Tuesday, October 25, 2016

The Law of Definite Composition

Looking ahead to chemical reactions, an important concept to understand is The Law of Definite Composition. To understand this, it is a good idea to begin by reviewing Dalton's Atomic Theory (the parts he got right!):

Dalton noticed that all compounds have something in common. No matter how large or small the sample, the ratio of the masses of the elements in the compound is always the same. Dalton's theory was developed based on this observation.

Dalton proposed the theory that all matter is made up of individual particles called atoms which cannot be divided.
    • All elements are composed of atoms.
    • All atoms of the same element have the same mass, and atoms of different elements have different masses.
    • Compounds contain atoms of more than one element.
    • In a particular compound, atoms of different elements always combine in the same ratios.
Dalton compared the mass ratios of elements when they reacted. He found that (for instance) no matter how much sodium he started with, if the mass of the chlorine didn't change, the exact same amount of sodium would be used up in the reaction.

So suppose he had sixty grams of sodium but only needed one mole (22.990 grams)? After the reaction was complete, he would have left over sodium. In fact, he would have exactly 37.01 grams of sodium left over.

The reason behind this is that compounds form in fixed, set ratios. Salt is always in a ratio of one atom of sodium to one atom of chlorine (1:1 ratio of atoms), or 22.99 grams of sodium (Na) to 35.453 grams of chlorine (Cl). Water is always in a ration of two hydrogen (H) atoms to one oxygen (O) atom (2:1 ratio of atoms). 

Every compound is made up of a fixed, specific ratio of the atoms that make it up. And because atoms of any element have the same atomic mass, compounds form in not only fixed ratios of the number of atoms, but also in fixed ratios of masses.

This Law of Definite Composition can be cited as evidence that supports the idea that some pure substances are combined of elements in a definite ratio. 

The Law of Definite Composition means that, when a chemical reaction is going on, when one of the elements is used up, the reaction ends, and any portion of the other element that is left over is—well, it is left over.

Tuesday, October 18, 2016

Subatomic Particle Configuration

As understanding of atoms increased, different experiments resulted in the discovery of the particles that of which an atom is made. Thompson is credited with first realizing that the atom is made of particles, some negatively charged and some positively charged. Rutherford was credited with identifying that the positive charge is located within a central, relatively dense nucleus. Later, in 1932, James Chadwick designed an experiment to show that neutrally charged particles also existed in atoms—neutrons.

To summarize the development of atomic models, the modern atom is understood to be made up of subatomic particles.

Three Types of Subatomic Particles.

Proton — a positively charged subatomic particle that is found in the nucleus of an atom.
Symbol: p+
Relative charge: +1
Relative mass: 1
Actual mass: 1.674 X 10^-24
Location: In the nucleus

Electron — a negatively charged subatomic particle is found in the space outside the nucleus.
Symbol: e-
Relative charge: -1 
Relative mass: 1/1836
Actual mass: 9.11 X 10^-28
Location: In distinct orbitals surrounding the nucleus

Neutron — a neutral subatomic particle that is found in the nucleus of an atom.
Symbol: n
Relative charge: 0
Relative mass: 1
Actual mass: 1.675 X 10^-24
Location: In the nucleus

From the information above, a few conclusions:
  • Protons and neutrons are almost identical in mass.
  • Electrons are about "1/2000" the mass of a proton. (closer to 1/1836)
  • The neutron does not affect the charge of an atom. They just add mass.

It is the unique ratio of the combination of subatomic particles in different atoms that give each element its characteristics. Thought all elements are built out of the same subatomic particles—protons, neutrons, and electrons—the many, many different combinations possible result in the many, many different elements that can be observed in creation.

The basic classification system for elements is the periodic table. The rules that identify exactly what we call a sample of an element is based on the configuration of the atoms which make it up.

Atomic Number
The basic classification is based on the number of protons. The number of protons in an atom is called its atomic number. Atoms of different elements all have different numbers of protons, and thus, different atomic numbers. If an atom has one proton, it is hydrogen. If it has six it is carbon. No matter what else happens, the number of protons determine the atomic number, and the atomic number determines the type of atom.

