ch. 6. what is a chemical bond? mutual electrical attraction between the nuclei and valence...
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CHEMICAL BONDINGCh. 6
What is a chemical bond? mutual electrical attraction between the nuclei and
valence electrons of different atoms that bind the atoms together
Why don’t noble gases do this?Already have filled s and p orbitals stable octet: 8 valence e- (or 2 if, you’re helium)
Atoms that don’t have a stable octet are more reactive
Key Point #1: By forming bonds with each other, most atoms reduce their potential energy, becoming more stable.
This is a chemical change! All chemical changes involve energy!
What types of bonds can be formed?
Metallic Bonding
• In a metal, the empty orbitals in the atoms’ outer energy levels overlap
Delocalized Electron: outer electron that does not belong to any one atom but can move freely through the metal’s network of empty atomic orbitals.
sea of electrons: mobile electrons around the metal atoms, which are packed together in a crystal lattice.
metallic bonding: chemical bonding that results from the attraction between metal atoms and surrounding sea of electrons
Key Point: In metallic bonding, valence electrons move freely throughout a network of metal atoms.
Unique Characteristics of Metals
Metals have many unique properties because of their sea of electrons
• Malleability: ability of a substance to be hammered or beaten into thin sheets
• Ductility: ability of a substance to be pulled into a thin wire
• Why? atoms can slide past one another along a plane without breaking bonds
Luster: shiny appearance• Why? Absorb a wide range of light frequencies,
many orbitals separated by small energy differences
Conductivity• Thermal: ability to conduct heat• Electrical: ability to conduct electricity
• Why? Electrons move easily through network of empty orbitals
Metallic Bonding Strength
The strength of metallic bonding is determined by the enthalpy of vaporization:
• the amount of energy required to vaporize (turn into a gas) 1 mol of a metal
In general, the strength of the metallic bond INCREASES moving left to right across the periodic table.• Soft metals (less dense) metals harder (more
dense) metals toward right
Properties of Metals: Malleability and Ductility
Properties of Metals: Surface Appearance
Properties of Metals: Electrical and Thermal Conductivity
Types of Bonds What type of bonds can be formed?
Ionic bond Covalent bond
○ Nonpolar covalent○ Polar covalent
Ionic bonding: bonds that result from electrical attractions between cations and anions
1 atom losses electrons 1 atom gains electrons
Covalent bonding: sharing of electrons between 2 or more atoms
Key Point 2: Rarely is bonding between atoms purely ionic or purely covalent. Instead, it usually falls somewhere between the two
extremes. Why?
Key Point 3: The extent of ionic or covalent bonding between two atoms can be estimated by calculating the difference in each elements’ electronegativity.
Covalent Bonding Large difference in E.N.: bond has more ionic
character Small difference in E.N: bond has more covalent
character
Types of Covalent Bonds Non-polar covalent bonding: both electrons equally
shared between atoms
Polar covalent bonding: unequal attraction for the shared electrons
6.1 Practice WorksheetPart 1
The property of electronegativity, which is the measure of an atom’s ability to attract electrons, can be used to predict the degree to which the bonding between atoms of two elements is ionic or covalent.
The greater the electronegativity difference, the more ionic the bonding is.
If the calculated electronegative difference is…
> 1.7 : ionic bond is formed> 0.3 , < 1.7 : polar-covalent bond0 – 0.3 : non-polar covalent bond
Increasing difference in electronegativity
NonpolarCovalentshare e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
Elements Electronegativity Electronegativity Difference
Bond Type
Element 1
Element 2
Mg to Cl
H to O
C to Cl
N to H
C to S
K to F
Na to Cl
H to H
1.2
2.5
3.0
2.5
.8
.9
2.1
2.1
3.0
3.5
3.0
2.1
2.5
4.0
3.0
2.1
1.8
1.4
.5
.9
0
3.2
2.1
0
Ionic
Ionic
Ionic
Polar covalent
Polar covalent
Polar covalent
Non-polar covalent
Non-polar covalent
Polyatomic Ions It is also possible if a compound contains polyatomic
ions, for both types of bonding to be present.
Monatomic Ions: Fe2+ , Na+, Cl-
Polyatomic Ions: PO43-, NH4
+ , NO-1
Groups of atoms are bonded covalent together, but because of few or more than expected valence electrons they have an overall charge (so they can also bond ionically with other ions)
Ex: Ca2+ and SO42- CaSO4 (metal & diff. nonmetals)
Classify the following as ionic, covalent, or both
1. CaCl2 = __________
(metal & nonmetal)
2. CO2 = __________
(nonmetal & nonmetal)
3. MgO = __________
(metal & nonmetal)
4. HCl = ___________
(nonmetal & nonmetal)
5. BaSO4 = ___________
(metal & diff. nonmetals)
6. H2O = ____________
(nonmetal & nonmetal)
7. SO3 = ___________
(nonmetal & nonmetal)
8. AlPO4 = ___________
(metal & diff. nonmetals)
Covalent
Covalent
Covalent
Ionic
Ionic
Both
Covalent
Both
Covalent BondingSection 6.2
What is a molecule? Neutral group of atoms that are held together by
covalent bonds.
Chemical formula: indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.
Formation of Covalent Bonds The electrons of one atom and protons of the other
atom attract each another.
The two nuclei and two electrons repel each other.
These two forces cancel out to
form a covalent bond at a length
where the potential energy is
at a minimum.
Bond Length vs. Bond Energy Bond length (pm): distance between two bonded
atoms at their minimum potential energy
Bond energy (kJ/mol): energy required to break a chemical bond and form neutral isolated atoms.○ Breaking bonds: absorbs (requires) energy○ Forming bonds: releases energy
Key Point: As you increase the number of bonds between 2 atoms the bond energy increases, while the bond length decreases. This is an inverse relationship.
