## Inorganic Chemistry Lesson of the Day: 5-Coordinated Complexes

There are 2 common geometries for 5-coordinated complexes:

• Square pyramid: The metal centre is coordinated to 4 ligands in a plane and a 5th ligand above the plane.
• Trigonal bipyramid: The metal centre is coordinated to 3 ligands in a plane and 2 lignads above and below the plane.

## Inorganic Chemistry Lesson of the Day: 2-Coordinated Complexes

Some coordination complexes have just 2 ligands attached to the metal centre.  These complexes have a linear geometry; this allows the greatest separation of the electron clouds in the metal-ligand bonds, which minimizes electron repulsion.

## Inorganic Chemistry Lesson of the Day: 4-Coordinated Complexes

My last lesson stated that the most common coordination number for coordination complexes is 6.  The next most common coordination number is 4, and complexes with this type of coordination adopt either the tetrahedral or the square planar geometry.  The tetrahedron is far more common than the square plane for 4-coordinated complexes, and the type of geometry depends a lot on the size and bonding strength of the ligands.  If the ligands are too big, then a tetrahedral geometry provides greater separation between ligands and minimizes electron repulsion.  If the ligands are too small, then there is room for 2 extra ligands to bond to the metal centre to form a 6-coordinated complex, and an octahedral geometry is adopted instead.

The square planar geometry is usually adopted by 4-coordinated complexes with metal ions that have a d8 electronic configuration.  Examples of such ions include Ni2+, Pd2+, Pt2+, and Au3+.

## Inorganic Chemistry Lesson of the Day: 6-Coordinated Complexes

The most common coordination number for inorganic coordination complexes is 6, and these complexes will most commonly adopt an octahedral geometry.  This geometry is especially common for coordination complexes with a first-row transition metal ion as the Lewis-acid centre.  It consists of 4 ligands forming a plane, and 2 ligands above and below the plane.  The “octa-” prefix in “octahedral” refers to the 8 faces that this geometry has.

Two alternative geometries of 6-coordinated complexes are the trigonal prism and hexagonal plane; these are far less common than the octahedron.

## Inorganic Chemistry Lesson of the Day – Coordination Complexes

A coordination complex is a compound that consists of Lewis bases bonded to a Lewis acid in its centre.  The charge of the complex can be neutral, positive, or negative; if the complex has a positive or a negative charge, then it is called a complex ion.  The Lewis acid is almost always a metal atom or a metal ion.  The Lewis bases are called ligands, and they are often covalently bonded to the Lewis acid.  Common ligands include carbon monoxide, water, and ammonia; what unifies them is the existence of at least one lone pair of electrons in their outermost energy level, and this lone pair of electrons is donated to the Lewis acid.

Some key terminology:

• The donor atom is the atom within the ligand that is attached to the Lewis acid centre.
• The coordination number is the number of donor atoms in the coordination complex.
• The denticity of a ligand is the number of bonds that it forms with the Lewis acid centre.
• If a ligand forms 1 bond with the Lewis acid centre, then it is monodentate (sometimes called unidentate).
• If a ligand forms multiple bonds with the Lewis acid centre, then the coordination complex is polydentate.  For example, a bidentate ligand forms 2 bonds with the Lewis acid centre.

In later Inorganic Chemistry Lessons of the Day, I will only refer to coordination complexes with metal atoms or metal ions as the Lewis acid centres.

## Physical Chemistry Lesson of the Day – Effective Nuclear Charge

Much of chemistry concerns the interactions of the outermost electrons between different chemical species, whether they are atoms or molecules.  The properties of these outermost electrons depends in large part to the charge that the protons in the nucleus exerts on them.  Generally speaking, an atom with more protons exerts a larger positive charge.  However, with the exception of hydrogen, this positive charge is always less than the full nuclear charge.  This is due to the negative charge of the electrons in the inner shells, which partially offsets the positive charge from the nucleus.  Thus, the net charge that the nucleus exerts on the outermost electrons – the effective nuclear charge – is less than the charge that the nucleus would exert if there were no inner electrons between them.

## Physical Chemistry Lesson of the Day – Standard Heats of Formation

The standard heat of formation, ΔHfº, of a chemical is the amount of heat absorbed or released from the formation of 1 mole of that chemical at 25 degrees Celsius and 1 bar from its elements in their standard states.  An element is in its standard state if it is in its most stable form and physical state (solid, liquid or gas) at 25 degrees Celsius and 1 bar.

