONIC AND COVALENT BONDS, NAMING, AND MOLECULAR

GEOMETRY

IONIC COMPOUNDSIONIC COMPOUNDS

What happens to electrons?e- are transferred from the metal to the

non-metal, +/- ions are formed and held together by electrostatic attraction.

What is the electronegativity difference?>1.7 Ex.

What type of elements are involved?Metals and non-metals or polyatomic ions

IONIC COMPOUNDSIONIC COMPOUNDS

What is the simplest unit of the compound called?

Formula unit. Many formula units combine to form a crystal lattice.

How is strength of the bond measured?Lattice energy – the energy required or

released when an ionic compound is formed from 2 ions in their gaseous state.

Higher charges and smaller ions lead to higher lattice energies.

COVALENT COMPOUNDSCOVALENT COMPOUNDS

What happens to electrons?Electrons are shared evenly – pure covalentElectrons are shared, but are pulled towards

the more electronegative atom – polar covalent.

What is the electronegativity difference?Pure – less than 0.3Polar – between 0.3 and 1.7

What type of elements are involved?2 non-metals together


What is the simplest unit of the compound called?

Non-polar molecule orPolar molecule

How is strength of the bond measured?Bond energy – the energy required to break a

covalent bond.SingleDoubleTriple Shorter/ strongerWeaker/ longer

IONIC PROPERTIESCOVALENT PROPERTIES

Very strong bond and Intermolecular forces due to electrostatic attraction

High melting point and boiling point

When (l) or dissolved in water (aq) ionic compounds will conduct electricity

Hard and brittleForm a crystal lattice

with unique crystal shapes

Weaker bonds/ e are shared

Low melting and boiling point

Soft (usually l or g)Non-conductorsLow intermolecular

forces ◦ H-bonding, dipole-

dipole in polar covalent◦ London dipersion forces

in non-polar covalent.

ION FORMATION REVIEWION FORMATION REVIEW

Why do ions form?

Atoms become stable by gaining or losing

VALENCE electrons to achieve a full outer ENERGY level.

◦ Losing electrons produces a POSITIVE charge and the

ion is called a CATION.

◦ Gaining electrons produces a NEGATIVE charge and the ion is called an ANION.

Normally, atoms will try to achieve a full octet

and noble gas electron configuration to become stable.

Ex. Li loses 1 electron and forms Li+ which has the same electron configuration as Helium.

[He]2s1 [He] Ex. S gains 2 electrons and forms S–2

which has the same configuration as Argon.

[Ne] 3s23p4 [Ne]3s23p6 [Ar]

Exceptions to the octet rule:

Hydrogen and He only need 2 (not 8) to be stable, since they only include an S orbital.

TRANSITION METALS do not have to

achieve noble gas configuration to become more stable.

Ex. Fe can have charges of +2, +3, or +6

Energy level diagram: 4s___ 3d___ ___ ___ ___ ___

Lose the 2 electrons in the 4 s orbital: +2Lose the 2 electrons in the 4s and 1 from 3d:

+3Lose the 6 electrons in the 3d: +6

Ions have different properties than the parent atom. ◦ Recall, the size of the ion changes.

◦ Cations are smaller than the parent atom.

Anions are larger than the parent atom.

◦ Ions are more stable than the parent atom. Ex. Once sodium loses an electron, it is no

longer reactive when it becomes +1. We can eat Na+ in NaCl, but we could never eat

Na (s) – it explodes in water!!!

Polyatomic ions

Covalently bonded group of atoms with an overall charge.

Ex. NH4+

FORMATION of a CRYSTAL LATTICEFORMATION of a CRYSTAL LATTICE

Label ionization energy step.Label electron affinity step.Label the bond energy (between Cl-Cl).Label the crystal lattice energy – released

when sodium and chlorine ions come together.

IONIC FORMULAS and NAMING IONIC FORMULAS and NAMING IONIC COMPOUNDSIONIC COMPOUNDS

Ionic compounds must have a net charge equal to ZERO.

