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Thermochemistry and Particles (Periodic Trends (Atomic Forces (Protons and…
Thermochemistry and Particles
Periodic Trends
Electron configuration
When non-metals gain electrons to form anions, they are added to the p sublevel.
When metals lose electrons, they are removed from sublevel furthest from the nucleus
Atomic Forces
Protons and electrons hold atom together
Protons act as point charge
Distance of valence electrons from nucleus
Shielding: Electron repulsion
Radius
Atomic
Distance between outermost electron and nucleus
Increase down a group
Decrease across a period
Ionic
Cation
Smaller than atom
Lost valence shell
Less shielding
Anion
Larger than atom
Increased electron-electron repulsion in same shell
Ionization Energy
Energy to remove n mol of electrons from 1 mol of gas atoms
Increase across the period
Increase in nuclear charge
No additional shielding
Increase in attraction between nucleus and valence electron
Decrease down a group
Increase in nuclear charge
Outweighed by energy level increase
Decrease in attraction between nucleus and valence electron
Electronegativity
Ability of an atom to attract electrons
Same trends as 1st IE
Shapes of Molecules
Bonds repel each other as far as possible
Formal Charge
For molecules with elements not obeying octet rule
Difference between number of valence electrons in a single atom of that element and the number of valence electrons that the atom has in the molecule
Each bond is counted as 1 electron
Polarity
Polar
Covalent bond
Asymmetrically distributed around central atom
Non-polar
No covalent bond
Or covalent bonds symmetrically distributed around central atom
Like dissolves like
Thermochemistry
Mixed characteristics
Depends on electronegativities
Metallic Bonds
Cations in lattice
Electrons delocalized
Between cations and electrons
Characteristics
Malleable
Ductile
No direction
Conductive
High melting point
Strength
Boiling point is a good measure
Increases with atomic number
Affected by number of valence electrons
The greater, the stronger the attraction
Ionic bonding
Bond strength according to Coulomb's law
Lattice energy determines strength
Increases with electronegativity difference
Lattice structure
Metal and non-metal
Electrostatic attraction between cations and anions
Characteristics
Conducts when dissolved
Hard but brittle
Directional bonds
Covalent Network Solids
Two atoms share electrons
"Full" outer shell achieved
Covalent bonds connecting all the atoms in solid
Characteristics
Very high boiling point
Hard and rigid
Intermolecular Forces
Temporary Dipole-Induced Dipole Attractions
Temporary distortions in electron cloud ditribution
Induces charges on neighbouring molecules
Strength
Electron cloud size
Increases with size as can create larger dipoles more frequently
Molecular shape
Increases with linear shapes, as can stack nicely
Permanent Dipole-dipole attractions
Dipoles arise from differing electronegativities
Must be polar shape
Only significant in comparison to molecules of similar electron number
In addition to Temporary dipole-induced dipole attractions
Greater if dipole is on exterior
Strength
Polarity of bonds
Shape of molecule
Hydrogen bonding
Occurs between H and F, O or N
Big difference in electronegativity between the atoms
H only has 1 electron, so it basically becomes a positive charge
F, O or N become very negative, because it is also very small
Enthalpy Changes
Hess's Law
Add up equations
Justify
Magnitude
Sign
Spontaneity
Enthalpy is a property of a particular state under a particular set of conditions and does not depend on how you
get to that state
General Enthalpies of Reaction
Enthalpy of reaction
1 mole of matter goes through a chemical reaction
Enthalpy of combustion
1 mole of compound undergoes combustion with oxygen
Enthalpy of formation
1 mole of compound formed from its elements
Bond dissociation enthalpy
Break 1 mole of a specific bond in specific compound
Phase changes
Latent heat of fusion
Latent heat of vaporisation
Latenet heat of sublimation
Specific heat capacity
Born-Haber Cycle and Lattice Enthalpy
Atomisation enthalpy
To find strength of ionic bonds
Ionisation energy
Electron affinity
Lattice formtion enthalpy
Lattice dissociation energy
Use Hess's law to connect reactants and products
