Science With Mr. Milstid

So far throughout this unit, we’ve been discussing atoms as uncombined, pure substances - essentially elements. Each elements, as a consequence of its unique combinations of protons, neutrons and electrons has unique characteristic properties. These properties, while unique to individual atoms/elements, occur in very predictable, ordered ways - as sequenced on the periodic table of elements. This table is a fascinating and important tool for scientists; it’s like a guide-book to every element in the universe: it tells you how its atoms are composed, what its atomic mass and number is, what it behaves like, what elements it’s related to, etc.

The Periodic Table
The table as we know it has been under development since 1869, following the first arrangement of known atoms by Dimitri Mendeleev. He arranged atoms by their accepted atomic masses, and began to see patterns emerge as he viewed the elements and their properties side by side. He, and other scientists were also able to use gaps in the table to predict other, yet undiscovered elements.

Since the first arrangement of the periodic table, there have been several advancements and rearrangements of the table based on discoveries of new naturally occurring elements, and laboratory-composed elements. Our modern table now contains all the elements known, arranged by atomic number from left to right.

Arranging the Table
The periodic table is arranged in several ways - the first, and most general is by rows and columns.
Rows = Periods
Each horizontal row of elements (from left to right) on the periodic table is called a period.
The properties of elements changes in similar ways across rows of the periodic table. As you move from left to right, elements become increasingly less reactive, turn to metalloids and non-metals, and end in noble gases.
This common trend of characteristic change across rows is called “periodicity.”

Columns = Groups
Each vertical column of elements (from top to bottom) on the periodic table is called a group.
Elements in the same group have very similar properties because they have the same number of electrons in their outermost shell. Thus Lithium and Rubidium will behave similarly, as will Neon and Argon.

The table is also broken down into categories. In general, there are 3 main categories on the periodic table, each of which may be broken down into subcategories.

1. Metals
2. Non-Metals
3. Metalloids

 
Metals
Most elements in the world are metals. Metals are found to the left of the “zigzag” line on the periodic table (the “B, Si, Ge-As, Sb-Te” metalloid divider).
In general, metals have the following properties: they are shiny, conduct heat and electricity well, malleable & ductile.
Examples of metals include: Iron, Nickel, Cobalt.


The Periodic Breakdown of Metals
Group 1: Alkali Metals
Alkali metals are the most reactive metals. They respond violently to water as they freely join with other elements to make a stable outer shell, often resulting in highly energetic reactions.

 
Group 2: Alkaline Earth Metals
Alkaline-earth metals are less reactive than alkali metals are. Soft, silvery metals that conduct electricity well.
 


 

Group 3-12: Transition Metals
Transition metals include those we are most familiar with: zinc, silver, gold, etc. They tend to be shiny, malleable and conduct thermal energy and electric current well. These metals are able to place up to 32 electrons in their second to last shell, and can use the two outermost shells/orbitals to bond with other elements. It’s a chemical trait that allows them to bond with many elements in a variety of shapes.

Lanthanides and Actinides
Elements in the first row that follow lanthanum are called lanthanides. These are also called rare-earth and inner-transition metals. These can be found naturally (though rarely, considering the name!) on earth.
Elements in the second row that follow actinium are called actinides.
These are are all radioactive and some are not found in nature. Some of the elements with higher atomic numbers have only been made in labs.

 
Poor Metals
The poor metals (post-transition metals) are some metallic elements of the p-block of the periodic table that are more electronegative than the transition metals. Their melting and boiling points are generally lower than those of the transition metals, and they are also softer. They are distinguished from the metalloids, however, by their significantly-greater boiling points in the same row.
These include: Aluminum, Gallium, Indium, Tin, Thallium, Lead, Bismuth, and Polonium.

Non-Metals
Nonmetals are found to the right of the zigzag line on the periodic table (the “B, Si, Ge-As, Sb-Te” metalloid divider). They are a varied group of elements, ranging from reactive solids to inert gases. In general though, they are all dull in appearance, do not conduct heat or electricity well, and are brittle and unmalleable. They are also somewhat rare.
Examples of this group include: Sulfur, Iodine, Neon.


