Chemistry

Selection of chemistry related illustrations:

Bohr Atom

The Bohr model of the atom, developed in the early twentieth century, was an attempt to explain observations about the way atoms and electrons absorb, retain and release energy. The model assumed that the atom has a structure similar to that of the solar system with the atomic nucleus at the center of the atom and the atom’s electrons move in circular orbits similar to the way planets orbit the sun. The Bohr model represented a great advancement in the understanding of atomic structure and contributed to the development of quantum mechanics.



Above is a Bohr Atom. Click on the grey rings to move the electron from orbital to orbital. Change the number of excitation states the electron has with the slider at the lower left and click on the hidden, visible and comment buttons to toggle information about the atom on and off.

The Bohr atom is popular as a teaching tool because it helps visualize the relationship between energy, electron position and the emission of electromagnetic energy. However, it is important to understand that the planet-like imagery is just a representation. The planetary model is not consistent with modern observations of the relationship between the nucleus of an atom and the electrons associated with that atom.

Exploration of the Bohr model of the atom contributed to the development of a framework for understanding how electrons absorb, and release discrete amounts (quanta) of energy by suggesting that the electrons associated with an atom do not have free range to be anywhere around that atom. Rather, electrons maintain discrete positions around the nucleus. In the Bohr Atom

  • Electrons travel in circular paths around the nucleus of an atom
  • Electrons can exist only in a finite number of orbitals.
  • Each orbital is at a different distances from the nucleus.
  • Electrons in each orbital contain a set quantity of energy.
  • As long as an electron remains in the same orbital, the energy content of that electron remains constant.
  • Electrons can move between orbits by releasing or absorbing energy.

The lowest energy level an electron can occupy is called the ground state. The higher orbitals represent higher excitation states. The higher the excitation state, the more energy the electron contains.

When an electron absorbs energy, it jumps to a higher orbital. An electron in an excited state can release energy and ‘fall’ to a lower state. When it does, the electron releases a photon of electromagnetic energy. The energy contained in that photon corresponds to the difference between the two states the electron moves between. When the electron returns to the ground state, it can no longer release energy, but can absorb quanta of energy and move up to excitation states (higher orbitals).

The number of movements an electron can make depends on the number of excitation states available. In the case of one ground state plus one excitation state, there is only one possible state change. The electron can absorb one quantum of energy and jump up to the excitation state. From that excitation state, the electron can then drop back down, releasing a photon with a predictable amount of energy in the process.

The addition of a second excitation state increases the number of moves possible from one to three. One associated with movement between the ground state and the lower excitation state, and two associated with movement between the ground state and the second excitation state.

The number of possible moves increases as an arithmetic series as the number of excitation states increase. With four excitation state, the number of state changes is 10, which is 4 plus 3 plus 2 plus 1. The Bohr representation of the atom is also makes it possible to visualize movements of electrons from particular states.

In a Bohr atom with six excitation states, an electron can jump from the ground state up to any one of those six states. An electron in the fifth excitation state, can absorb energy and jump up to the highest excitation state, or fall to any one of the five lower energy states, releasing a photon in the process.

It is important to remember that the Bohr atom is not an accurate representation of atomic structure. However, this model helps illustration some basic concepts of energy absorption and release by atoms and their electrons.

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Covalent Bond Energy

Eating, putting gas in a car and throwing a log on a campfire all involve adding energy to a system. In each case, the energy is added in the form of covalent bond*s that hold atoms together in molecules.

Covalent bonds are one of four types of chemical bonds. The other three are ionic bonds, metallic bonds and hydrogen bonds. Each bond type differs in the way atom share electrons. In covalent bonds, two atoms completely share one or more pairs of electrons. These bonds are quite strong.

Covalent bonds form between atoms when the total energy present in the newly formed molecule is lower than the energy present in each of the atoms alone. The lower energy when bonded results from the fact that atoms are more stable when their outer electron shells are full. Atoms can fill their outer shells by sharing electrons with other atoms though the formation of covalent bonds.

There is a symmetrical relationship between the amount of energy released during the formation of a covalent bond the amount of energy needed to break the bond. Note the flow of energy. Breaking covalent bonds requires energy, and covalent bond formation releases energy.

The term used to describe the energy in a system is Gibbs Free Energy. Gibbs Free Energy can be thought of as energy released during bond formation. When released, this energy is free to do other work. This energy is measured as heat using the units joules or calories or kilocalories.

The amount of energy released during molecule formation can be estimated by counting the number and types of bond in a molecule. For example, a methane molecule has one carbon atom bound to four hydrogen atoms via four single carbon-hydrogen covalent bonds. Carbon-hydrogen bonds release 100 kcal/mole of energy when formed, so the total energy needed to break all the bonds in a methane molecule is 100 kcal x 4 or 400 kcal.

n the following illustration, explores the amount of energy associated with covalent bonds in a selection of molecules containing between one and five carbon atoms. You can also test your understanding with covalent bond energy calculation practice problems.



Balancing Chemical Equations

In chemical reactions, sets of compounds interact with each other to form new compounds. Chemists use equations to describe these interactions. Like mathematical equations, chemical equations conform to a set of rules. This allows equations to provide detailed information about a reaction.



A chemical equation can be thought of as a recipe for making a set of chemical compound using other compounds as starting material. A properly formed chemical equation contains:

  • Reactants - the substances transformed by the reaction. These go on the left.
  • Products - the substances formed by the reaction. These go on the right.
  • An arrow separates reactants from products.
  • In reactions with more than one reactant or product, plus signs separate the individual products and reactants from each other.
  • Numbers in front of each compound specify how many of each is required to convert all of the reactants to products.
The chemical equation for the formation of sugar from water and carbon dioxide is:

6CO2 + 6H2O -> C6H12O6 + 6O2

Properly written, the equation obey's the law of conservation of mass*. The law states that the mass of the reactants going into a reaction must be equal to the mass of the products. This mean that nothing can be gained or lost in the process. A chemical reaction involves the rearrangement of atoms between molecules, not the creation or destruction of atoms.

