CHAPTER 5 and 6 OVERVIEW

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ACADEMIC CHEMISTRY NAME:
PERIOD: DATE:
TYPES OF BONDING in COMPOUNDS

IONIC BONDS COVALENT BONDS

What happens to
the electrons in POLAR COVALENT PURE COVALENT
this bond?

What is the
electro-negativity
difference of the
atoms involved?

What type of
elements are
involved?

What is the
simplest unit of
this compound
called?

How is bond
strength
measured?
ACADEMIC CHEMISTRY NAME:
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IOINIC BONDS COVALENT BONDS

What are
some
properties of
compounds
with this
type of
bonding?
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IONIC COMPOUNDS
Ion Formation
Why do ions form?

 Atoms become stable by gaining or losing VALENCE electrons to achieve a full outer ENERGY level.

o Losing electrons produces a POSITIVE charge and the ion is called a CATION.

o 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]1s1  [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:

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

2. TRANSITION METALS do not have to achieve noble gas configuration to become more stable.
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Ex. Fe can have charges of +2, +3, or +6

Energy level diagram: (see board)

Lose the 2 electrons in the 4 s orbital: +2
Lose the 2 electrons in the 4s and 1 from 3d: +3
Lose the 6 electrons in the 3d: +6

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

 Cations are smaller than the parent atom.

 Anions are larger than the parent atom.

o 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!!!

POLYATIOMIC IONS
 A group of two or more atoms COVALENTLY bonded together with an OVERALL CHARGE.

See polyatomic ion sheet.
Memorize those with *. You will be quizzed on their names and formulas (with charges).
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Ex. NH4+

Nitrogen and hydrogen share electrons.
(copy drawing from the board)

NH4+ - covalent inside

NH4Cl – ionic compound

*** THE BONDING WITHIN A POLYATOMIC ION IS COVALENT….but WHEN A POLYATOMIC ION IS IN A COMPOUND, IT WILL
BE AN IONIC COMPOUND

IONIC FORMULAS AND NAMING IONIC COMPOUNDS

 Ionic compounds must have a net charge equal to zero.

 Subscripts are used to balance positive and negative charges. Use parentheses 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 with its charge.

Ex.

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

Ex.
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Rules for Naming Ionic Compounds

1. Name the cation first.

a. Use the name of the element from the periodic table.

Ex.

b. If the cation is a transition metal, you must indicate the charge (oxidation state) on the ion using Roman numerals in

parentheses following the cation name.

2. Name the anion second.

a. Use the name of the anion form the periodic table, drop the ending and add –ide.

Ex. Oxygen becomes oxide.

Phosphorus becomes phosphide.

Nitrogen becomes nitride.

Sulfur becomes sulfide.

3. If a polyatomic ion is present, use the name given on the polyatomic ion sheet.

Ex.

Write the names of the following compounds.

BeI2
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NiS
NiCl3
SnO
(NH4)Cr2O7

Writing Formulas From the Name

1. Write the cation first with its charge.

a. The charge on transition metals is given in the name.

b. If a polyatomic ion exists, use your polyatomic ion sheet.

2. Write the anion next with its charge.

a. Find the charge using your periodic table.

b. If a polyatomic ion exists, use your polyatomic ion sheet.

3. Criss-cross the charges to find subscripts.

4. Rewrite the fomula without charges.

Ex. Write the formulas for the following compounds.

1. cesium nitride
2. strontium sulfate
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3. copper (II) oxide
4. ammonium fluoride
5. barium carbonate

COVALENT COMPOUNDS

 Formed when two atoms share electrons
o Usually involve 2 or more non-metal atoms
o Similar 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)

o Called molecules

o There are more covalent compounds than ionic compounds.

 Binary covalent – include 2 elements only

 Organic molecules (which include carbon) are covalent
 Glucose C6H12O6
 Sucrose C12H22O11
 Hydrocarbon C6H6 etc.
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 The shared valence electrons are found in a molecular orbital, formed by overlapping atomic orbitals of the atoms involved in
bonding.

