Download BONDING AND GEOMETRY

BONDING AND
GEOMETRY
Unit 8
Chemistry
ATOMS AND IONS REVIEW
 Atoms are neutral
 They have the same number of protons and electrons
 Number of positives = number of negatives
 Example: Na 11 protons, 11 electrons  11 – 11 = 0
 Ions have a charge
 They have a different number of protons and electrons
 Example: Na+111 protons, 10 electrons 11 – 10 =
+1
 If an atom GAINS an electron  becomes negatively
charged  ANION
 If an atom LOSES an electron  becomes positively
charged  CATION
TYPES OF BONDS
 Bonding occurs because every element is
either trying to get to 0 electrons in the valence
or 8 electrons in the valence (zero and 8 are
both stable)
 Valence is the outer electron shell—place where
bonding occurs
 Ionic – Bonding between a metal and a
nonmetal
 Metallic – Bonding between two metals
 Covalent – Bonding between two nonmetals
IONIC BONDING
 Very stable and strong
 Strongest possible bond
 Requires a large amount of energy to break an
ionic bond
 Forms compounds known as “ionic
compounds”
 All ionic compounds will dissolve in water and
carry a current (electrolyte)
 Generally have high melting and boiling points
 Compounds are generally hard and brittle
METALLIC BONDING
 Metal atoms are pieces of metal that consist of closely
packed cations (positively ions)
 Cations are surrounded by mobile valence electrons that are
free to drift from one part of the metal to another
 Metal atoms are arranged in very compact and orderly
(crystalline) patterns
 Metallic bonding is the electrostatic attraction between
conduction electrons, and the metallic ions within the
metals, because it involves the sharing of free
electrons among a lattice of positively-charged metal
ions
 Occurs between 2 or more metals
 Result of the attraction of free floating valence electrons for the
positive ion
 These bonds hold metals together
COVALENT BONDING
 Covalent:
 Covalent bonds are when atoms SHARE
VALENCE electrons
 A covalent compound is called a
molecule
 Covalent bond ALWAYS occurs between
2 nonmetals
TYPES OF COVALENT BONDS
 Single Bond
 Covalent bond where one pair of electrons (2
electrons total) are shared between 2 atoms
 Atoms share electrons so that each has a full octet
(8 valence)
 Electrons that are shared count as valence
electrons for both atoms
 Examples
 Cl2
COVALENT BONDING
 Double Bonds
 Bond in which two pairs of electrons (4
electrons total) are shared between 2 atoms
 Examples
 O2
 Triple Bonds
 Bond in which 3 pairs of electrons (6 total
electrons) are shared between atoms
 Examples
 N2
COVALENT BONDING
 Electron Pairs
 Electron pairs involved in the actual bond are called
BONDING PAIR or SHARED PAIR electrons
 Electrons not involved in the actual bond, those
surrounding the rest of each element are called
LONE PAIR electrons
POLAR BONDS AND
MOLECULES
 Covalent bonds are formed by sharing
electrons between two atoms
 The bonding pair of electrons is shared
between both elements, but each atom is
tugging on the bonding pair
 When atoms in a molecule are the same (diatomic)
the bonding pair is shared equallythis bond is
called non polar covalent
 When atoms in a molecule are different, the bonding
pair of electrons are not shared equallythis is
called a polar covalent bond
POLAR BONDS AND
MOLECULES
 Why is the bonding pair not shared
equally?
 The answer lies within electronegativity
 One of the elements is more electronegative than
the other and therefore has a greater desire for
the shared pair
 The MORE electronegative element tends to pull
the electrons closer and thus has a slightly
negative charge
 The LESS electronegative element has a slightly
positive charge since the shared pair is being
pulled away
POLAR BONDS AND
MOLECULES
 Polar Molecules
 Molecule in which one end of the molecule is
slightly negative and the other end is slightly
positive
 Just because a molecule contains a polar
bond DOES NOT mean the entire molecule
is polar
 The effect of polar bonds on the polarity of
an entire molecule depends on the shape of
the molecule and the orientation of the polar
bonds
POLAR BONDS AND
MOLECULES
 Example: CO2
O=C=O
 Carbon and Oxygen lie along the same axis.