SUBSCRIPTS are used to balance positive and negative charges. Use parenthesis when a POLYATOMIC ION has a subscript.

A short cut method for balancing charges is the CRISS CROSS method.

1. Write the cation and its charge followed by the anion and its charge.

Ex. Fe+2O-2 Ba+2 PO4-3 Sn+4O-2

2. Criss cross the charges to form subscripts on the opposite ion.

Ex. FeO Ba3(PO4)2 SnO2

More examples:Determine the formula of the compound

that forms when:Sodium bonds with phosphate ion

Magnesium bonds with nitrogen

Aluminum bonds with hydroxide

RULES for NAMING IONIC RULES for NAMING IONIC COMPOUNDSCOMPOUNDS

Name the CATION first.a.Use the name of the element from the

periodic table.

b. If the cation is a metal with multiple charges, you must indicate the charge on the ion using ROMAN NUMERALS in parenthesis following the metal name.

Ex.

2. Name the ANION seconda.Use the name of the anion from the

periodic table, drop the ending and add –IDE

Ex. O-2 becomes oxide P-3 becomes phosphide

N-3 becomes nitride S-2 becomes sulfide If a polyatomic ion is present, use thename given on the polyatomic ion sheet.Ex. SO4 -2 sulfate OH-1 hydroxide

Name the following:1.BeI2 beryllium iodide2.NiS nickel (II) sulfide3.NiCl3 nickel (III) chloride4.SnO tin (II) oxide5.(NH4)2Cr2O7 ammonium dichromate

WRITING FORMULASWRITING FORMULAS

1. Write the cation first with its charge.a. The charge on a transition metal cation is

given in the name.b. If a polyatomic ion exists, use your

polyatomic ion sheet.

2. Write the name of the anion next with its charge.a. find the charge using your periodic table

b. if a polyatomic ion exits, use your polyatomic ion sheet

3. Criss cross the charges to find subscripts.4. Rewrite the formula without charges.

Write formulas for the following compounds:

1. Cesium nitride Cs3N2. Strontium sulfate SrSO4

3. Copper (II) oxide CuO4. Ammonium fluoride NH4F5. Barium carbonate BaCO3


Formed when two atoms share electrons Usually involve 2 or more non-metal atomsSimilar electronegativity values

(electronegativity difference is less than 1.7).Can be pure covalent – equal sharing of e (<.3)Polar covalent – uneven sharing of electrons

(.3-1.7)Called molecules There are more covalent compounds than

ionic compounds.

Covalent CompoundsCovalent Compounds

Binary covalent – include 2 elements onlyOrganic molecules (which include carbon) are

covalent◦Glucose C6H12O6

◦Sucrose C12H22O11

◦Hydrocarbon C6H6 etc.

The shared valence electrons are found in a molecular orbital, formed by overlapping atomic orbitals of the atoms involved in bonding.

Naming Covalent CompoundsNaming Covalent Compounds

Covalent molecules are NOT named like ionic compounds!!!

Steps to naming binary covalent molecules: (only include 2 elements)


1. the first element (least electronegative) is named first

2. the second element has the ending –ide1. Prefixes are used before the element

name to indicate the number of each atom in the molecule (the subscript)


# of atoms Prefix

1 mono

2 di

3 tri

4 tetra

5 penta

6 hexa

7 hepta

8 octa

9 nona

10 deca


Do NOT use a prefix for the first atom when there is only one present.

When the element name begins with a vowel, the (a) and (o) are dropped from the end of the prefix


Examples:CCl4 carbon tetrachloride CO carbon monoxide (eliminate the o from the prefix)P2Cl5 diphosphorus pentachloride (more

than one of the first element)N2O3 dinitrogen trioxide Si3N4 trisilicon tetranitride H2O dihydrogen monoxide


Steps for writing covalent formulas from the name.

the elements appear in the same order as in the name

the prefix indicates the subscript in the chemical formula

Examples: Boron trifluoride BF3

Dinitrogen monoxide N2O Dinitrogen tetroxide N2O4

ENERGY and STABILITY of ENERGY and STABILITY of COVALENT BONDSCOVALENT BONDS

Most atoms have relatively LOW STABILITY and HIGH POTENTIAL ENERGY.

