Lecture: The Periodic System and Atomic Structure
1. Classification of the vast number of experimental facts that make up the science of chemistry is very important. Otherwise scientists would be handicapped when trying to understand the similarities and differences that commonly occur throughout chemistry.
2. The periodic system was developed to bring some organization to chemistry. It is in the form of a table designed to bring together elements in columns that have similar chemical and physical properties.
3. It is called a periodic system because many of the properties of elements are periodic functions of their atomic numbers.
Early attempts at classification
1. Metals and nonmetals
a. The first attempt.
b. It did help chemists separate elements into these two groups.
2. Atomic weights
a. J. W. Dobereiner (1780 - 1849) was the first of a long series of chemists who recognised a relation between atomic weights and chemical properties. He observed that, in a set of three elements whose chemical properties were similar, the atomic weight of the second member of the "triad" was almost exactly the mean of the atomic weights of the first and third elements. This striking observation attracted much attention, for it seemed to show a numerical law governing chemical behaviour.
3. Law of Octaves
a. After Dobereiner’s ideas were ruled out, an English chemist John Alexander Reina Newlands in 1864 noted that every eighth element showed similar physical and chemical properties, when the elements are placed in the increasing order of their atomic masses. This was called as the Newlands’ law of octaves. The law states that when elements are placed in the increasing order of atomic masses, the properties of the eight element are repeated.
- Newlands arranged the elements then known in the following manner.
Li Be B C N O F
Na Mg Al Si P S Cl
K Ca
- Each row of elements had seven elements and the eighth fell under the first element. In those days, the number of elements known was very limited and no elements from the noble or inert gas elements such as helium (He), neon (Ne), argon (Ar), etc. were known.
- First let us see the elements in the first column. Li is the first element. The eighth element after Li is Na. Similarly, the eighth element after Na is K. So from the Newlands' law of octaves, we should expect the elements Li, Na and K to have similar chemical properties. This they do have. All the elements are metallic, highly reactive and show a valence of +1. They are known as alkali elements.
- Next, if we take beryllium (Be) as the first element, the eighth element from Be is magnesium (Mg). If we continue in the similar fashion, the eighth element after Mg is calcium Ca. According to Newlands’ law, the elements Be, Mg and Ca should display similar chemical and physical properties. They do. The elements Be, Mg, Ca fall under the group of alkali-earth metals. All these elements are metallic in nature, their oxides are alkaline in nature and they have a valence of +2.
- Now look at another vertical column that has carbon (C) as the first element. The eighth element from C is silicon (Si). It is seen that C and Si are similar in properties. Both of them show tetra-valency. They are non-metals and form oxides easily. Thus Newlands' law of octaves hold good.
- Similarly let’s see the last group of halogens starting with flourine (F). The eighth element after F is chlorine (Cl). As we have already seen that F and Cl display similar properties. Both of them are highly reactive, when dissolved in water, form acids, and have valence of -1. Thus the Newlands' law of octaves was obeyed.
- Newlands' Law of octaves failed for the following reasons :
1. It was not valid for elements that had atomic masses higher than Ca.
2. When more elements were discovered, such as elements from the noble gases such as He, Ne, Ar, they could not be accommodated in his table.
- But the most important contribution in the process of classification of elements was the periodicity he saw in every eighth element. The modern periodic table, drew heavily from the concept of periods of eight. Also it must be noted that Dobereiner’s triads occurred in the octaves of Newlands.
4. Atomic weights and periodic variations
a. When scientists arranged the elements in order according to their atomic weights in a chart with rows and columns, similar elements occupied the columns.
5. The Periodic Law
a. Newly discovered elements fitted neatly into the gaps left in the table leading to acceptance of the “periodic law”.
b. When atomic numbers are used in lieu of atomic weights, the table is even more accurate.
c. The Law states that “The physical and chemical properties of the elements are periodic functions of their atomic numbers.”
6. The Periodic Table
a. Series or periods make up the 7 horizontal rows.
i. Series 1 is the hydrogen-helium series and is called the very short series.
