C4- Chemical Patterns

The periodic table

In the periodic table elements are represented by a symbol and sometimes their names. They are arranged in order of proton number, otherwise known as atomic number. The proton number is the number of positive protons in each atom, it is the bottom number.

Arranging the elements by proton number gives a repeating pattern of propertie: in the table each element is placed beneath those with similar properties.

1 of 31

History of the periodic table

During the 18th century the mass of elements could not be measured, so they were compared against the mass of hydrogen. This is known as relative atomic mass.

Dobereiner- 1817- 'Law of triads', 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.

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.

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.

2 of 31

Data from the table

Relative atomic mass (RAM)- Top number, above the proton number. 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

3 of 31

Data from the table

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.

Periods are the rows of elements on the periodic table.

The elements to the left of the red line are metals, those to the right are non metals. The elements are mainly metals.

4 of 31

Making predictions

Because there are patterns in the way the elements are arranged in the periodic table, it can be used to predict their properties and interpret data. It is possible to predict both physical and chemical properties for elements in the same group.

5 of 31

Group 1 alkali metals- appearance

• 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.
• They are stored in oil to prevent unwanted 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.

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

6 of 31

Alkali metals-physical properties

• The alkali metals have low melting and boiling points compared to most other metals.
• Melting and boiling points decrease down group 1.
• The metals have low densities, compared with others.
• The densities in general increase down the group.
• The alkali metal are very soft.
• They become softer down the group.
7 of 31

Group one reactions with water

The alkali metals react vigorously with cold water. In each reaction hydrogen gas is given off and a metal oxide is produced. The speed and violonce of the reaction increases down the group.

Lithium- Floats, fizzes and and becomes smaller until in disapears.

Sodium- Melts to form a ball that skates on the surface, fizzes rapidly and the hydrogen may burn with an orange flame, before the sodium disapears.

Potassium- Melts, floats, skates around quickly on the surface of the water. The hydrogen ignites instantly. The metal is set on fire and sparks with a purple flame. There is someties a small explosion at the end of the reaction.

The hydroxides formed in all of these reactions dissolve in water to form alkaline solutions. These solutions turn universal indicator purple, showing they are strongly alkaline. Strong alkalis are corrosive, so care must be taken when they are used - for example, by using goggles and gloves.

8 of 31

Group 1- reactions with chlorine

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- If a piece of hot lithium is lowered into a jar of chlorine, white powder is produced and settles on the sides of the jar. This is the salt lithium chloride.

Sodium-If a piece of hot sodium is lowered into a jar of chlorine, the sodium burns with a bright yellow flame. Clouds of white powder are produced and settle on the sides of the jar. This is the salt sodium chloride.

Potassium- Potassium reacts more violently with chlorine than sodium does

9 of 31

The halogens physical properties

The halogens have low melting points and boiling points. The melting points and boiling points increase as you go down the group.

Room temperature is usually taken as being 25°C. At this temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids. There is therefore a trend in state from gas to liquid to solid down the group

The halogens become darker as you go down the group. Fluorine is very pale yellow, chlorine is yellow-green, and bromine is red-brown. Iodine crystals are shiny purple - but easily turn into a dark purple vapour when they are warmed up.

When we can see a trend in the properties of some of the elements in a group, it is possible to predict the properties of other elements in that group. Astatine is below iodine in group 7. The colour of these elements gets darker as you go down the group. Iodine is purple, and, as we would expect, astatine is black.

10 of 31

Reactions of Halogens

The halogens become less reactive as you go down the group. Fluorine, at the top of the group, is the most reactive halogen. It is extremely dangerous, causing severe chemical burns on contact with skin.

The halogens react with metals to make salts called metal halides.

Metal + halogen → metal halide

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.

Because the halogens are very reactive and poisonous, care must be taken when using them. Chlorine is used only in a fume cupboard. Iodine should not be handled (it will damage the skin). Gloves may be used, and goggles should be worn.

11 of 31

Displaement reactions

The reactivity of the halogens decreases as we move down the group. This can be shown by looking at displacement reactions.

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.

12 of 31

Reactivity series

This type of reaction happens with all of the halogens. A more reactive halogen displaces a less reactive halogen from a solution of one of its salts.

If you test different combinations of the halogens and their salts, you can work out a reactivity series for group 7. The most reactive halogen displaces all of the other halogens from solutions of their salts, and is itself displaced by none of the others. The least reactive halogen displaces none of the others, and is itself displaced by all of the others. It works just the same whether you use sodium salts or potassium salts.

