BONDING AND GEOMETRY

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BONDING AND GEOMETRY

• Unit 8
• Chemistry

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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 Na1?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

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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 shellplace where
  bonding occurs
• Ionic Bonding between a metal and a nonmetal
• Metallic Bonding between two metals
• Covalent Bonding between two nonmetals

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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

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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

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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

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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

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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

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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

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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

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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

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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

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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
  non-polar

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POLAR BONDS AND MOLECULES

• Example H2O

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INTRAMOLECULAR FORCES

• Intramolecular forces- (attraction is within the
  molecule)
• Types of forces covalent, Ionic, metallic

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INTERMOLECULAR FORCES

• Intermolecular Forces- (attraction is between
  molecules)
• Hydrogen bonding
• Dipole-Dipole Bonding
• Dispersion (London/van der Waals)

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INTRA/INTERMOLECULAR FORCES

• Which is a greater force Intermolecular or
  intramolecular?
• Rank the strength of each kind of inter and intra
  molecular force.

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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

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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

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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

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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

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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

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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

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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 1 3
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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

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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 molecules 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

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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

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VSEPR THEORY

• Tetrahedral
• Central atom bonds to 4 atoms and has zero lone
  pairs
• CH4

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VSEPR THEORY

• Pyramidal
• The central atom bonds to 3 atoms and has 1 lone
  pair of electons
• NH3

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VSEPR THEORY

• Trigonal Planar
• The central atom bonds to 3 atoms and has zero
  lone pairs
• BF3

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VSEPR THEORY

• Bent Triatomic
• The central atom bonds to 2 atoms and has 2 lone
  pair of electrons
• H2O

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VSEPR THEORY

• Linear Triatomic
• The central atom bonds to 2 atoms and has zero
  lone pair of electrons
• CO2

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VSEPR THEORY

• Linear
• One bond between 2 atoms
• HCl