Bond Energies & Bond Lengths A. How many electrons are shared in a
single bond: double bond: triple bond:
B. Which bond is shorter? C – C or C = C C. Which bond requires more energy to break?
In addition to finding an ideal bond length, atoms also lower their potential energy by achieving a stable octet of 8 valence electrons
Bond Energies & Bond Lengths A. How many electrons are shared in a
single bond: 2 e- double bond: 4 e- triple bond: 6e-
B. Which bond is shorter? C – C or C = C C. Which bond requires more energy to break? =
In addition to finding an ideal bond length, atoms also lower their potential energy by achieving a stable octet of 8 valence electrons
Octet Rule Octet Rule: chemical compounds tend to form so that
each atom has an octet of e-’s in its highest occupied energy level
Exceptions to the octet rule:Atoms that cannot fit eight electrons Atoms that can fit more than eight electrons
Hydrogen: 2e-Boron: 6e-Phosphorus, Sulfur, & Xenon: expanded
valence, more than 8e-
How can we represent molecules? Lewis Structures: formulas in which atomic symbols
represent nuclei and inner shells, which are surrounded by dot-pairs/dashes represent valence electrons
MOLECULAR GEOMETRY
Chapter 6.5
VSEPR THEORY
Lewis Structures are 2D but we live in a 3D world!molecular geometry: the three-dimensional
arrangement of a molecule’s atoms
What do those 3D structure/shapes look like??Follow the Valance Shell Electron Pair Repulsion
Theory or VSEPR○ Repulsion between the sets of valence electrons
surrounding an atom causes them to be oriented as far away from each other as possible
Why use VSEPR Theory? Key Point: VSERP Theory is used to predict the shape
of molecules based on the fact that electron pairs strongly repel each other.
Following VSEPR allows us to predict bond polarity:uneven distribution of electrons
AB2 – Linear
Atoms bonded to
central atom (B)
Number of Lone Pairs on central
atom (E)
Bond Angle
2 0 180˚
Central atom
Atoms/group of atoms attached to central atom
Other Linear Geometries The shape of two atoms bonded together is not
given in the chart.Ex: F2
What is the only possible shape a binary compound can have?
○ LINEAR!
AB2E1 – BentAtoms
bonded to central atom
(B)
Number of Lone Pairs
(E)Bond Angle
2 1 <120˚What happens to the bond angle between atoms as you increase the number of “lone pair electrons” on the central atom?
Bond angles decrease!
AB2E2 – Bent
Atoms bonded to
central atom (B)
Number of Lone Pairs
(E)Bond Angle
2 2 104.5˚
AB3 – Trigonal Planar
Atoms bonded to
central atom (B)
Number of Lone Pairs
(E)Bond Angle
3 0 120˚Shape is often associated with atoms that break octet rule, but doesn’t have to be
AB3E1 – Trigonal Pyramidal
Atoms bonded to
central atom (B)
Number of Lone Pairs
(E)Bond Angle
3 1 107˚
AB4 – TetrahedralAtoms
bonded to central atom
(B)
Number of Lone Pairs
(E)Bond Angle
4 0 109.5˚
Predicting Molecular Geometry
1. Draw Lewis structure for molecule.
VSEPR theory: if any lone pairs of electrons are found on the central atom, these electrons
decrease the bond angles of atoms attached to it.
2. Draw a revised Lewis structure to show more accurate geometry
SO O
AB2E
bent
S
F
F
F F
AB4E
tetrahedral
3. To indicate the polarity of the bonds, we use this symbol: __________________ , which always points toward the more electronegative element.
4. When multiple bonds are found in a molecule, we must identify polarity of each bond.
Predicting Molecular Polarity
H F
electron richregion
electron poorregion FH
e- riche- poor
d+ d-
5. Observe the overall polarity of the molecule. Think of it as “tug-of-war” for valence electrons between the various atoms.
Non-polar covalent molecules: If the atom is symmetrical and all atoms have an equal pull on electrons
Polar covalent molecules: If the atom is not symmetrical and/or the atoms do not all have an equal chance of winning the tug of war for electrons
Intermolecular Forces
Intermolecular forces: attractive forces between molecules.
Intramolecular forces: attractive forces within a molecule (the bonds)
intermolecular forces are much weaker than intramolecular forces
Intramolecular Forces (bond)
Intramolecular Forces
Intermolecular Forces
Strength of IMF
Hydrogen Bond Dipole – Dipole Induced Dipole London Dispersion
Forces
strongest
weakest
H F
Dipoles What is a dipole?
A polar moleculeUneven sharing of electrons so there is a
separation of charge
FH
electron richregion
electron poorregion
e- poor
d+ d-
Dipole-Dipole Forces
Attraction between two polar molecules
— + — +
Hydrogen Bonding
Special type of Dipole – Dipole Attraction between:
Hydrogen & Nitrogen/Oxygen/Fluorine
Induced Dipole Attraction between one polar and one nonpolar
molecule
— +
— + — +
Electrons shift toward positive end
of dipole
London Dispersion Forces
Attraction between two nonpolar molecules
— + — +
Electrons become
uneven and form a dipole
What does IMF effect? Viscosity Surface Tension Boiling Point
Boiling Point
Point at which liquid particles escape the surface of the liquid into the gas phase
Stronger IMF Higher Boiling Point
Stronger IMF Higher Surface Tension
Surface Tension
result of an imbalance of forces at the surface of a liquid.
Stronger IMF Higher Viscosity
Viscosity
Measures a fluid’s resistance to flow