1. Hydrogen exists as a diatomic molecule.
2. Hydrogen forms mainly covalent compounds with most of the elements and is assigned an oxidation number of +1
3. With highly electropositive metals hydrogen has an oxidation number of -1 and forms ionic hydrides.
4. Helium molecules are monotomic and it has no combining power.
ii. Series 2 contains eight elements and is often called the first short series
1. Lithium through Neon
2. Lithium is very electropositive while fluorine is very electronegative and neon is inert, having a complete outer shell.
3. The next is beryllium and it has metallic characteristics and tends to be a dipositive ion in compounds.
4. Boron is nonmetallic and its oxidation number is usually +3 in compounds.
5. Carbon and nitrogen are both nonmetallic.
6. Oxygen is a very electronegative non-metal and combines directly with most of the other elements. It forms ionic compounds with metals and covalent compounds with other nonmetals.
7. From left to right the elements progressively lose metallic properties and gain nonmetallic properties to an inert element, neon.
iii. Series 3 has 8 elements and these elements are referred to as the typical elements.
1. Sodium through chlorine with argon last.
2. Each element has physical and chemical properties similar to the element directly above it.
iv. Series 4 contains 18 elements and is referred to as the first long series.
1. The first 7 elements appear in a regular sequence from Group IA through Group VIIA.
2. From Group IA through Group VIIA in Series 4 through 7 the elements are referred to as the A Group Elements.
3. Group VIII has three elements because they are so similar to each other. These are often referred to as triads and transition triads.
4. Groups IB to VIIB are the B Group Elements.
5. The element groups refer only to Series 4 through 7 but Series 2 and 3 are often included due to property similarities.
6. The transition elements have less of a change in properties as we go through the sequence from left to right than do the regular series.
v. Series 5 mimics Series 4 in properties, etc. and is called the second long series.
vi. Series 6 contains 32 elements and is the first very long series.
1. The first 3 elements mimic the ones above them.
2. The next 14 elements are very much alike in physical and chemical properties.
a. They are similar to lanthanide so are called the lanthanide elements.
b. They are also called the rare-earth elements.
3. The last 15 elements of Series 6 are done the same way as Periods 4 and 5.
vii. Series 7 is referred to as the incomplete series.
1. The elements from actinium through lawrencium are called the actinide elements.
2. The elements show a gradation of properties typical to a series.
3. The first 6 elements exist in nature, the others are manmade.
4. All of the elements in this series demonstrate natural levels of radioactivity.
Groups make up the 16 vertical columns.
When working with Series, we look at the differences across the row. When working with Groups, we look at the similarities down the column.
1. Group IA has the alkali metals and is called the alkali-metal family. Hydrogen is kept separate.
a. Good conductors of electricity, soft enough to cut with a knife, and are very light and electropositive.
b. Melting points are low and decrease as the atomic weight increases.
c. All react with cold water to release hydrogen.
d. They are electropositive and become stronger down the column.
e. They have an oxidation number of +1.
2. Group IB is made up of elements called the coinage metals.
a. They resemble each other closely but are physically very different from Group IA.
b. Oxidation n umber can be +1 in compounds.
c. They have a metallic luster, are hard and malleable.
d. Excellent condustors of heat and electricity.
e. They do not react with water or non-oxidizing acids.
3. Group VIIB is made up of non-metals called the halogens.
a. They are highly electronegative, becoming less so as you go down the column.
b. All exist as covalent molecules, changing from gas to solid as you go down the column.
c. They all have low boiling points and are non-conductors of electricity.
4. Group 0 is a very unique group.
a. The members are all monotomic gases with low boiling points and chemical inertness.
b. The inert gases are very unreactive and the molecules have weak van der Waals forces.