To make sure that the equation conforms to the law of conservation of mass the equation must be balanced. A balanced chemical equation is one where there are the same number of atoms of each element on either side of the equation. This is the significance of the numbers written before each compound in the reaction. Without the proper number of reactants and products, a chemical equation is not a complete representation of the reaction.

The process of balancing an equation involves adding to each side of the equation until there are the same number of atoms of each element present on both sides.

The illustration explores the process of balancing a chemical equation using the oxidation of a number of different hydrocarbons containing between one and five carbon atoms. A video demonstration of the illustration can be found below.

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Boyle's Law

On Earth, matter exists in one of three states: solid, liquid or gas. Matter in each state exhibits distinct characteristics. Gases, for example, do not have a fixed volume* or shape. As a result, gases respond to pressure changes by changing their volume. In other words, gases are compressible. In contrast, liquids and solids are not compressible: their volume does not change in response to changing pressure. This is the reason air-filled spaces in our ears “pop” during airplane takeoff and landings while the liquid filled spaces in our bodies do not. Boyle's Law* describes the relationship between pressure and volume at constant temperature for a fixed mass* (number of molecules) of a gas.

To understand Boyle's law, it helps to visualize the behavior of gas particles (or molecules) in an enclosed space. In a closed container, individual molecules are constantly hitting and bouncing off the container walls. Each time a gas molecule bounces, it imparts a force on the wall.1 In a flexible container such as a balloon, the force of the molecules hitting the inside of the balloon hold the balloon inflated. The force of each impact is small, but the sheer number of collisions create enough force to keep the balloon open.2

Pressure in a closed container changes if

  1. temperature changes
  2. number of molecules increases or decreases
  3. volume changes
Boyle’s Law deals with number 3; the relationship between volume and pressure when the other two remain constant.

According to Boyle’s Law, the amount a gas will compress is proportional to the pressure applied. Its mathematical expression is:

P1V1=P2V2

Where, P1 is the pressure of a quantity of gas with a volume of V1 and P2 is the pressure of the same quantity of gas when it has a volume of V2. This means that if nothing else changes, the volume of a given mass of gas is inversely proportional to pressure it is under. It is a linear relationship. If pressure on a gas doubles, its volume will decrease by 1/2. An alternative expression of the law is:



PV = C

The product of the volume and pressure equals a constant.

The relationship between pressure and volume results from the influence volume has on the rate at which gas molecules collide with the container walls. If the volume decreases, the molecules encounter the container walls more often. This is true even though the speed (temperature) of the individual molecules has not changed. An increased number of collisions equates to higher pressure. Conversely, if volume increases, the rate of collisions and the pressure decreases.

The following illustration shows this relationship with a container of gas with a fixed temperature and number of molecules. Below the illustration is a youtube video demonstrating its use. You can also test your understanding of the concepts with Concept and Calculation practice problems.



Video demonstration:

1This is Newton’s third law of motion which states that when two objects interact, they exert equal and opposite forces on each other. When gas molecules collide with the wall both the wall and the particle experience the force of the impact. 2A typical party balloon has a volume of 10 to 15 liters. Fully inflated it contains approximately 3×1023 molecules of air. At normal* room temperature, these air molecules are moving at about 300 to 500 m/s. At these speeds, each molecule hits the walls of the balloon thousands of times a second.


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Charles' Law

The Ideal Gas* Laws describe the relationship between pressure, volume* and temperature in a fixed mass* of gas. Charles' law describes the relationship between temperature and volume at a constant pressure.

The ability to visualize the behavior of individual gas particles in an enclosed space helps in understanding the mechanism underlying Charles’ Law. The molecules that make up a gas are moving in straight lines until they encounter another molecule, or a wall. When a molecule encounters a wall, it bounces off and moves off in a different direction. When this happens, Newton's Third Law of motion says that both the molecule and the wall will experience a force. In a flexible container such as a balloon, molecules hitting the inside of the of the balloon are what keep the balloon inflated. In a rigid, but adjustable container such as a sealed syringe, the collisions of the moving gas molecules with the syringe walls provide the force that resists efforts to move the syringe plunger.

Increasing the temperature of a volume of gas causes individual gas molecules to move faster. As the molecules move faster, they encounter the walls of the container more often and with more force. In an inflexible container, the more frequent and forceful collisions result in higher pressure. However, if the container volume is adjustable, the volume will increase, and the pressure will remain the same. Charles’ Law is the formal description of this relationship, allowing change in volume to be calculated if the temperature change is known.

The equation describing Charles’ Law is:

V1/T1 = V2/T2

assuming no change in pressure.

The relationship is linear, if the temperature of a volume of gas doubles, the volume doubles. Temperature has to be measured in Kelvin, the absolute* temperature scale, for the relationship to hold.

While Charles’ Law describes the behavior of ideal gases, not real ones, the law does have real world applications. Real gas*es behave in accordance with Charles' Law at temperatures well above the gases condensation point. Closer to the condensation point, the linear relationship does not hold up; volume decreases more rapidly than temperature.

The following illustration explores the relationship between the temperature and volume of an ideal gas in a container that can adjust to allow pressure to remain constant. This is done by observing changes in the volume of a sealed syringe as the temperature varies. A video demonstrating the illustration can be found at the bottom of the post.



Video demonstration:




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