For example, the standard heat of formation for carbon dioxide involves oxygen and carbon as the reactants.  Oxygen is most stable as O2 gas molecules, whereas carbon is most stable as solid graphite.  (Graphite is more stable than diamond under standard conditions.)

To phrase the definition in another way, the standard heat of formation is a special type of standard heat of reaction; the reaction is the formation of 1 mole of a chemical from its elements in their standard states under standard conditions.  The standard heat of formation is also called the standard enthalpy of formation (even though it really is a change in enthalpy).

By definition, the formation of an element from itself would yield no change in enthalpy, so the standard heat of reaction for all elements is zero.

## Physical Chemistry Lesson of the Day – Hess’s Law

Hess’s law states that the change in enthalpy of a multi-stage chemical reaction is just the sum of the changes of enthalpy of the individual stages.  Thus, if a chemical reaction can be written as a sum of multiple intermediate reactions, then its change in enthalpy can be easily calculated.  This is especially helpful for a reaction whose change in enthalpy is difficult to measure experimentally.

Hess’s law is a consequence of the fact that enthalpy is a state function; the path between the reactants and the products is irrelevant to the change in enthalpy – only the initial and final values matter.  Thus, if there is a path for which the intermediate values of $\Delta H$ are easy to obtain experimentally, then their sum equal the $\Delta H$ for the overall reaction.

## Physical Chemistry Lesson of the Day – The Perpetual Motion Machine

A thermochemical equation is a chemical equation that also shows the standard heat of reaction.  Recall that the value given by ΔHº is only true when the coefficients of the reactants and the products represent the number of moles of the corresponding substances.

The law of conservation of energy ensures that the standard heat of reaction for the reverse reaction of a thermochemical equation is just the forward reaction’s ΔHº multiplied by -1.  Let’s consider a thought experiment to show why this must be the case.

Imagine if a forward reaction is exothermic and has a ΔHº = -150 kJ, and its endothermic reverse reaction has a ΔHº = 100 kJ.  Then, by carrying out the exothermic forward reaction, 150 kJ is released from the reaction.  Out of that released heat, 100 kJ can be used to fuel the reverse reaction, and 50 kJ can be saved as a “profit” for doing something else, such as moving a machine.  This can be done perpetually, and energy can be created forever – of course, this has never been observed to happen, and the law of conservation of energy prevents such a perpetual motion machine from being made.  Thus, the standard heats of reaction for the forward and reverse reactions of the same thermochemical equation have the same magnitudes but opposite signs.

Regardless of how hard the reverse reaction may be to carry out, its ΔHº can still be written.

## Physical Chemistry Lesson of the Day – State Functions vs. Path Functions

Today’s lesson may seem mundane; despite its subtlety, it is actually quite important.  I needed to spend some time to learn it and digest it, and it was time well spent – these concepts are essential for understanding much of thermodynamics.  For brevity, I have not dived into the detailed mathematics of exact differentials, though I highly recommend you to learn it and review the necessary calculus.

Some thermodynamic properties of a system can be described by state variables, while others can be described by path variables.

A state variable is a variable that depends only on the final and initial states of a system and not on the path connecting these states.  Internal energy and enthalpy are examples of state functions.  For example, in a previous post on the First Law of Thermodynamics, I defined the change in internal energy, $\Delta U$, as

$\Delta U = \int_{i}^{f} dU = U_f - U_i$.

State variables can be calculated by exact differentials.

A path variable is a variable that depends on the sequence of steps that takes the system from the initial state to the final state.  This sequence of steps is called the path.  Heat and work are examples of path variables.  Path variables cannot be calculated by exact differentials.  In fact, the following quantities may seem to have plausible interpretations, but they actually do not exist:

• change in heat ($\Delta q$)
• initial heat ($q_i$)
• final heat ($q_f$)
• change in work ($\Delta w$)
• initial work ($w_i$)
• final work ($w_f$)

There is no such thing as heat or work being possessed by a system.  Heat and work can be transferred between the system and the surroundings, but the end result is an increase or decrease in internal energy; neither the system or the surroundings possesses heat or work.

A state/path variable is also often called a state/path function or a state/path quantity.

