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

Click here for the next section

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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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Basic Calculations in A-level Chemistry

The Mole

Balancing chemical Equations

Molecular and empirical formulas

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Bond Angles

Valence-shell electron-pair repulsion (VSEPR) theory

• Assumes that each atom in a molecule will be positioned so that there is minimal repulsion between the valence electrons of that atom.

In simple molecules in which there are no nonbonding electrons, there are five basic shapes:

1. LINEAR - Bond angle = 180
   • All diatomic molecules are linear.
   • Molecules with two atoms around a central atom such as BF₂ are linear because positioning the two attachments at opposite ends of the central atom minimizes electron repulsion.
   • Generic Formula:  MX or MX₂ (where M is the central atom and X is are the bonding atoms).
     
2. TRIGONAL PLANAR - Bond angle = 120
   • Molecules with three atoms around a central atom such as BF₃ are trigonal planar because electron repulsion is minimized by positioning the three attachments toward the corners of an equilateral triangle.
   • Generic Formula:  MX₃ (where M is the central atom and X is are the bonding atoms).
     
3. TETRAHEDRAL - Bond angle = 109.5
   • Molecules with four atoms around a central atom such as CH₄ are tetrahedral because electron repulsion is minimized by position the four attachments toward the corners of a tetrahedron.
   • Generic Formula:  MX₄ (where M is the central atom and X is are the bonding atoms).
     
4. TRIGONAL BIPYRAMIDAL
   • Bond angle within the equatorial plane = 120
   • Bond angle between equatorial and axial plane = 90
   • Molecules with five atoms around a central atom such as PF₅ are trigonal bipyramidal.  Three of the attachments  are positioned in a trigonal plane with 120 bond angles.  The remaining two attachments are positioned perpendicular (90) to the trigonal plane at opposite ends of the central atom.  This arrangement of atoms minimizes electron repulsion.
   • Generic Formula:  MX₅ (where M is the central atom and X is are the bonding atoms).
     
5. OCTAHEDRAL - Bond angle = 90
   • Molecules with six atoms around a central atom such as SF₆ are octahedral.  Four of the attachments are positioned in a square plane with 90 bond angles.  The remaining two attachments are positioned perpendicular (90) to the square plane at opposite ends of the central atom.  This arrangement of atoms minimizes repulsion. 
   • Generic Formula:  MX₆ (where M is the central atom and X is are the bonding atoms).
      
     
     

Other shapes

BENT (ANGULAR or V-SHAPED)
• Molecules with two atoms and one or two pairs of nonbonding electrons around a central atom such as H₂O are bent.  It can be imagined that a linear molecule with two atoms attached to a central atom is altered when electrons are added to the top of the central atom.  The repulsion caused by the addition of these extra electrons causes the molecule to become bent.  The angle of bent molecules is less than 120 if there is one pair of nonbonding electrons and is less than 109.5 if there are two pairs of nonbonding electrons.
• Some molecules, such as NO₂ have two atoms and a single unpaired electron around a central atom.  These molecules are also bent due to the repulsion of the single atom added to the central atom.
• Generic Formula:  MX₂E or MXE₂ (where M is the central atom, X is are the bonding atoms, and E are nonbonding pairs of electrons).
  
TRIGONAL PYRAMIDAL
• Molecules with three atoms and one pair of nonbonding electrons around a central atom such as NH₃ are trigonal pyramidal.  These molecules are essentially tetrahedral molecules with one of the attached atoms replaced by a pair of nonbonding electrons.  The force of repulsion of these electrons makes the bond angle between the attached atoms less than 109.5.  For example, in NH₃, the H-N-H bond is 107.5.
• Generic Formula:  MX₃E (where M is the central atom, X is are the bonding atoms, and E are nonbonding pairs of electrons).
  

 

SQUARE PYRAMIDAL

• Molecules with five atoms and one pair of nonbonding electrons around a central atom such as BrF₅ are square pyramidal.  These molecules are essentially octahedral molecules with one of the attached atoms replaced by a pair of nonbonding electrons.  This leaves four atoms in a plane as a square base and one atom positioned perpendicular (90)  to this plane.
• Generic Formula:  MX₅E (where M is the central atom, X is are the bonding atoms, and E are nonbonding pairs of electrons).
  

SQUARE PLANAR

• Molecules with four atoms and two pairs of nonbonding electrons around a central atom such as XeF₄ are square planar.  These molecules are essentially octahedral molecules with two of the attached atoms opposite each other around the central atom each replaced by a pair of nonbonding electrons.  This leaves four atoms in a square plane.
• Generic Formula:  MX₄E₂ (where M is the central atom, X is are the bonding atoms, and E are nonbonding pairs of electrons).
  