 Bond polarities are going to cancel out
because they are in opposite directions
 Carbon dioxide is a nonpolar molecule even
though there are two polar bonds present
 Would cancel out if the polarities moved towards
each other as well
 When polarities cancel out, the molecule is nonpolar
POLAR BONDS AND
MOLECULES
 Example: H2O
INTRAMOLECULAR
FORCES
 Intramolecular forces- (attraction is
within the molecule)
 Types of forces: covalent, Ionic, metallic
INTERMOLECULAR
FORCES
 Intermolecular Forces- (attraction is
between molecules)
 Hydrogen bonding
 Dipole-Dipole Bonding
 Dispersion (London/van der Waals)
INTRA/INTERMOLECULAR
FORCES
 Which is a greater force: Intermolecular
or intramolecular?
 Rank the strength of each kind of inter
and intra molecular force.
FORCES IN A MOLECULE
 Dipole-Dipole Forces
 Dipoles are created when equal but opposite
charges are separated by a short distance
 Have to have a positive and a negative end so
that one of the elements is pulling on the electron
 Only happens in polar molecules
 Dipole forces are extremely strong and lead
to high melting and boiling points
FORCES IN A MOLECULE
 Hydrogen Bonding
 Very strong type of dipole force
 Only occurs when hydrogen is covalently
bonded to a highly electronegative atom
 Always involves hydrogen
 Example: HF, H2O, NH3
FORCES IN A MOLECULE
 H – bonding and boiling point
 Why does boiling point increase as you go
down a group?
 The increase in boiling point happens because the
molecules are getting larger with more electrons, and
so dispersion forces become greater
FORCES IN A MOLECULE
 London Dispersion Forces
 Electrons are in constant motion around a
nucleus
 At any given time there might be more
electrons on one side of an atom than on the
other
 For a split second, the side with more
electrons is negative, and the side with less
electrons is positive
FORCES IN A MOLECULE
 London Dispersion Forces
 Recall that Noble Gases have a full outer
shell and you have been told they are
unreactive BUT due to London Dispersion
Forces, they COULD bond for an instant
 Example: Ar2
 London Forces are very weak
 The smaller the mass of the atom, the
smaller the London Force
BOND DETAILS
 Terminology
 Bond strength-energy required to break a bond
 Bond axis-imaginary line joining two bonded atoms
(example: C-C)
 Bond length-the distance between two bonded
atoms at their minimum potential enery; the average
distance between two bonded atoms
 Bond energy-energy required to break a chemical
bond and form neutral isolated atoms
 Chemical compound tend to form so that each
atom, by gaining, losing, or sharing electrons,
has an octet of electrons in its highest
occupied energy level
BOND DETAILS
 Comparison of Bond Length/Strength for
Covalent Bond Types:
 Longer bond = less bond strength
 Rating 1-3 (with 3 as the largest and 1 as
the smallest)
Bond
Length Strength
Single
3
1
Double 2
2
Triple
3
1
BOND DETAILS
 Coordinate Covalent Bonds
 Very rare
 Tend to form harmful molecules
 Occurs when both of the bonding pair of
electrons in a covalent bond come from only
ONE of the atoms
 Example: CO
BOND DETAILS
 Resonance
 Occurs when there are more than one possible
structures for a molecule
 Refers to bonding in molecules or ions that cannot
be correctly represented by a single Lewis structure
 Example: CO2
 To indicate resonance, a double-headed arrow is placed
between a molecule’s resonance structures
 Even though all of the structures are different, the
number of bonding pair of electrons and lone pair of
electrons stay the same in each structure
VSEPR THEORY
 Valence Shell Electron Pair Repulsion
Theory
 Allows us to picture molecules in 3
dimensions
 Centers around the fact that electrons
have negative charges and repel one
another
 So electron pairs within a structure try to
arrange themselves to be as far away
from other pairs as possible
VSEPR THEORY
 Tetrahedral
 Central atom bonds to 4 atoms and has zero
lone pairs
 CH4
VSEPR THEORY
 Pyramidal
 The central atom bonds to 3 atoms and has
1 lone pair of electons
 NH3
VSEPR THEORY
 Trigonal Planar
 The central atom bonds to 3 atoms and has
zero lone pairs
 BF3
VSEPR THEORY
 Bent Triatomic
 The central atom bonds to 2 atoms and has
2 lone pair of electrons
 H2O
VSEPR THEORY
 Linear Triatomic
 The central atom bonds to 2 atoms and has
zero lone pair of electrons
 CO2
VSEPR THEORY
 Linear
 One bond between 2 atoms
 HCl