## Physical Chemistry Lesson of the Day – Standard Heats of Reaction

The change in enthalpy of a chemical reaction indicates how much heat is absorbed or released by the system.  This is valuable information in chemistry, because the exchange in heat affects the reaction conditions and the surroundings, and that needs to be managed and taken into account – in theory, in the laboratory, in industry or in nature in general.

Chemists often want to compare the changes in enthalpy between different reactions.  Since changes in enthalpy depend on both temperature and pressure, we need to control for these 2 confounding variables by using a reference set of temperature and pressure.  This set of conditions is called the standard conditions, and it sets the standard temperature at 298 degrees Kelvin and the standard pressure at 1 bar.  (IUPAC changed the definition of standard pressure from 1 atmosphere to 1 bar in 1982.  The actual difference in pressure between these 2 definitions is very small.)

The standard enthalpy of reaction (or standard heat of reaction) is the change in enthalpy of a chemical reaction under standard conditions; the actual number of moles are specified by the coefficients of the balanced chemical equation.  (Since enthalpy is an extensive property, the same reaction under standard conditions could have different changes in enthalpy with different amounts of the reactants and products.  Thus, the number of moles of the reaction must be standardized somehow when defining the standard enthalpy of reaction.)  The standard enthalpy of reaction has the symbol ΔHº; the º symbol indicates the standard conditions.

## Physical Chemistry Lesson of the Day – Intensive vs. Extensive Properties

An extensive property is a property that depends on the size of the system.  Examples include

An intensive property is a property that does not depend on the size of the system.  Examples include

As you can see, some intensive properties can be derived from extensive properties by dividing an extensive property by the mass, volume, or number of moles of the system.

## Physical Chemistry Lesson of the Day – The Effect of Temperature on Changes in Internal Energy and Enthalpy

When the temperature of a system increases, the kinetic and potential energies of the atoms and molecules in the system increase.  Thus, the internal energy of the system increases, which means that the enthalpy of the system increases – this is true under constant pressure or constant volume.

Recall that the heat capacity of a system is the amount of energy that is required to raise the system’s temperature by 1 degree Kelvin.  Since the heat absorbed by the system in a thermodynamic process is the increase in enthalpy of the system, the heat capacity is just the change in enthalpy divided by the change in temperature.

$C = \Delta H \div \Delta T$.

## Statistics and Chemistry Lesson of the Day – Illustrating Basic Concepts in Experimental Design with the Synthesis of Ammonia

To summarize what we have learned about experimental design in the past few Applied Statistics Lessons of the Day, let’s use an example from physical chemistry to illustrate these basic principles.

Ammonia (NH3) is widely used as a fertilizer in industry.  It is commonly synthesized by the Haber process, which involves a reaction between hydrogen gas and nitrogen gas.

N2 + 3 H2 → 2 NH3   (ΔH = −92.4 kJ·mol−1)

Recall that ΔH is the change in enthalpy.  Under constant pressure (which is the case for most chemical reactions), ΔH is the heat absorbed or released by the system.

## Physical Chemistry Lesson of the Day – The Difference Between Changes in Enthalpy and Changes in Internal Energy

Let’s examine the difference between a change enthalpy and a change in internal energy.  It helps to think of the following 2 scenarios.

• If the chemical reaction releases a gas but occurs at constant volume, then there is no pressure-volume work.  The only way for energy to be transferred between the system and the surroundings is through heat.  An example of a system under constant volume is a bomb calorimeter.  In this case,

$\Delta H = \Delta U + P \Delta V = \Delta U + 0 = q - w + 0 = q - 0 + 0 = q$

This heat is denoted as $q_v$ to indicate that this is heat transferred under constant volume.  In this case, the change in enthalpy is the same as the change in internal energy.

• If the chemical reaction releases a gas and occurs at constant pressure, then energy can be transferred between the system and the surroundings through heat and/or work.  Thus,

$\Delta H = \Delta U + P \Delta V = q - w + P \Delta V = q$

This heat is denoted as $q_p$ to indicate that this is heat transferred under constant pressure.  Thus, as the gas forms inside the cylinder, the piston pushes against the constant pressure that the atmosphere exerts on it.  The total energy released by the chemical reaction allows some energy to be used for the pressure-volume work, with the remaining energy being released via heat.  (Recall that these are the 2 ways for internal energy to be changed according to the First Law of Thermodynamics.)  Thus, the difference between enthalpy and internal energy arises under constant pressure – the difference is the pressure-volume work.