When a compound forms, the atoms become MORE STABLE and the potential energy is at a MINIMUM.

HYDROGEN POTENTIAL ENERGY HYDROGEN POTENTIAL ENERGY CURVECURVE


Hydrogen Potential Energy CurveWhen the nuclei are farthest apart, the

potential energy is ZERO.As they get closer, the energy

DECREASES.When the potential energy is LOWEST (at -

436 kJ//mole), the atoms bond.


The distance between the two nuclei at this lowest energy is called the BOND LENGTH and is 75 pm for the H2 molecule.

When the REPULSION of the two atoms perfectly balance the ATTRACTIVE FORCES between the two nuclei, a COVALENT BOND forms.

Since the potential energy DECREASED, energy has been RELEASED. For H2, the energy released is -436 kj/mol.


The energy released when forming the bond is the exact amount of energy that would be needed to break the bond. Energy needed = +436 kj/mol

The energy required to break a bond is known as the BOND ENERGY.

BOND ENERGY and BOND LENGTHBOND ENERGY and BOND LENGTH

Bond energy

(kj/ mol)

Bond length (pm)

Electronegativity

Difference

HF 570 92 1.8CF 552 138OO 498 121HH 436 75HCl 432 127 1.0CCl 397 177HBr 366 141 0.8HI 299 161 0.5

BOND ENERGY and BOND LENGTHBOND ENERGY and BOND LENGTH

BOND ENERGY (STRENGTH) is inversely related to the BOND LENGTH. The higher the bond energy, the SHORTER the bond.

Bond length is actually an AVERAGE DISTANCE between the two nuclei since the distance is constantly changing due to the bond being able to vibrate and bend.

BOND STRENGTH and POLARITYBOND STRENGTH and POLARITY

The HIGHER the electronegativity difference, the STRONGER the bond.

Ex. HF has a much stronger bond than HI.

HF is nearly IONIC and HI is almost PURE COVALENT

LEWIS DOT STRUCTURESLEWIS DOT STRUCTURES

Named after Gilbert Newton Lewis, who in 1920 came up with a way to represent valence electrons in an atom using dots.

SINGLE ATOM LEWIS STRUCTURES

Only show the valence electrons around the element symbol.

A maximum of 2 electrons per side of the element symbol.

To find valence electrons, use the group number from the periodic table.

LEWIS DOT STRUCTURES for LEWIS DOT STRUCTURES for COVALENTLY BONDED MOLECULESCOVALENTLY BONDED MOLECULES

Recall the octet rule – most atoms need 8 valance electrons to be satisfied.

Exceptions to the octet rule:◦H and He◦B◦Elements beyond period 3 can EXCEED the octet rule. This can happen since these atoms have d-orbitals available to store the extra electrons.

DRAWING LEWIS DOT STRUCTURESDRAWING LEWIS DOT STRUCTURES

1. Use a PENCIL. You may have to move electrons around and this becomes frustrating and messy if you use a pen.

2. Count the total number of valence electrons available.

-For ions, a negative charge increases the number of electrons and a positive charge decreases the number of electrons.


3. Try to find a central atom and place all other atoms around it. Use symmetry whenever possible. Other times, the order in which the formula is written will allow you to determine which atoms should be connected.

4. Use dots to represent the electrons shared between atoms. 2 electrons represent a BONDING PAIR or a single bond. Use the available valence electrons to make bonds between the atoms. These electrons would actually exist in molecular orbitals.


5. Fill in the leftover electrons as LONE PAIRS around the atoms to achieve an octet (remember the exceptions). Electrons must always be paired. Lone pair electrons do not participate in the bonds, they are held in the atomic orbitals of the atoms involved.

