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Chemistry

Home » Student Resources » Online Chemistry Textbooks » CH150: Preparatory Chemistry » CH150: Chapter 2 – Atoms and Periodic Table

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CH150: CHAPTER 2 – ATOMS AND PERIODIC TABLE

Chapter 2: Atoms and the Periodic Table

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2.1 Atomic Theory with Historical Perspectives

2.2 Introduction to Elements and the Periodic Table

2.3 Dmitri Mendeleev and the development of the periodic table

2.4 Families of the Periodic Table

2.5 Defining the Atom

Basic Atomic Structure – electrons, neutrons, and protons

2.6 Atomic Number – Protons Determine the Identity of an Element

2.7 Atomic Mass, Isotopes, and Allotropes

2.8 Electronic Structure of Atoms

The Four Electronic Quantum Numbers

Electron Orbital Filling Rules

Electron Configurations and the Periodic Table

Electron Configuration Solitaire

Electron Configuration Shorthand

Electron-Dot Symbols

2.9 Periodic Table Trends

Atomic Size

Electronegativity

Ionization Energy

Metallic and Nonmetallic Character

2.10 Chapter Summary and Homework

2.11 References

2.1 Atomic Theory with Historical Perspectives

What are the smallest building blocks of everyday objects? This is a question that has interested man since the age of the Greek philosophers. Like the ancient Greeks we can perform a simple thought experiment that raises a very important question for modern chemistry: suppose you were given a piece of aluminum foil and asked to cut the foil in half over and over. How long could you continue cutting, assuming that you had no limitations based on your own abilities? Is there a limit on how small matter can be broken up into, or could you infinitely divide matter into smaller and smaller pieces? This argument dates as far back as the Greek philosophers. Most, like Aristotle, argued that matter could be divided infinitely. However, one brilliant philosopher, Democritus, argued that there is a limit. He proposed that the smallest piece that any element (like aluminum) can be divided into and still be recognized as that element is an Atom, a word derived from the Greek word atomos, meaning “indivisible”.

Figure 2.1: Democritus
Photo taken from: Public Domain
Philosophers, like Democritus, based most of their ideas off of thought experiments like the one above instead of actual observations and experimentation. It is for this reason that Democritus’ ideas on atoms were dismissed until 1808, when John Dalton, an English scientist, proposed four fundamental assumptions based upon observations that we call Dalton’s Atomic Theory.

Dalton proposed that:

1. 
   

Matter is made up of tiny particles called atoms

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3. 
   

Atoms cannot be broken into smaller pieces. During a chemical reaction, atoms are rearranged, but they do not break apart, nor are they created or destroyed

4. 
   
5. 
   

All atoms of the same element are identical in mass and other properties

6. 
   
7. 
   

Atoms of different elements differ in mass and other properties

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2.2 Introduction to Elements and the Periodic Table

An element is a substance that cannot be broken down into simpler chemical substances. There are about 90 naturally occurring elements known on Earth. Using technology, scientists have been able to create nearly 30 additional elements that are not readily found in nature. Today, chemistry recognizes a total of 118 elements which are all represented on a standard chart of the elements, called the Periodic Table of Elements. Each element is represented by a one or two letter code, where the first letter is always capitalized and, if a second letter is present, it is written in lowercase. For example, the symbol for Hydrogen is H, and the symbol for carbon is C. Some of the elements have seemingly strange letter codes, such as sodium which is Na. These letter codes are derived from latin terminology. For example, the symbol for sodium (Na) is derived from the latin word, natrium, which means sodium carbonate. Elements in the periodic table can be broken up into different general classes based upon similarities in their properties. Going from left to right across the periodic table, the elements can be broken up into metals, metalloids, and nonmetals.
Metals are typically shiny, very dense, and have high melting points. Most metals are ductile (can be drawn out into thin wires), malleable (can be hammered into thin sheets), and good conductors of both heat as well as electricity. All metals are solids at room temperature except for mercury. In chemical reactions, metals easily lose electrons to form positive ions. Examples of metals are silver, gold, and zinc.
Nonmetals are generally brittle, dull, have low melting points, and they are generally poor conductors of heat as well as electricity. In chemical reactions, they tend to gain electrons to form negative ions. Examples of nonmetals are hydrogen, carbon, and nitrogen.
Metalloids have properties of both metals and nonmetals. Metalloids can be shiny or dull. Electricity and heat can travel through metalloids, although not as easily as they can through metals. They are also called semimetals. They are typically semi-conductors, which means that they are elements that conduct electricity better than insulators, but not as well as conductors. They are valuable in the computer chip industry. Examples of metalloids are silicon and boron.

A.

Periodic Table Downloadable PDF Version
Figure 2.2: Periodic Table of the Elements. All of the known chemical elements are arranged in the format of a table. The table has been set up in such a way that the characteristics of each different element can be predicted by their position on the table. (A) On this rendition of the periodic table, you can see that the pink elements on the lefthand side of the table are the metals, while the blue elements on the right are the non-metals (Hydrogen is the only exception to this rule and will be explained in the subsequent sections). The metalloids (also termed semi-metals) occur in a stairstep pattern between the metals and nonmetals and are represented in this diagram by the green elements. (B) Shows the positions of the metals, nonmetals and metalloids on the periodic table. During this chapter, you will learn more about these unique characteristics, called periodic trends.
The elements vary widely in abundance. In the universe as a whole, the most common element is hydrogen (about 90%), followed by helium (most of the remaining 10%). All other elements are present in relatively minuscule amounts, as far as we can detect. On the planet Earth, however, the situation is rather different. Oxygen makes up 46.1% of the mass of Earth’s crust (the relatively thin layer of rock forming Earth’s surface), mostly in combination with other elements, while silicon makes up 28.5%. Hydrogen, the most abundant element in the universe, makes up only 0.14% of Earth’s crust. Table 2.1 lists the relative abundances of elements on Earth as a whole and in Earth’s crust. Table 2.2 lists the relative abundances of elements in the human body. If you compare Table 2.1 and 2.2, you will find disparities between the percentage of each element in the human body and on Earth. Oxygen has the highest percentage in both cases, but carbon, the element with the second highest percentage in the body, is relatively rare on Earth and does not even appear as a separate entry in Table 2.1; carbon is part of the 0.174% representing “other” elements. How does the human body concentrate so many apparently rare elements?
The relative amounts of elements in the body have less to do with their abundances on Earth than with their availability in a form we can assimilate. We obtain oxygen from the air we breathe and the water we drink. We also obtain hydrogen from water. On the other hand, although carbon is present in the atmosphere as carbon dioxide, and about 80% of the atmosphere is nitrogen, we obtain those two elements from the food we eat, not the air we breathe.


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2.3 Dmitri Mendeleev and the development of the periodic table

In the 19th century, many previously unknown elements were discovered, and scientists noted that certain sets of elements had similar chemical properties. For example, chlorine, bromine, and iodine react with other elements (such as sodium) to make similar compounds. Likewise, lithium, sodium, and potassium react with other elements (such as oxygen) to make similar compounds. Why is this so?
In 1864, Julius Lothar Meyer, a German chemist, organized the elements by atomic mass and grouped them according to their chemical properties. Later that decade, Dmitri Mendeleev, a Russian chemist, organized all the known elements according to similar properties. He left gaps in his table for what he thought were undiscovered elements, and he made some bold predictions regarding the properties of those undiscovered elements. Later, when elements were discovered whose properties closely matched Mendeleev’s predictions, his version of the table gained favor in the scientific community. Because certain properties of the elements repeat on a regular basis throughout the table (that is, they are periodic), it became known as the periodic table.