Atomic Mass
Another very important difference among atoms is its mass. The mass number is the sum of the masses of all the protons and neutrons in the nucleus of an atom. Not all elements have equal numbers of protons and neutrons, so to find out how many neutrons are in an atom:

Number of neutrons = Mass Number - Atomic Number

For example, the Atomic Number of chlorine (Cl) is 17, and its Mass Number is 35.45. Subtracting, it becomes apparent that there are 18 neutrons (generally, you ignore the decimals in this process).

The mass number also tells us how many grams one mole of the elements' atoms would weigh. NOTE: Some elements do not exist as solo atoms—oxygen, chlorine, and others are always paired, so one mole of the gas will be two moles of the atoms. Looking at the periodic table, we can know how many grams of something would be needed in order to have one mole of it. Doing math, we can, therefore, calculate the mass of any fraction of a mole of any element. So 1/10 of a mole of sodium (Na, Mass Number = 22.99) would weigh 2.299 grams. 

It is the number of protons in a atom that determine what element it is. But in some cases, it is possible to add an neutron to the nucleus, which will change its atomic weight. For instance, a neutron could be added to hydrogen, given it a mass of around 2, yet with only one proton, it would still be hydrogen, atomic number 1. Atoms with the same number of protons, but different numbers of neutrons are called isotopes. Isotopes of an element have the same atomic number, but different mass numbers because they have different numbers of neutrons.

Understanding just the three previous ideas, any substance can be classified and fit into the periodic table as either an element or an isotope of an element. Understanding these concepts opens the door to working with elements in vast and complex ways.

Some definitions and content from:

Physical Science Concepts in Action, Pearson

Saturday, October 15, 2016

Atomic Theory

Understanding what things were made of goes back far into history. The ancient Greeks had ideas and debated as long ago as 2500 years.
Democritus believed that all matter consisted of extremely small particles that could not be divided. He called these particles atoms from the greek word transliterated as atomos, which means "uncut" or "indivisible." He thought there were different types of atoms with specific sets of properties. The atoms in liquids, for example, were round and smooth, but the atoms in solids were rough and prickly.

Aristotle did not think there was a limit to the number of times matter could be divided. He proposed that all matter was made of four elements: earth, fire, air, and water. For centuries, most people accepted Aristotle's views on the structure of matter.

But by the 1800s, scientists had enough data from experiments to support a more substantial atomic model of matter. John Dalton, born in 1766, developed one early model.

Dalton's Theory

Dalton noticed that all compounds have something in common. No matter how large or small the sample, the ratio of the masses of the elements in the compound is always the same. Dalton's theory was developed based on this observation.

Dalton proposed the theory that all matter is made up of individual particles called atoms which cannot be divided.
  • All elements are composed of atoms.
  • All atoms of the same element have the same mass, and atoms of different elements have different masses.
  • Compounds contain atoms of more than one element.
  • In a particular compound, atoms of different elements always combine in the same ratios.
According to Dalton's theory, atom were pictured as solid spheres, each one a tiny, solid sphere with a different mass. His theory satisfied what had, up to that point, been observed and was widely accepted. While incomplete, much of Dalton's theory is still useful in modeling how elements combine to form different compounds. 

However, in time, scientists found that not all of Dalton's ideas about atoms were completely correct. His views were not discarded, but instead, they were revised to take into account new discoveries.

Thompson's Model

A scientist, J. J. Thompson (1856-1940) studied atoms by putting a gas between metal plates and applying an electrical charge. No matter what material was used for the charged metal plates, a beam would appear in the gas, and the beam always behaved in the same way.

Thompson was able to conclude that the beam was made of negatively charged particles that had a mass of about 1/2000 that of a hydrogen atom, the lightest of all atoms. Thompson's experiments provided the first evidence that atoms are made of even smaller particles. 

