4. Bonding Flashcards
Ionic Bonding
The electrostatic attraction between oppositely charged ions
- Cations: positive charge, lose valence electrons (metals)
- Anions: negative charge, gain valence electrons (non-metals)
Ionic Bond Properties
- non directional bond - strength of bond eaqual in all directions
- conducts electricity when molten/ in solution
- high melting/ boiling points
- hard solids
- low volatility
Covalent Bonding
The electrostatic attraction between a pair of electrons and positively charged nuclei
- bonding electron pairs
- electrons not involved in bonding are called non-bonding e- pairs
Electronegativity Difference
- Covalent - high electronegativity locks the electrons in covalent bonds
- Metalic - Low electronegativity allows electrons to float in seas
- Ionic - high electronegativity difference means uneven sharing
< 0.4 ——————————————— 1.8<
nonpolar polar ionic
covalent bond covalent bond
Polarity
- Partial positive charge (positive/ negative)
- The overall charge = Dipole movement
- The molecular is polar if the electron densities are not symmetrical (eg H2O)
Ionic Bond strength
- electrostatic attraction between positive and negative ions
- the smaller the ions / the greater the charge on the ions, the stronger the attraction between the ions
- greater charge density within the structure
Covalent Bond strength
- shared pair of electrons between atoms
-
the shorter the bond, the stronger it is
- triple bonds are shorter and stronger than double bonds
- triple > double > single
Metallic Bond strength
- attraction between lattice of positive metal ions and delocalized outer shell electrons
- the smaller the metal ions, the greater the charge on ions
- the more delocalised outer shell electrons there are, the stronger the attraction between the ions and electrons
- greater charge density within the structure
Lewis Structures
- representations of molecules showing all electrons, bonding and nonbinding
- each bond contains 2 electrons
- elements pair up according to the octet rule
Exceptions to the Octet rule
-
Boron - forms stable compounds with just 3 valence electrons
- highly reactive
- Beryllium - forms stable compounds with just 2 valence electrons
- compounds that break the octet rule are often toxic and dangerous, ready to react so the octet rule is obeyed
VSEPR Theory
Valence Shell Electron Pair Repulsion
Electron domain - the number of bonds/ electron pairs there are on the central atom
Shapes w/ non-bonding electron pairs
If a molecule / ion has lone pairs on the central atom, the shapes are slightly distorted away from the regular shapes
because of extra repulsion caused by the lone pairs
Shape Angles - Linear
180
Shape Angles - Bent
105
Shape Angles - Trigonal Planar
120
Shape Angles - pyramidal
107
Shape Angles - Tetrahedral
109.5
Shape Angles - Octahedral
90
Resonance structures
When a lewis structure allows for the same arrangement of atoms but a different but equally valid arrangement of electrons, resonance occurs
- Electrons are delocalized and are spread over 2 bonding orbitals
- Located in the p orbitals - overlapping p orbitals create pi bonds with delocalized electrons
Allotropes
Compounds of the same element that differ in structure
eg Carbon
Diamond (structure & properties)
- Each carbon bonded to 4 other carbons by sp3 hybridization
- High melting point
- strong directional covalent bond
- Extremely hard
- difficult to break atoms apart/ move them in relation to one another
- No electrical conductivity
- electrons are localized in specific bonds
- Insoluble in polar and non-polar solvents because molecular bonds are stronger than any IMF
Graphite (structure)
- Carbon atoms bonded via sp2 hybridization
- form sheets of 6 sided rings with p-orbitals perpendicular from plane of ringt
- presence of p orbitals allows for stronger london forces that hold the sheets together
Graphite (properties)
- Conducts electricity - delocalized electrons
- Slippery - sheets can slip past each other
- High melting point - delocalized electrons
- Insoluble
Fullerenes
Buckyballs
- carbon atoms bond in units of 60 atoms, with interlocking 6 sided and 5 sided rings
- sp2 hybridization - extra p orbitals form pi bonds resulting in the properties
Properties
- Super strong
- conduct electricity and heat with low resistance
- free radical scavenger
Silicon structures