  
  

  Source: http://library.thinkquest.org/C006669/data/Chem/bonding/shapes.html

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Chemical Bonding (Section 1 of 2)

Contents of this Page:

1. The Bond Energy
2. Factors affecting the Bond energy
3. The Four Main Types of Bonding
1. Ionic Bonding
2. Covalent Bonding
3. Dative Covalent Bonding
4. Metallic Bonding
4. Section 2

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The Bond Energy
The energy required to break one mole of covalent bonds in gaseous state is called bond energy
For Example
H=H Bond Energy = +436 KJ/mol
I-I Bond Energy = +151 KJ/mol


Factors affecting Bond Energy

• Number of Bonds: Triple bonds are stronger than double bonds and double bonds are stronger than single bonds.
• Bond Length: Bond length is the distance between two nuclei of covalently bonded atoms as the bond length increases bond strength decreases, hence bond energy decreases

Ionic
•  between atoms on Left Hand Side and atoms on Right Hand Side of Periodic Table
•  electrons are TRANSFERRED between atoms
•  atoms end up as ions
•  strong electrostatic attraction between ions of opposite charge
•  giant ionic crystal lattice structure
•  Properties ... high melting points.....brittle.....water soluble....conduct when molten or in aqueous solution

 
Ionic Bonding; a video demonstration:

Drawing dot and cross diagrams

http://www.youtube.com/watch?v=8BZBzFwVXl4

Covalent Bonding
•  between atoms of the same element;   (e.g. in N2, O2, diamond, graphite)
•  between atoms of different elements on the RHS of table; (e.g. CO2,  SO2)
•  when one of the elements is in the middle of the table;  (e.g. C, Si)
•  head-of-the-group elements with high I.E.’s , (e.g. Be in BeCl2)
•  consists of a shared pair of electrons, one electron coming from each atom
•  atoms share to try and get an ‘octet’ of electrons
•  leads to the formation of simple molecules and giant molecules (e.g. silica)

Polar Covalent Bonds
In bonds between atoms of the same element the sharing of the electrons is equal between the two atoms. When two atoms of different elements make a bond, the electrons will not usually be shared equally. The electrons are pulled more toward the more electronegative element.

What is electronegativity?
Electronegativity is the measure of the ability of an atom in a molecule to draw bonding electrons to itself. In general, electronegativity increases from bottom to top and left to right on the periodic table. Fluorine is the most electronegative element since it has a tendency to pick up electrons easily and hold on to them strongly. An element like cesium has a low electronegativity. The unequal sharing of electrons is called a polar covalent bond. The definition of a polar covalent bond is a covalent bond in which the bonding electrons spend more time near one atom than the other.

 

Dative Covalent Bonding
•  consists of a shared pair of electrons, both electrons from one atom
•  one species is a lone pair donor - LEWIS BASE
•  other species has space in outer shell to accept a lone pair - LEWIS ACID
•  once the bond has been formed it is the same as a covalent bond

Metallic Bonding
•  metal atoms arranged in regular lattice give up outer shell electrons
•  electrons form a mobile ‘cloud’ which binds metal ions together
•  strength of bond depends on number of electrons and size of ions
•  mobile electrons ... allow electricity to be conducted

Strength of a Metallic Bond:
As we go down the group the atomic size increases, the attraction between the nucleus and free electrons decreases and therefore the melting melting point also decreases down the group.
As we move across the period, atomic size decreases and the strength of the metallic bond increases therefore melting point also increases

Sources:
Images and definitions: www.knockhardy.org.uk
Videos: Youtube.com

The VSEPR theory 
VSEPR theory proposes that the geometric arrangement of terminal atoms, or groups of atoms about a central atom in a covalent compound, or charged ion, is determined solely by the repulsions between electron pairs present in the valence shell of the central atom. To view more about VSEPR theory visit this page

The Sigma and Pie bonds

Click here for Section 2

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Atoms molecules and stoichiometry

Before proceeding with this section please make sure you have seen all the lectures on the Basic Calculations in A-level Chemistry page

Mass number and atomic number

Atomic Number (Z)      Number of protonsin the nucleus of an atom
Mass Number (A)       Sum of the protons and neutronsin the nucleus

What is the avogadro's constant?
The Avogadro Constant is a constant number used to refer to atoms, molecules, ions and electrons. Its value is 6.023 x 10²³ mol^-1 . I.E 1 Mole of Sodium contains  6.023 x 10²³ number of particles.

What is the Relative Atomic Mass (Ar)?


The mass of an atom relative to the 12C isotope having a value of 12.000 is called as the Ar.

Relative Isotopic Mass Similar, but uses the mass of an isotope Eg 238U
Relative Molecular Mass (Mr) Similar, but uses the mass of a molecule Eg CO2,  N2
Relative Formula Mass Used for any formula of a species or ion E.g NaCl,  OH¯




What are Isotopes?Definition Atoms with ...  the  same atomic number but different mass number                                                                                                              or
                         the  same number of protons but different numbers of neutrons.

Important: Chemical properties of isotopes are identical 

Example: Chlorine contains 75% by mass Cl (35) and 25% by mass Cl (37). Calculate the relative atomic mass of Chlorine.

Ar of Cl = (75 x 35) + (25 x 37)  = 35.5

                              100

What is the mass spectra?

