Periodic Table - Proton Number
In the periodic table, the elements are arranged in order of proton number, also called atomic number. This is the number of positive protons in each atom.
Periodic Table - Development .1
During, the period of the 18th century, actual masses could not be measured so they were compared against the mass of hydrogen (the lightest). This is called the relative atomic mass.
In 1817 Johann Döbereiner proposed his 'law of triads'. He realised the relative atomic mass of the middle element in a group of three elements (that had similar properties) was close to the average of the other two elements.
Periodic Table - Development .2
The next significant stage in the development of the periodic table came from John Newlands. He arranged the known elements in order of their atomic masses. He proposed a 'law of octaves', meaning every eighth element had similar properties. This did not work for all the known elements, so was dismissed by the scientific community.
The most significant development of the periodic table was due to Dmitri Mendeleev. He put the known elements in order of relative atomic mass but left gaps for undiscovered elements. He also predicted what the properties of these undiscovered elements might be.
Periodic Table - Relative Atomic Mass
As well as the proton number shown below each element, another number is shown above it. This is the relative atomic mass of the element. It is a comparative measurement of the mass of one atom of the element. You can use it to see how much heavier an atom of one element is compared with an atom of another element.
For example a magnesium atom has a relative atomic mass of 24. So we know it is twice as heavy as a carbon atom, which has a relative atomic mass of 12.
Periodic Table - Groups
Each column in the table contains elements with similar properties, called a group. Each has a group number, shown across the top of the table. So group 1 contains the elements lithium (Li) to francium (Fr), and group 7 contains the elements fluorine (F) to astatine (At).
Reactions of Elements - Group 1 Alkali Metals .1
Lithium, sodium and potassium are all soft metals that are easily cut with a scalpel or knife. The freshly cut surface is a shiny, silver colour, but this tarnishes quickly to a dull grey as the metal reacts with oxygen and water in the air. Pieces of such metals are stored in oil to prevent these reactions.
The shiny surface of sodium tarnishes more quickly than that of lithium. And potassium tarnishes more quickly than sodium. This shows the increasing reactivity of the metals as we go down the group.
Reactions of Elements - Group 1 Alkali Metals .2
Because the alkali metals are so reactive, care has to be taken when using them. They must not be touched because they will react with the water in sweat on the skin. Gloves may be used, and goggles should be worn.
Reactions of Elements - Alkali Metal Reactions
All the alkali metals react vigorously with cold water. In each reaction, hydrogen gas is given off and the metal hydroxide is produced. The speed and violence of the reaction increases as you go down the group. This shows that the reactivity of the alkali metals increases as you go down group 1.
Lithium + water → lithium hydroxide + hydrogen
2Li(s) + 2H2O(l) → 2LiOH(aq) + H2(g)
Sodium + water → sodium hydroxide + hydrogen
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Potassium + water → potassium hydroxide + hydrogen
2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)
Reactions of Elements - Alkali Metal Reactions .Ch
All of the alkali metals react vigorously with chlorine gas. Each reaction produces a white crystalline salt. The reaction gets more violent as you move down group 1, showing how reactivity increases down the group.
Lithium + chlorine → lithium chloride
2Li(s) + Cl2(g) → 2LiCl(s)
Sodium + chlorine → sodium chloride
2Na(s) + Cl2(g) → 2NaCl(s)
Potassium + chlorine → potassium chloride
2K(s) + Cl2(g) → 2KCl(s)
Reactions of Elements - Group 7 Halogens
The halogens have low melting points and boiling points. This is a typical property of non-metals. Fluorine has the lowest melting point and boiling point. The melting points and boiling points then increase as you go down the group.
Reactions of Elements - Group 7 Halogens .2
The halogens react with metals to make salts called metal halides.
Metal + halogen → metal halide
EG. Sodium + chlorine → sodium chloride
2Na(s) + Cl2(g) → 2NaCl(s)
The reaction between sodium and a halogen becomes less vigorous as we move down group 7. Fluorine reacts violently with sodium at room temperature. Chlorine reacts very vigorously when in contact with hot sodium. Iodine reacts slowly with hot sodium.
Reactions of Elements - Uses of Halogens
Halogens are bleaching agents. They will remove the colour of dyes. Chlorine is used to bleach wood pulp to make white paper.
Halogens kill bacteria. Chlorine is added to drinking water at very low concentrations. This kills any harmful bacteria in the water, making it safe to drink. Chlorine is also added to the water in swimming pools.
Reactions of Elements - Displacement
When chlorine (as a gas or dissolved in water) is added to sodium bromide solution the chlorine takes the place of the bromine. Because chlorine is more reactive than bromine, it displaces bromine from sodium bromide. The solution turns brown. This brown colour is the displaced bromine. The chlorine has gone to form sodium chloride.
If you look at the equation, you can see that the Cl and Br have swapped places.
Chlorine + sodium bromide → sodium chloride + bromine
Cl2(aq) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)
Explaining Patterns - Protons, Neutrons and Electr
At the centre of every atom is a nucleus containing protons and neutrons. All atoms of the same element have the same number of protons. This number is used to arrange the elements in the periodic table, beginning with hydrogen, which has just one proton.
Electrons are contained in shells around the nucleus. In a neutral atom the total number of electrons is always the same as the number of protons in the nucleus.