- SiO2 repeating in a tetrahedral structure
- Si is larger than C so the Si-Si bond length is greater
- the greater the bond length, the lower the bond enthalpy
- easier to break
- more reactive than diamond
Intermolecular Forces - London Forces
- between 2 non-polar molecules
- caused by the motion of electrons around a nucleus
- creation of temporary dipoles / instantaneous induced dipoles
- increases with atomic radii
Intermolecular Forces - Dipole Dipole Forces
- between 2 polar molecules
- permanent dipole
- the partially charged ends of the molecules attract and repel each other
- LDFs are always acting as well
-
usually asymmetrical molecules
- except for bent structured molecules eg water
Intermolecular Forces - Hydrogen Bonds
- between an electronegative atom from one molecule and a hydrogen atom from another
- N, O, F
- super strong dipole dipole bond
- hydrogen bond 5% stronger than an average covalent bond due to lack of underlying layer of electrons
Effect of IMF on Boiling Points
- phase change is when the IMF are overcome
- liquid turns into a gas when the attractive forces between the particles are completely broken
- Covalent macromolecular structure - extremely high melting / boiling points
- Metals/ Ionic compounds - high boiling points due to ionic attractions
- The weaker the attractive forces, the more volatile the substance
Metallic Bonding
The electrostatic attraction between a lattice of positive ions and delocalized electrons
- strength of bond increases with charge of ion
- conducts electricty - delocalized electrons
- malleable - close-packed layers of positive ions can slide over each other without breaking more bonds than are made
Alloys
Mixtures of metals and one or more other elements
Produces increased
- strength
- durability
- hardness
- resistance to corrosion
- magnetic properties
- due to the different elements being of different sizes preventing the easy sliding between 2 atoms
Composition of Brass and Steel
Brass
Copper and Zinc
Steel
Iron and Carbon
Formal Charge (FC)
FC = valence e- - 0.5bonding e- - non-binding e-
Most favored Lewis Structure:
Structure with the FC closest to zero
The structure with the negative FC on the most electronegative atom
Overlapping Orbitals
- Sigma bonds result from the axial overlap of orbitals
- Pi bonds result from the sideways overlap of parallel p orbitals
- double bonds fomed by 1 sigma and 1 pi bond
- triple bonds formed by 1 sigma and 2 pi bonds
Sigma Bonds
- Head to head overlap
- Cylindrical symmetry of electron density around the internuclear axis
Pi Bonds
- Side-to-side overlap
- Electron density above and below the internuclear axis
- electron overlap is weaker, yet there is an extra bond –> energy required to break them is more
Importance of Ozone
- At ground level reacts with chemicals to form smog, harms respiratory systems and degrades materials
- At atmospheric levels is essential for life
- absorbs UV radiation
- exothermic reactions cause a temp. inversion in the stratosphere –> prevents convection keeping the layers of atmosphere stable
Ozone Formation Equation
O2(g) + UV (242nm) –> O•(g) + O•(g)
O•(g) + O2(g) –> O3(g)
exothermic
Ozone Depletion Equation
O3(g) + UV(330nm) –> O2(g) + O•(g)
O3(g) + O•(g) –> 2O2(g)
Catalytic Ozone Destruction - Nitrogen Oxides
NO•(g) + O3(g) –> NO2•(g) + O2(g)
NO2•(g) + O• –> NO•(g) + O2(g)
Catalytic Ozone Destruction - Chlorofluorocarbons
CCl2F2(g) –> CClF2•(g) + Cl•(g)
Cl•(g) + O3(g) –> O2(g) + ClO•(g)
ClO•(g) + O•(g) –> O2(g) + Cl•(g)
Bond Order
The measurement of the number of e- involved in bonds between two atoms in a molecule (if bond order = 0, there is no bonding)
Bond Order = Total bonding pairs / total no. of positions
eg Ozone O3 = 3/2
Hybridization (definition)
The mixing of different types of orbitals to produce new types of orbitals –> most hybrid orbitals are combinations of s and p orbitals that then form sigma bonds
number of electron domains = amount of hybridization there is
2: sp // 3: sp2 // 4: sp3
Two sp orbitals
- Triple bonds
- 2 sp orbitals form a sigma bond between the carbons, and two pairs of p orbitals overlap to form the 2 pi bonds
Three sp orbitals
- sp2 orbital on carbon overlaps to form a sigma bond with the orbital in oxygen
- the unhybridized p orbitals overlap in pi fashion
Four sp orbitals
- sp3 orbitals on carbon overlap to form a sigma bond with orbitals in H