Thompson's model of the atom resulted in what was called the "plum pudding" model of the atom: since an atom is neutrally charged, yet contains some negatively charged particles, it must also contain positively charged particles, and these particles are mixed together and spread throughout the atom.

As with Dalton's theory, Thompson's model fit what had been observed. Scientists briefly used Thompson's model to guide their investigations, but in 1909, a new discovery led to a even more useful model of the atom.

Rutherford's Atomic Theory

After discovering radioactive alpha particles, Rutherford wondered if they would pass through thin sheets of metal, like gold. Based on Thompson's "plum pudding" model, Rutherford believed that the mass and charge of the particles that make up an atom would, at any given point, be unable to stop the alpha particle.

What he discovered was that, although most of the alpha particles passed through without deviation of course, some of them were turned sharply, and some even bounced off the gold foil and reflected back the way they had come. For this to happen, the particles that make up an atom could not be evenly distributed. Thompson's model would have to be adjusted.

According to Rutherford's model, all of an atom's positive charge is concentrated in its nucleus, and it was the collision with the nucleus that caused the alpha particles to deflect and rebound. According to Rutherford, an atom has a dense, positively charged nucleus and electrons move randomly in the space around the nucleus.

Rutherford's model extended what was previously understood by identifying that atoms have a central, relatively dense (compared to the entire volume of an atom) nucleus around which electrons move, but other observations led to continued development. By 1913, Niels Bohr had provided additional insight into how atoms were constructed.

Bohr's Model

Bohr, who had worked with Rutherford, extended Rutherford's insight by examining more closely the electrons. In Bohr's model, electrons move with constant speed in fixed orbits (rather than randomly) around the nucleus—like planets around a sun. Each electron in an atom has a specific amount of energy. Electrons must orbit the nucleus in one of several fixed, specific orbits, and each orbit represents a specific energy level. The first orbital represents the lowest electron energy level, and the other orbitals represent progressively higher and higher energy levels.

It can be compared to stairs. Electrons can exist on any of the different stairs, and they can move between stairs, but they cannot exist between different stairs. Higher stairs represent higher energy levels.

Electrons can be, by increasing their energy, jump to higher levels, or if they give off energy, move to lower levels. An electron in an atom can move from one energy level to another when the atom gains or loses energy.

Bohr's model provided a great deal of insight in to how elements combine. Picturing the atom as a nucleus around which electrons orbit at different specific energy levels opened a vast amount of understanding. However, further discoveries called for additional refinement.

Electron Cloud Model

Evidence following Bohr's work led to the understanding that the electrons do not orbit the nucleus like a planet. While they do exist at specific energy levels and occupy orbitals, their position in the orbital is never 100% certain. They are somewhere in the orbital, but exactly where cannot be known specifically. Each orbital can be, therefore, conceived as an electron cloud. The concept can be considered analogous to the blur of an object spun at the end of a string. At any given moment, it is somewhere in the path, but the eye cannot pinpoint it.

Along the way, the model of the atom changed with each new discovery. The discovery of neutrons as part of the nucleus led to the understanding that the atom is made of positively charged protons, negatively charged electrons, and neutrally charged neutrons and that each element is composed of a specific combination of those sub-atomic particles. 

Ultimately, a model emerged that is very useful in explaining how chemical reactions take place (as well as in explaining many other aspects of chemistry and physics). A major part of this is understanding energy levels and orbitals for the electrons.

Atomic Orbitals

For any element, all electrons must exist in a set, specific orbital—though exactly where in that orbital is a matter of probability, not certainty. Various orbitals have different energy levels and can hold only a certain number of electrons:

Energy Level Number of Orbitals Maximum Electrons
1 1 2
2 4 8
3 9 18
4 16 32

This information directly leads to understanding of chemical reactions. 


Essentials of Atomic Theory



Definitions and content from:

Physical Science Concepts in Action, Pearson

Image from:
Physical Science Concepts in Action, Pearson, p 100