 Recall the properties of covalent molecules listed previously in chapter 5 notes (flow chart).

NAMING 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 –ide
3. 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
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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
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Si3N4
trisilicon tetranitride

H2O
dihydrogen monoxide

Steps for writing covalent formulas from the name.

1. the elements appear in the same order as in the name
2. the prefix indicates the subscript in the chemical formula

Examples:

Boron trifluoride BF3

Dinitrogen monoxide N2O

Dinitrogen tetroxide N2O4
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ENERGY and STABILITY of COVALENT 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.

Example:

The potential energy curve for hydrogen.

 When 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 H 2, the energy released is -436 kj/mol.
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 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 LENGTH

Bond energy Bond length (pm) Electronegativity Difference
(kj/ mol)
HF 570 92 1.8
CF 552 138
OO 498 121
HH 436 75
HCl 432 127 1.0
CCl 397 177
HBr 366 141 0.8
HI 299 161 0.5

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

 The HIGHER the electronegativity difference, the STRONGER the bond.

o Ex. HF has a much stronger bond than HI.

o HF is nearly IONIC and HI is almost PURE COVALENT.

LEWIS 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.

Ex.
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LEWIS DOT STRUCTURES for COVALENTLY BONDED MOLECULES

 Recall the octet rule – most atoms need 8 valance electrons to be satisfied.
 Exceptions to the octet rule:
o H and He
o B
o Elements beyond period 3 can EXCEED the octet rule. This can happen since these atoms have d-orbitals available to
store the extra electrons.

Ex.

Steps to completing a Lewis dot structure:
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.
a. For ions, a negative charge increases the number of electrons and a positive charge decreases the number of electrons.

Ex.

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.
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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.

RESONANCE

 when 2 or more equivalent Lewis structures exist for a molecule or compound

 all structures should be represented with a double arrow between
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 the actual bond strength is an average of all the bonds.

 Examples: SO2 NO2

(see back of notes)

MOLECULARGEOMETRY: ARRANGEMENT and SHAPE

 The three dimensional shape of molecules can be predicted using the VALENCE SHELL ELECTRON PAIR REPULSION THEORY.

o maximizes space between electron pairs to predict shape.

o 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.

 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.

o 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.
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VSEPR and MOLECULAR GEOMETRY

BONDING PAIRS LONE PAIR TOTAL NUMBER ARRANGEMENT SHAPE SKETCH and
of ELECTRONS ELECTRONS OF ELECTRON BOND ANGLES
PAIRS
1 0 1

2 0 2

3 0 3

2 1 3
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4 0 4

3 1 4

2 2 4

5 0 5
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4 1 5

3 2 5

2 3 5
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BONDING PAIRS LONE PAIR TOTAL NUMBER ARRANGEMENT SHAPE SKETCH
of ELECTRONS ELECTRONS OF ELECTRON and
PAIRS BOND ANGLES
6 0 6

5 1 6

4 2 6
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3 3 6

2 4 6

HYBRIDIZATION
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.
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Total number of e pairs around a Arrangement Hybridization
central atom
2 Linear sp
3 Trigonal planar sp2
4 Tetrahedral sp3
5 Trigonal sp3d
bipyramidal
6 octahedral sp3d2

POLARITY 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 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.

Molecule Sketch Direction of dipole or no net dipole
HBr

BeCl2
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BBr3

SeCl2

CO2

H2O

POLARITY AFFECTS PROPERTIES:

 For example:
o CO2 is non-polar, so the ATTRACTIVE FORCE between CO2 atoms is VERY WEAK.
o Only London Disperson Forces exist. Weakest intermolecular force (IMF).
o This results in a lower MELTING POINT and BOILING POINT.

 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 BONDING

 It is not the bond within water….it is a bond between water molecules.