The Periodic Breakdown of Non-Metals
Group VII A: Halogens
These elements are all one electron away from having full outer shells. Because they are so close to being complete, they have freely combine with many different elements (are very reactive). They often bond with metals and elements from Group One of the periodic table, forming Halides.

Group VIIIA: Noble Gases
All inert gases (or noble gases) are located in the far right column of the periodic table, in Group Zero (Group 0) [or Group Eighteen (Group XVIII)]. Elements in this group are unreactive - they do not combine with other elements freely because they have full outer shells with eight electrons (except for Helium, which has 2).

Metalloids
Metalloids, also called semiconductors, are the elements that form the “zigzag” line on the periodic table (the “B, Si, Ge-As, Sb-Te” metalloid divider).
These elements have some properties of both metals and nonmetals.
Some are shiny, some are dull, they are somewhat malleable and ductile, and some conduct heat and electricity. In general, they take on the properties of their group.
Examples of metalloids include: Silicon, Boron, Antimony.



 
An interactive view of the periodic table is below.
Click the image below to explore the properties of individual elements, bonding and practice your skills with atomic number, mass and electron configuration.

We’ve learned so far that all matter is made of atoms. We’ve looked at the history of the science surrounding the atom. Now, we will look at the atom itself - its structure, variations of atoms and how they all fit together.

 


 
Atomic Structure
Atoms are composed of 3 main subatomic particles: Electrons, Protons and Neutrons. There are even smaller subatomic particles, such as quarks and gluons, but for the purposes of understanding the basic composition of matter, these 3 are sufficient.

The Nucleus
At the very center of an atom is the nucleus.
It is a small, dense, positively charged “core” that contains most of the atom’s mass, and is composed of the atom’s protons and neutrons.

Protons
Protons are tiny, positively charged particles found in the nucleus.
Their mass is only about 1.7 x 10^-24g, or 1 atomic mass unit (amu).
Protons are incredibly important to the identity of an atom/element - each elemental atom has a unique number of protons in its nucleus. Boron, for example has 5 protons, while Carbon has 6. Variations on atoms (which we will discuss later as isotopes and ions) always have the same number of protons.
The number of protons an atom has is called its atomic number.

Neutrons
Housed next to the protons are particles called neutrons.
These particles have no charge, and are considered to have about the same mass as a proton. In a general, model atom, neutrons exist in the same number as protons. That is, in a model atom with 5 protons, there will be 5 neutrons. 6 protons, 6 neutrons, and so on.

Electrons
Swirling outside the nucleus in clouds are extremely small, negatively charged particles called electrons.
Compared to protons and neutrons, electrons are very, very small in mass.
It takes more than 1,800 electrons to equal the mass of one proton – in fact, the mass is so small, that we typically describe electrons as having no mass whatsoever.
In a general, model atom, electrons exist in the same number as protons - their charges (negative and positive) work to cancel each other out and provide a net charge of zero (neutral) to the atom. So for instance, if there are 5 protons in a model atom, there will be 5 electrons.

 
In the real world, however, this is not always the case. Sometimes, atoms have an unequal number of electrons and protons, which creates a positively or negatively charged atom.
We call this type of atom an ion.

An atom that has fewer electrons than protons and thus an overall positive charge, is called a cation.
An atom with more electrons than protons and thus an overall negative charge, is called an anion.
Ions are particularly important in chemistry because they allow for the joining of multiple atoms together to form compounds.

 
Occasionally, an atom will have an abnormal amount of neutrons in its nucleus.
This type of atom is called an isotope.

An atom is still the same element if it is missing an electron. The same goes for isotopes - they are still the same element, they are just a little different from every other atom of the same element.
Isotopes are pretty common.
In fact, no elements in the world are made up of purely neutral, model atoms. They are made of a mixture of different isotopes of the same atom.
As such, the atomic mass (the recorded measurement of the mass of protons and neutrons in an atom combine) that appears in the periodic table is an average number. Not all isotopes’ masses will equal this. Some will be higher, some lower.

