Chemical energetics (Section 2 of 2)

What does Lattice enthalpy depend upon?

• Charge of the ions: As the charge increases the lattice enthalpy becomes more exothermic For Example: Lattice enthalpies of  (Mg has a 2+ charge while Li has a 1+ charge)
• MgO.......-1050 KJ/mol
• LiO.........- 3923 KJ/mol
• Size of the ions: the lattice enthalpy becomes less exothermic as the size of the ion decreases, because as the size increases the the attraction between the two oppositely charged ions decreases as both their centres of charge are further away from each other

Stabilities of Group II carbonates:

The trend: The carbonates get more and more stable as we move down the group because down the group polarization power of the cation decreases

What is polarization?

The tendency for an ion to reshape (or redistribution of charge) due to external electric fields is termed polarization.

How does polarization affect the stabilities of group II carbonates?

As we move down a group the size of the ion increases  hence the charge density decreases which thus causes a decrease in polarizing power of the cation and when this happens the Carbonates becomes more stable as we go down the group and hence their decomposition temperatures increase down the group.

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Chemical energetics (Section 1 of 2)

The Born Haber Cycle

What is the Born Haber Cycle?

The born Haber cycle can be used to determine lattice enthalpy of a compound. In the born Haber cycle every step from the elements to the ionic compound can be measured experimentally except the lattice enthalpy which therefore can be calculated using the Hess's law

What is Lattice enthalpy?

Lattice enhalpy is the energy change that occurs when 1 mole of an ionic compound is formed from its gaseous ions under standard conditions (1 atm Pressure and 298 K Temperature)

The Born Haber Cycle of NaCl
Some terms used in the following slides:

• ΔH atomization: The energy change(enthalpy change) that occurs when one mole of gaseous atoms are formed from an element in its standard state
• ΔH First electron affinity: The enthalpy change that occurs when one electron is added to each gaseous atom to form one mole of 1- gaseous ions.

Video explanation of the above cycle

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Chemical Bonding (Section 2 of 2)

Contents of this page

1. Intermolecular forces of attraction
1. Hydrogen Bonding
2. Vander Waals Forces
1. Temporary Dipole induced dipole Forces
2. Dipole Dipole forces of attraction

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Van der Waals Forces and Hydrogen Bonding

Hydrogen Bonding (The strongest Force)
What is Hydrogen Bonding?

• In hydrogen bonds, the positively charged hydrogen end of one molecule is attracted to the negatively charged end of another molecule which must be an extremely electronegative element (fluorine, oxygen, or nitrogen - FON)
  • e.g.  H₂O, HF, and NH_3.
  • Alcohols can also form hydrogen bonds with water in which the O in the OH group of the alcohol bonds to the positively charged H end of the water molecule and the H in the OH group of the alcohol bonds to the negatively charged O of the water molecule.
  
• Hydrogen bonds are the strongest weak bond because the H atom essentially gives its single electron to form a bond and is therefore left unshielded.  The relatively strength of hydrogen bonds results in higher melting and boiling point temperatures than those in molecules with other van der Waals forces of attraction.
• Hydrogen bonding can explain why water is less dense in the solid phase than it is in the liquid phase (contrary to most other substances).  The hydrogen bonds between water molecules in ice to form a crystal structure, keeping them further apart than they are in the liquid phase.

Dipole dipole forces of attraction (The Medium Weak Force)

• Dipole-dipole forces exist between neutral, polar molecules where the positive end of one molecule is attracted to the negative end of another molecule.

• The greater the polarity (difference in electronegativity of the atoms in the molecule), the stronger the dipole-dipole attraction.
• Dipole-dipole attractions are very weak and substances held together by these forces have low melting and boiling point temperatures. Generally, substances held together by dipole-dipole attractions are gases at room temperature.

Dipole Induced Dipole Forces of attraction/London Dispersion forces (The weakest)

• London dispersion forces (LDF) occur between neutral, nonpolar molecules.  LDF occur due to the "random motion of electrons."  At any moment, one atom may be surrounded by an extra electron from a neighboring atom, resulting in an instantaneous polarity on the atom.  During that instant, the "polarized" atom will act as a very weak dipole.
• Since LDF are dependent upon the random motion of electrons, the more electrons an atom or molecule has, the greater the LDF between them.
• LDF as a whole are extremely weak, so substances held together by these forces have extremely low melting and boiling point temperatures.  These substances tend to be gases at room temperature.

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States of Matter

The Kinetic Theory of Matter

i      All matter is composed of small particles.

ii     The particles of matter are in constant motion.

iii     All collisions between the particles of matter are perfectly elastic

Particle motion: atoms move in a straight line between collisions.

The Solid State

Gases: The Kinetic Theory
Two of the basic assumptions of this theory are:

• Actual volume of gas molecules in negligible as compared to the volume occupied by the gas
• Intermolecular forces between gas molecules are negligible

What are Ideal Gases?
These are the gases which obey the kinetic theory of gases. 

Under what conditions of temperature and pressure do gases deviate from ideal behaviour?

1. Low Temperature
2. High pressure

Smaller molecules like Helium, Nitrogen and Hydrogen behave more ideally due to their smaller molecular size, hence minimal intermolecular forces of attraction. 
Othher gases like Ammonia, Hydrogen Chloride and Sulphur dioxide deviate more from the ideal behavior due to much stronger intermolecular forces.


The Ideal Gas Equation

PV=nRT

wWhere,

P=Pressure of gas in Pascals or Newton per metre squared

V=Volume of gas in m(3) (metre cubed)

n= Number of moles of gas

R= Ideal gas constant 8.31 J /K mol

T= Temperature of gas in kelvins

Calculating Mr of a gas from the ideal gas equation

Mr = mass x (R) (T)

              (P) x (V)

The Liquid State: Vaporization and Vapor pressure

Vapor Pressure

It is the pressure exerted by the vapors of a liquid when they are in a state of dynamic equilibrium
Factors affecting vapor pressure;

• Temperature
• Intermolecular forces
  

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