6. Count the electrons and make sure all atoms have been satisfied. You may need to make some lone pair electrons into bonding electrons (in double or triple bonds) to satisfy some of the atoms involved.


7. Any extra electrons should be placed as lone pairs on the central atom. (This will happen if the central atom is one that can exceed the octet rule.)

8. Polyatomic ions should have their structure placed in a bracket with the overall charge placed outside of the bracket.

RESONANCERESONANCE

when 2 or more equivalent Lewis structures exist for

a molecule or compound

all structures should be represented with a double

arrow between

the actual bond strength is an average of all the

bonds.

Examples:

(see back of notes)

MOLECULARGEOMETRY: MOLECULARGEOMETRY:

ARRANGEMENT and SHAPEARRANGEMENT and SHAPE

The three dimensional shape of molecules can be predicted

using the VALENCE SHELL ELECTRON PAIR REPULSION

THEORY.

◦maximizes space between electron pairs to predict shape.

◦VSEPR

Arrangement of electrons – based on the number of regions

of electron density around the central atom.

Shape of the molecule – can be predicted by counting the

number of bonding pair electrons and lone pair electrons.

MOLECULARGEOMETRY: ARRANGEMENT and SHAPEMOLECULARGEOMETRY: ARRANGEMENT and SHAPE

Treat double and triple bonds as single electron pairs when

determining VSEPR shape.

once the shape is determined, certain bond ANGLES can be

predicted. Bond angles are altered by the present of LONE PAIR

ELECTRONS. When present, the bond angles are compressed

due to the REPULSION of the.

lone pairs require more space since they are only controlled by

one nucleus.

The arrangement of electrons and the shape of the molecule

with be the same when there are NO LONE PAIR ELECTRONS.

HYBRIDIZATIONHYBRIDIZATION

In order for bonds to have equivalent energy, mixing of orbitals must occur.

Ex. Methane (CH4) predicts 4 equivalent bonds. Which orbitals of C are used?

mix the orbitals used into a new “hybrid orbital” with a new name.

HybridizationHybridization

Total number of e- pairs around a central atom

Arrangement Hybridization

2 Linear sp3 Trigonal planar sp2

4 Tetrahedral sp3

5 Trigonal bipyramidal

sp3d

6 octahedral sp3d2

POLARITY and DIPOLE MOMENTPOLARITY and DIPOLE MOMENT

POLAR BONDS:

Result from high ELECTRONEGATIVITY DIFFERENCES between atoms of a molecule.

Causes PARTIAL POSITIVE and PARTIAL NEGATIVE ends of the molecule (called DIPOLES).

SHAPE AFFECTS POLARITYSHAPE AFFECTS POLARITY

DIPOLE MOMENT – overall direction of electron “pull” within a molecule. Show using a molecular model.

Sometimes dipoles will CANCEL each other, and the result will be a molecule with NO NET DIPOLE due to the shape.

Ex. HBr, BeCl2, BBr3, SeCl2, CO2, H2O

POLARITY AFFECTS PROPERTIESPOLARITY AFFECTS PROPERTIES

For example: ◦CO2 is non-polar, so the ATTRACTIVE FORCE between CO2 atoms is VERY WEAK.

◦Only London Disperson Forces exist. Weakest intermolecular force (IMF).

◦This results in a lower MELTING POINT and BOILING POINT.

POLARITY AFFECTS PROPERTIESPOLARITY AFFECTS PROPERTIES

H2O is POLAR, the molecules interact with each other and attractive forces are greater.

This results in a higher MELTING POINT and BOILING POINT.

A very strong INTERMOLECULAR force (called a dipole-dipole interaction) exists between water molecules due to its polarity.

This force is called HYDROGEN BONDINGIt is not the bond within water….it is a bond

between water molecules.

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CHAPTER 6 IONIC AND COVALENT BONDS, NAMING, AND MOLECULAR GEOMETRY