Chemistry

Development of the periodic table

Imagine trying to understand the chemical elements:

-       - Without knowing much about atoms

-       - With some chemical compinds mistakenly through to be elements

-       - Without knowing a complete list of elements.

During the 19^th century, chemists were new elements almost every year. They were also trying very hard to find patterns in the behaviour of the elements. This would allow them to organise the elements and understand more about chemistry.  One of the first suggestions can from John Dalton. He arranged the elements in order of their atomic weights, which had been measured in various chemical reactions. In 1808 he published a table of elements in his book. A New System of Chemical Philosophy.

Mendeleev’s breakthrough

In 1869, the Russian chemist Dmitri Mendeleev cracked the problem. At this time around 50 elements had been identified. Mendeleev arranged all of these in a table. He placed them in the order of their atomic weights. Then he arranged them so that a periodic (regular occurring) pattern in their properties could be seen.

He left gaps for elements that had not yet been discovered. Then he used his table to predict what their properties should be. A few years later, new elements were discovered with properties that closely matched Mendeleev’s predictions. Then there were not many doubts left that his table was a breakthrough in scientific understanding.

Metals, non – metals, and electronic structures

There are different properties of metals and non – metals. The main difference is that metals conduct electricity but non – metals generally are electrical insulators. Notable exceptions are some forms of carbon. In general, metals also have much higher melting points. Comparing solid examples, you find metals are ductile (can be drawn out into wires) and malleable ( can be hammered into shapes without smashing), whereas non – metal solids are brittle.  

Group 0 – the noble gases

The atoms of noble gases have eight electrons in their outermost shell, making the atoms very stable. The exception is the first of the noble gases, helium, which has just two electrons but this complete first shell is also a very stable electronic structure.  

Properties of the alkali metals

All the alkali metals are very reactive. They have to be stored in oil. This stops them reacting with oxygen in the air. Their reactivity increases as you go down the group. So lithium is the least reactive alkali metal and francium is the most reactive.  

Melting points and boiling points

The group 1 metals melt and boil a relatively low temperatures for metals. Going down the group, the melting points and boiling points get lower and lower. In fact, caesium turns into a liquid at just 29°C.  

Reaction with water

When you add lithium, sodium, or potassium to water, the metal floats on the water, moving around and fizzing. The fizzing happens because the metal reacts with the water from the hydrogen gas. Potassium reacts so vigorously with the water that the hydrogen produced ignites. It burns with a lilac flame, coloured by the potassium ions formed in the reaction. The reaction between an alkali metal and water also produces a metal hydroxide. This is why they are called alkali metals. They hydroxides of the alkali metals are all soluble in water. The solution is colourless with a high pH.

Properties of the halogens

The group 7 elements are called halogens. They are a group of toxic non – metals that have coloured vapours. They have fairly typical properties of non – metals.

-     -  They have low melting points and boiling points. Their melting and boiling points increase going down the group.

-     -  They are poor conductors of heat and electricity.

As elements, the halogens all exist as molecules made up of pairs of atoms. These are called diatomic molecules. The atoms in each pair are joined to each other by a covalent bond.  

Reactions of the halogens

The electronic structure of the halogens determines the way they react with other elements. They all have seven electrons in their outermost shell (the highest energy level). So they need to gain just one more electron to achieve the stable electronic structure of a noble gas. When they react with non – metals they gain an extra electron by sharing a pair of electrons with another atom, for example, with hydrogen.  

States of matter

From an early age, anyone can tell the difference between solids, liquids, and gases by using your senses. Later you learn that the majority of substances are classified as solids, liquids, or gasses, and that these are called the three states of matter.

Solids have s fixed shape and volume. They cannot be compressed. Liquids have a fixed volume but they can flow and change their shape. Liquids occupy just slightly more space than when solid (water and ice are exceptions). Gases have no fixed shape or volume. They can be compressed easily.

To explain the properties of solids, liquids, and gases, the particle theory is used. It I based on the fact that all matter is made up of tiny particles and describes;

-        - The movement of the particles,

-        - The average distance between particles.  

Losing electrons to form positive ions

In ionic bonding, the atoms involved loss or gain electrons to form charged particles called ions. The ions have the electronic structure of a noble gas.  So, for example, if sodium from group 1 in the periodic table, loses one electron, it is left with the stable electronic structure of neon.

However, it is also left with one more proton in its nucleus than there are electrons around the nucleus. The proton has a positive charge, so the sodium atom has now become a positively charged ion. The sodium ion has a single positive charge.

You have seen how atoms in group 1 of the periodic table have one electron in their outermost shell (or highest energy level) and form 1+ ions. Group 7 atoms have seven electrons in their outermost shell and form 1- ions. The group number gives the number of electrons in the outermost shell. So how does the group number relate to the charges on the ions formed from atoms? Sometimes the atoms reacting need to gain or lose two electrons to gain the stable electronic structure of a noble gas. An example is when magnesium from group 2 reacts with oxygen from group 6. When these two elements react they form magnesium oxide. This is made up of magnesium ions with a double positive charge, and oxide ions with a double negative charge. So you say that when the atoms form ionic bonds, atoms from

-       

-       Group 1 form 1+ ions

-       Group 2 form 2+ ions

-       Group 3 form 3+ ions, when they form ions as opposed to sharing electrons

-       Group 4 don’t form ions

-       Group 5 form 3- ions, when they form ions as opposed to sharing electrons

-       Group 6 form 1- ions, when they form ions as opposed to sharing electrons

-       Group 7 form 1- ions, when they form ions as opposed to sharing electrons

-       Group 0 never form ions in compounds.

Simple molecules

The atoms of non – metals generally tend to gain electrons to achieve stable electron structures. When they react together, neither atom can give away electrons. So they get the electronic structure of a noble gas by the sharing electrons. The atoms in the molecules are then held together by the shared pairs of electrons. These strong bonds between the atoms are called covalent bonds.