A mass spectrum is an intensity(Abundance) vs. m/z (mass-to-charge ratio) plot representing a chemical analysis. Hence, the mass spectrum of a sample is a pattern representing the distribution of ions by mass (more correctly: mass-to-charge ratio) in a sample.

How to calculate the Relative Atomic mass from a mass spectra

What is the empirical formula

A formula that gives the simplest whole-number ratio of atoms in a compound.

How to calculate the empirical formula?

1. Start with the number of grams of each element, given in the problem. 
• If percentages are given, assume that the total mass is 100 grams so that 
the mass of each element = the percent given.
2. Convert the mass of each element to moles using the molar mass from the periodic table. 
3. Divide each mole value by the smallest number of moles calculated. 
4. Round to the nearest whole number.  This is the mole ratio of the elements and is 
represented by subscripts in the empirical formula. 
• If the number is too far to round (x.1 ~ x.9), then multiply each solution by the same 
factor to get the lowest whole number multiple. 
• e.g.  If one solution is 1.5, then multiply each solution in the problem by 2 to get 3. 
• e.g.  If one solution is 1.25, then multiply each solution in the problem by 4 to get 5.

Example;
A compound was found to contain 13.5 g Ca, 10.8 g O, and 0.675 g H.  What is the empirical formula of the compound?

Calculating empirical formula from combustion data
Example : Calculating the empirical formula of a hydrocarbon

The general procedure:

• Calculate the number of grams of carbon in the compound by calculating the number of grams of carbon in the given amount of CO₂.
• Calculate the number of grams of hydrogen in the compound by calculating the number of grams of hydrogen in the given amount of H₂O.
• If the compound contains oxygen, calculate the number of grams of oxygen in it by subtracting the masses of carbon and hydrogen from the given total mass of compound.

? g O  =  (given) g total  -  (calculated) g C  -  (calculated) g H

•  Calculate the empirical formula of the compound from the grams of carbon, hydrogen, and oxygen.

Sample Question: Dianabol is one of the anabolic steroids that has been used by some athletes to increase the size and strength of their muscles.The molecular formula of Dianabol, which consists of carbon, hydrogen, and oxygen, can be determined using the data from two different experiments. In the first experiment, 14.765 g of Dianabol is burned, and 43.257 g CO₂ and 12.395 g H₂O are formed. In the second experiment, the molecular mass of Dianabol is found to be 300.44.  What is the molecular formula for Dianabol?

Solution: Obtained via http://www.mpcfaculty.net/

? g O  =  14.765 g total  - 11.805 g C  -  1.3870 g H  =  1.573 g O

We now calculate the empirical formula.

The empirical formula is C₁₀H₁₄O. We use it to calculate the molecular formula:

Empirical formula mass  =  10(12.011)  +  14(1.00794)  +  1(15.9994)  =  150.22

molecular formula         C₂₀H₂₈O₂

 

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Atomic Orbitals and electronic configurations

Atomic Orbitals

An orbital subdivision of a subshell. A region of space around the nucleus of an atom where there is maximum probability of finding an electron
The following video explains the above concept extremely well:

Electronic configurations

• Not more than two electrons to be placed in one orbital
• If two or more electrons are to be placed in one orbital they must be placed with opposite spin to minimize repulsion

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Ionization Energy

Ionization Energy:

Energy required to remove one mole of electrons from one mole of gaseous atoms

Factors affecting Ionization Energy

• Atomic size
• As atomic size increases the ionization energy increases
• Shielding affect
• Ionization energy decreases as shielding affect increases

Trend of ionization energy in the periodic table:

• Ionization energy decreases down the group

• As we move down the group shell size increases and so the distance from the nucleus increases hence ionization energy decreases
• Also as we move down the group more electrons are being added, hence shielding affect increases.
• Ionization energy increases across the period
• As we move across the period the shielding affect remains constant
• BUT across the period atomic size decreases due to INCREASE in EFFECTIVE NUCLEAR CHARGE which causes the ionization energy to increase

 The following videos will help you understand the trend :

Ionisation Energy - Variation

He > H
One extra proton therefore the
nuclear charge is greater and the
extra electron has gone into the
same energy level.  Increased
attraction makes the electron
harder to remove.

Li < He
Despite the increased nuclear
charge, the outer electron is held
less strongly because it is shielded
by full inner level of electrons and is
further away - easier to remove

Be > Li
Increased nuclear charge plus the
electrons in the same energy level

O < N
Despite the increased nuclear charge the electron is easier to remove.  This is because,
in N the three electrons in the 2p level are in separate orbitals whereas in O two of the
four electrons are in the same orbital.  This leads to repulsion so less energy is needed
for the removal of one of them.

Na < Li
Despite the increased nuclear charge due to the larger number of protons in the nucleus
the increased shielding due to filled inner energy levels coupled with the greater distance
from the nucleus means that the outer electron is held less strongly and easier to remove
ATOMIC NUMBER

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