 
Because there are so many different forms of isotopes in the universe, they get unique isotope names, based on their atomic number, combined with the number of neutrons the atom has.
For example:
A normal, neutral Carbon atom has an atomic number of 6, indicating it would have 6 protons.
In a model atom, the neutrons and protons would be the same, adding up to a total of 12 protons and neutrons.
Its atomic name would be: Carbon-12.

We can use isotopes to calculate the average atomic mass of an element using the percent composition of an atom (which scientists have figured out over years of painstaking research).
The steps to do so are as follows:

1. Multiply the mass number of each isotope by the percentage it exists in an element.
2. Add those numbers together to get a final atomic mass for our isotope.

 
The simulation at this link allows you to build atoms and isotopes of each of the first 10 elements on the periodic table.

 
This simulation asks you to use your understanding of atomic number, mass and electron configuration to build isotopes of an element (but in an altogether different way than the one above).

Here’s a link to the big version.

About 440BC, a Greek scientist named Democritus came up with the idea that eventually, all objects could be reduces to a single particle that could not be reduced any further.
He called this particle an atom, from the Greek word atomos which meant “not able to be divided.”
From this, the idea of the atom – the basic building block of all matter – was born.

While not exactly correct about the size of atoms, or the ability to be subdivided, Democritus discovered one of the most important “things” in science, the atom.
By definition, an atom is the smallest particle into which an element can be divided and still be the same substance. Atoms are the basis of chemistry. They are the basis for everything in the Universe. They are the basis for you! (Obviously, they are a little bit important.)

Our next stop in the history of atoms brings us all the way up to around 1700. By this point, scientists understanding of molecular composition of matter had grown considerably since Democitus’ time; they had figured out that elements combine together in specific ratios to form compounds. One scientist, John Dalton, wanted to know why.
He began to experiment with many different substances, and finally came to the conclusion that compounds come together in specific ratios because they are made of single atoms.
This was a breakthrough in the scientific community’s understanding of the role that atoms played in our world.
In 1803, Dalton came up with a theory about atoms:

1. All substances are made of small particles that can’t be created, divided, or destroyed called atoms.
2. Atoms of the same element are exactly alike, and atoms of different elements are different from each other. (So, atoms of gold are exactly like gold atoms, but different than aluminum atoms).
3. Atoms join with other atoms to make new substances.

For a time, Dalton’s theories were the scientific standard, but, by the end of the 1800’s, scientists had decided that, though much of what he had theorized was correct, new information they could observe about the way matter behaved proved some of Dalton’s theory wrong. So they changed atomic theory, step by step, until they reached the ideas we understand to be true today.

One major change to Dalton’s atomic theory was the discovery of electrons.
In 1897, a British scientist named JJ Thomson discovered small particles inside an atom – which means that atoms can be divided into even smaller parts.
Thomson used something called a cathode-ray tube to discover electrons, which he thought were spread throughout an atom like “plums in a pudding”.



Thus, we consider Thomson’s model of the atom the “plum pudding” model.

 


 
In 1909, a scientist named Ernest Rutherford decided to test the Thomson theory, and designed an experiment to examine the parts of an atom.
He aimed a beam of small, positively charged particles at a sheet of gold foil with a special coating on the back. This coating glowed when hit by the positively charged particles, so Rutherford could see where the particles were going after they hit the gold sheet they were aimed at.
He began with the idea that atoms were soft blobs of matter with electrons spread throughout them, but when he saw the way some of the particles he was shooting at his machine bouncing around in different directions, he realized that atoms must be made differently.
If they were bouncing around and reflecting, then they must be much more dense than he originally thought.

 


 
Following Rutherford’s experiments, Niels Bohr proposed in 1913 that electrons move around the nucleus of an atom in specific paths, on different levels of energy.
He proposed that there were several layers of electrons surrounding a nucleus, all of which travel around in an orbit.



By now, we have come to understand that electrons do not, in fact, travel in specific paths or orbits around the nucleus of an atom.
Instead, they move randomly around the nucleus, with no specific pattern or path, and they do it in a general range of area called an electron cloud (the area where an electron is likely to be found around a nucleus).