WEEK TWO
Oxidation numbers and position in the Periodic Table
The Theory of Atomic Structure
1. General structure of the atom
a. A small nucleus surrounded by a diffuse electron cloud.
b. The nucleus
i. Proton has an electrical charge of +1 and a mass of 1.
ii. Neutron has no electrical charge and a mass about equal to that of the proton.
c. The electron has an electrical charge of -1 and a mass 1/1850 the mass of the proton.
d. The number of electrons = the number of protons.
e. The atomic weight equals the number of protons plus the number of neutrons since the electrons weight so little.
f. Each element has a unique charge in its nucleus and so must have a unique number of electrons in the electron cloud.
g. Chemical interactions involve interactions between electrons of different atoms.
Electronic Structure of the Atom
The electron cloud is made up of shells or energy levels. All the electrons in any particular shell possess virtually the same energy and are acted upon by the positive nucleus with nearly equal force. The energy of the electrons becomes greater as the distance between the shell and the nucleus increases. The electrons in the outer shells are said to be at higher energy levels.
Principle Quantum Numbers
The electrons in the various shells are characterized in terms of principle quantum numbers. The principle quantum number is represented by the lettr n, which has values of 1,2,3,4, etc.
1. There are 7 possible energy levels in the electron cloud: n = 1, n =2, n = 3, etc.
2. Each shell is also designated with a letter K, L, M, N, O, P, and Q.
3. Each energy-level shell is divided into sub-levels or sub-shells.
a. The total number of possible sub-levels in a shell is equal to the quantum number of the shell.
b. The K (n=1) shell has only 1 sub-level which is designated as s. This s sub-level contains only one orbital and is called the 1s orbital. (2 e maximum)
c. The L (n=2) shell has two sub-levels which are designated as s and p. The s is lower in energy than the p and has only one orbital called the 2s orbital. The p sub-level has three orbitals each of which is called a 2p orbital (2p, 2p, 2p). (8 e maximum)
d. The M (n=3) shell has 3 sub-levels: s, p, and d. The s sub-shell has the lowest energy and 1 orbital (3s). The p sub-shell is next and has 3 orbitals (3p, 3p, 3p). The d sub-shell has the highest energy and 5 orbitals (each labeled 3d). (18 e maximum)
e. The N (n=4) shell has 4 sub-levels: s, p, d, and f. The s sub-shell has the lowest energy and 1 orbital (4s). The p sub-shell is next and has 3 orbitals (4p, 4p, 4p). The d sub-shell has heigher energy and 5 orbitals (each labeled 4d). The f shell has 7 orbitals (each labled 4f). (32 e maximum: 4s2, 4p6, 4d10, 4f14 – coefficient is the principle quantum number and each superscript is the total number of electrons in each orbital.
4. Each single orbital can contain a maximum of 2 electrons.
5. This system allows the labeling of electrons
a. A 1s electron is in the n=1 shell and lower energy than a 3p electron which is in the n=3 shell.
Distribution of electrons in the Periodic Table
1. The maximum number of electrons that can be contained in a given shell is twice the square of the n value of that shell.
2. The outermost shell of a neutral atom is limited to 8 electrons and the next to the outermost shell is limited to 18 electrons.
3. Series 1 is hydrogen and helium. Hydrogen has 1 electron 1s1 and helium has 2 electrons 1s2. The next electron has to go into a higher shell so a new series begins.
4. Series 2 is lithium through neon. 3Li has 3 electrons 1s2, 2s1. 4Be has 4 electrons 1s2, 2s2. The next element, boron, will have its 5th electron in the 2p sub-shell. This continues until neon fills the n=2 energy level.
5. Series 3 has its outermost electrons in the n=3 energy level. There are only 8 elements in this series because the outermost shell is limited to 8 electrons. 18Ar has 18 electrons 1s2, 2s2, 2p6, 3s2, 3p6. The p level is filled so electrons start to fill the n=4 energy level since an electron in the d sub-orbital would put 9 electrons into the outer shell. A new series begins with potassium 19K.
6. Series 4 has its outermost electrons in the n=4 energy level. There are 18 elements in this series because they work on filling the 4s, 3d, and 4p orbitals. The element 36Kr is the last in the first long series and has 8 in its outermost shell.