13 of 31

State symbols

(s) Solid (l) Liquid (g) Gas

(aq) Aqueous (dissolved in water)

14 of 31

Protons, neutrons and electrons

At the centre of an atom is a nucleus, which contains protons and neutrons. All atoms of the same element have the same number of protons.

Electrons are arranged in shells around the nucleus. In a neutral atom the number of protons is the same as the number of electrons.

The shells can also be referred to as energy levels. The number of the shells and the number of electrons in the outer shell varies from one element to another.

15 of 31

Relative masses and charges

Protons and neutrons have the same mass, which is larger than the mass of an electron.

Protons and electrons have an electric charge, which is the same size for both. Except protons have a positive charge and electrons have a negative charge.

Neutrons have no charge, they are neutral.

16 of 31

Flame colours

When the atoms of some metals are heated they give off coloured light. The colour given off is different for every metal and can be used to identify them.

Lithium- Red flame

Sodium- Yellow flame

Potassium- Lilac flame

17 of 31

Line spectra

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. Some examples of what line spectra look like are shown here:

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.

18 of 31

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.

19 of 31

Dot and cross

The electronic structure of each element can be shown simply as the number of electrons in each shell. For example, lithium is 2.1, neon is 2.8.8, and calcium is 2.8.8.2.

The arrangement of electrons can also be shown using a 'Dot-and-cross' diagram. Electron shells are drawn as circles, with the electrons on each shown as dots or crosses.

20 of 31

Electronic arrangement and group number

Electronic structure- the way electrons are arranged on an atom.

The period number of an atom is the same as the number of occupied shells.

The group number is the same as the number of electrons inthe outer shell.

Group 0 is a partial exception to this rule, since although it comes after group 7 it is not called 'group 8', and it contains helium, which has only two electrons in its outer shell.

21 of 31

Working out electronic structure from the periodic

Here's how to use the periodic table to work out an electronic structure:

1. Find the element in the periodic table. Work out which period it's in, and draw that number of circles around the nucleus.

2. Work out which group the element is in and draw that number of electrons in the outer circle - with eight for group 0 elements (except helium).

3. Fill the other circles with electrons (remember: two in the first, eight in the second and third, then the forth begins to be filled for potassium and calcium.

4. Finally, count your electrons and check that they equal the atomic number.
22 of 31

Working out electronic structure from the atomic n

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.

23 of 31

Electron arrangement and chemical patterns of grou

Elements of group one all have 1 electron in their outer shell. This is why their chemical properties are similar.

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.

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.

24 of 31

Electronic arrangement and chemical patterns of gr

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.

In a reaction with a metal, each halogen atom gains an outer electron and becomes an ion with a single negative charge, -1.

As you go down group 7 the atoms become larger, and the highest occupied energy level (the outer shell) becomes further from the nucleus. The force of attraction between the positively-charged nucleus and a negatively-charged electron from another atom becomes weaker.

As a result, it becomes harder to attract and gain an electron the larger the atom becomes. The more difficult it is to gain these outer electrons, the less reactive a halogen is.

For example chlorine is less reactive than fluorine because the outer electrons in a chlorine atom are further from the nucleus than the outer electrons in a fluorine atom. It is harder for a chlorine atom to gain an electron than it is for a fluorine atom.

Astatine atoms, with 6 shells, are the largest atoms in group 7. They are very unreactive.

25 of 31

Metal ions and non-metal ions

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.

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.

26 of 31

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.

Here are some reactions between metals and non-metals:

• 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.

27 of 31

Dot and cross diagrams for ionic compounds

You need to be able to draw dot-and-cross diagrams to show the ions in some common ionic compounds. For example:

Diagram of bonding in sodium chloride. A sodium atom gives an electron to a chlorine atom. The result is a sodium ion (2,8)+ and a chloride ion (2,8,8)-. Both ions have full outer shells.

Sodium ions have the formula Na+, and chloride ions have the formula Cl. You need to show one sodium ion and one chloride ion.

28 of 31

Properties of ionic compounds

Ionic bonds form when a metal reacts with a non-metal. Metals form positive ions and non-metals form negative ions. Ionic bonds are the electrostatic forces of attraction between oppositely-charged.

Ionic bonds form when a metal reacts with a non-metal. Metals form positive ions and non-metals form negative ions. Ionic bonds are the electrostatic forces of attraction between oppositely-charged ions. The oppositely-charged ions are arranged in a regular way to form giant ionic lattices. Ionic compounds often form crystals as a result. The illustration shows part of a sodium chloride (NaCl) ionic lattice ions.

29 of 31

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.

30 of 31

Working out 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-.

31 of 31