Reactions under constant pressure are often illustrated by a reaction that releases a gas in cylinder with a movable piston, but they are actually quite common.  In fact, in chemistry, reactions under constant pressure are much more common than reactions under constant volume.  Chemical reactions often happen in beakers, flasks or any container open to the constant pressure of the atmosphere.

## Physical Chemistry Lesson of the Day – Heat Capacity

The heat capacity of a system is the amount of heat required to increase the temperature of the system by 1 degree.  Heat is measured in joules (J) in the SI system, and heat capacity is dependent on each substance.  To make heat capacities comparable between substances, molar heat capacity or specific heat capacity are often used.

• Molar heat capacity is the amount of heat required to increase the temperature of 1 mole of a substance by 1 degree.
• Specific heat capacity is the amount of heat required to increase the temperature of 1 gram of a substance by 1 degree.

For example, over the range 0 to 100 degrees Celsius (or 273.15 to 373.15 degrees Kelvin), 4.18 J of heat on average is required to increase the temperature of 1 gram of water by 1 degree Kelvin.  Thus, the average specific heat capacity of water in that temperature range is 4.18 J/(g·K).

## Physical Chemistry Lesson of the Day – Enthalpy

The enthalpy of a system is the system’s internal energy plus the product of the pressure and the volume of the system.

$H = U + PV$.

Just like internal energy, the enthalpy of a system cannot be measured, but a change in enthalpy can be measured.  Suppose that the only type of work that can be performed on the system is pressure-volume work; this is a realistic assumption in many chemical reactions that occur in a beaker, a flask, or any container that is open to the constant pressure of the atmosphere.  Then, the change in enthalpy of a system is the change in internal energy plus the pressure-volume work done on the system.

$\Delta H = \Delta U + P\Delta V$.

## Physical Chemistry Lesson of the Day: Pressure-Volume Work

In chemistry, a common type of work is the expansion or compression of a gas under constant pressure.  Recall from physics that pressure is defined as force applied per unit of area.

$P = F \div A$

$P \times A = F$

Consider a chemical reaction that releases a gas as its product inside a sealed cylinder with a movable piston.

Image from Dpumroy via Wikimedia.

As the gas expands inside the cylinder, it pushes against the piston, and work is done by the system against the surroundings.  The atmospheric pressure on the cylinder remains constant while the cylinder expands, and the volume of the cylinder increases as a result.  The volume of the cylinder at any given point is the area of the piston times the length of the cylinder.  The change in volume is equal to the area of the piston times the distance along which the piston was pushed by the expanding gas.

$w = -P \times \Delta V$

$w = -P \times A \times \Delta L$

$w = -F \times \Delta L$

Note that this last line is just the definition of work under constant force in the same direction as the displacement, multiplied by the negative sign to follow the sign convention in chemistry.

## Physical Chemistry Lesson of the Day – The First Law of Thermodynamics

The change in internal energy of a system is defined to be the internal energy of a system in its final state subtracted by the internal energy of the system in its initial state.

$\Delta U = U_{final} - U_{initial}$.

However, since we cannot measure the internal energy of a system directly at any point in time, how can we calculate the change in internal energy?

The First Law of Thermodynamics states that any change in the internal energy of a system is equal to the heat absorbed the system plus any work done on the system.  Mathematically,

$\Delta U = q + w$.

Recall that I am using the sign convention in chemistry.

The value of $q$ and $w$ can be positive or negative.

• A negative $q$ denotes heat released by the system.
• A negative $w$ denotes work done by the system.

## Physical Chemistry Lesson of the Day – Isolated, Closed and Open Systems

A thermodynamic system can be one of three types:

• An isolated system cannot transfer energy or matter with its surroundings.  Other than the universe itself, an isolated system does not exist in practice.  However, a very well insulated and bounded system with negligible loss of heat is roughly an isolated system, especially when considered within a very short amount of time.  In mathematical and physical modelling, a hydrogen atom’s proton, a hydrogen atom’s electron and a planet are often treated each as an isolated system.
• A closed system can transfer energy (heat or work) but not matter with its surroundings.
• A system that allows work but not heat to be transferred with its surroundings is an adiabatically isolated system.
• A system that allows heat but not work to be transferred with its surroundings is a mechanically isolated system.
• An open system can transfer both energy and matter with its surroundings.