Explaining Patterns - Energy Levels
These shells are also called energy levels. The number of shells, and the number of electrons in the outer shell, varies from one element to another. For example, a lithium atom has two shells, with two electrons in the inner shell and one in the outer shell. A carbon atom also has two shells, but with two electrons in the inner shell and four in the outer shell.
Explaining Patterns - Relative Masses and Charges
Protons and neutrons have the same mass, which is much larger than the mass of an electron.
Protons and electrons have an electrical charge. This electrical charge is the same size for both, but protons are positive and electrons are negative.
Neutrons have no electrical charge; they are neutral.
Explaining Patterns - Line Sprectra
All atoms give off light when heated, although sometimes this light is not visible to the human eye. A prism can be used to split this light to form a spectrum, and each element has its own distinctive line spectrum. This technique is known as spectroscopy.
Scientists have used line spectra to discover new elements. In fact, the discovery of some elements, such as rubidium and caesium, was not possible until the development of spectroscopy. The element helium was discovered by studying line spectra emitted by the Sun.
Explaining Patterns - Electron Arrangement
Electrons are arranged in shells at different distances around the nucleus. As we move across each row of the Periodic Table the proton number increases by one for each element. This means the number of electrons also increases by one for each element.Starting from the simplest element, hydrogen, and moving through the elements in order we can see how the electrons fill the shells. The innermost shell (or lowest energy level) of electrons is filled first. This shell can contain a maximum of two electrons.Next, the second shell fills with electrons. This can hold a maximum of eight electrons. When this is filled, electrons go into the third shell, which also holds a maximum of eight electrons. Then the fourth shell begins to fill.
Explaining Patterns - Electronic Structure
The atomic number of an atom is the number of protons it has. This is the same as the number of electrons. If we know the atomic number we can work out the arrangement of the electrons. Fill the shells starting from the smallest and going outward.
For example silicon has atomic number 16. So we have to fill the shells with 16 electrons. That makes 2 in the first (to fill it), 8 in the second shell (to fill that) and 6 left to go into the third shell. So silicon has electronic structure 2.8.6
Explaining Patterns - Group 1
The atoms of the elements in group 1 all have one electron in their highest occupied energy level (the outer shell). This is why their chemical properties are similar.
When you write the electronic structure of an alkali metal, the last number must be a 1. When you draw or complete a diagram showing the electronic structure of an alkali metal, there must only be one dot or cross in the outer circle.
In a reaction with a non-metal, each alkali metal atom loses its outer electron and becomes an ion with a single positive charge, +1.
Explaining Patterns - Larger Atoms = More Reactive
As you go down group 1 the atoms become larger and the outer electron is further from the nucleus. The force of attraction between the positively-charged nucleus and the negatively-charged outer electron becomes weaker, which is why the outer electron is more easily lost.
So potassium is more reactive than lithium because the outer electron of a potassium atom is further from its nucleus than the outer electron of a lithium atom.
Francium atoms, with 7 shells, are the largest atoms in group 1. They are very reactive - now you know why.
Explaining Patterns - Group 7
The atoms of the elements in group 7 (also called the halogens) have seven electrons in their highest occupied energy level (the outer shell). This is why their chemical properties are similar.
When you write the electronic structure of a halogen, the last number must be a 7. When you draw or complete a diagram showing the electronic structure of a halogen, there must be seven dots or crosses in the outer circle.
In a reaction with a metal, each halogen atom gains an outer electron and becomes an ion with a single negative charge, -1.
Group 1&7 Elements - Metal Ions
When a metal reacts with a non-metal, each metal atom loses the electron, or electrons, from its outer shell. The atom loses negative electrons but still has the same number of positive protons, so it has an overall positive charge. It's not an atom now. Instead it is called an ion.
Group 1&7 Elements - Non-Metal Ions
When a metal reacts with a non-metal, each non-metal atom gains the number of electrons needed to fill its outer shell. The atom gains negative electrons, but still has the same number of positive protons, so it becomes an ion with a negative charge
Group 1&7 Elements - Ionic Compounds
When metals react with non-metals, electrons are transferred from the metal atoms to the non-metal atoms, forming ions. The resulting compound is called an ionic compound. EG.
sodium + chlorine → sodium chloride
magnesium + oxygen → magnesium oxide
calcium + chlorine → calcium chloride.
In each of these reactions, the metal atoms give electrons to the non-metal atoms, so that the metal atoms become positive ions and the non-metal atoms become negative ions. There is a strong electrostatic force of attraction between these oppositely-charged ions, called an ionic bond.
Group 1&7 Elements - Properties of Ionic Compounds
High melting and boiling points - ionic bonds are very strong and a lot of energy is needed to break them, so ionic compounds have high melting points and boiling points.
Conductive when liquid - ions are charged particles, but ionic compounds can only conduct electricity if their ions are free to move. So ionic compounds do not conduct electricity when they are solid, but they do conduct electricity when they are dissolved in water or when they are melted.
Group 1&7 Elements - Formulae and Charges
If you know the charges on the ions in an ionic compound, you can work out its formula.
For example, calcium has an ion with two positive charges, Ca2+, and chlorine has an ion with a single negative charge, Cl-. To balance the positive and negative charges, one calcium ion will need to be with two chloride ions, so the formula of calcium chloride is CaCl2.
If you have the formula of an ionic compound and the charge on one of the two ions, you can work out the charge on the other ion.
For example, sodium oxide has the formula Na2O, and the charge on a sodium ion, Na+, is +1. To balance up the charges from two sodium ions, the oxygen ion must have two negative charges, O2-.