3.1-The History of the Atom

The theory of atomic structure has changed over time

• At the start of the 19th century John Dalton described atoms as solid spheres and said that different sphere made up the different elements

• In 1897 J J Thompson concluded from his experiments that atoms weren’t solid spheres.

• His measurements of charge and mass showed that an atom must contain even smaller, negatively charged particles-electrons.

• The ''solid sphere’ idea of atomic structure had to be changed. The new theory was known as the ‘plum pudding model’.

  • The plum pudding model showed the atom as a ball of positive charge with electrons stuck in it

Rutherford showed that the plum pudding model was wrong

• In 1909 Ernest Rutherford and his student Ernest Marden conducted the famous alpha particle scattering experiments.

• They fired positively charged alpha particles at an extremely thin sheet of gold

• From the plum pudding model, they were expecting the particles to pass straight through the sheet or be slightly deflected at most.

• This was because the positive charge of each atom was thought to be very spread out through the ‘pudding’ of the atom.

• But, whilst most of the particles did go straight through the gold sheet, some were deflected more than expected, and a small number were deflected backwards.

  • So the plum pudding model couldn’t be right

• Rutherford came up with an idea to explain his new evidence- the nuclear model of the atom.

• In this, there’s a tiny, positively charged nucleus at the centre, where most of the mass is concentrated. A ‘cloud’ of negative electrons surrounds this nucleus-so most of the atom is empty space

• When alpha particles came near the concentrated, positive charge of the nucleus, they were deflected. If they were fired directly at the nucleus, they were deflected backwards.

  • Otherwise, they passed through the empty space

Bohr’s nuclear model explains a lot

• Scientists realised that electrons in a ‘cloud’ around the nucleus of an atom, as Rutherford described, would be attracted to the nucleus, causing the atom to collapse.

• Niels Bohr’s nuclear model of the atom suggested that all the electrons were contained in shells.

• Bohr proposed that electrons orbit the nucleus in fixed shells and aren’t anywhere in between. Each shell is a fixed distance from the nucleus

• Bohr’s theory of atomic structure was supported by many experiments and it helped to explain lots of other scientists’ observations at the time

Further experiments showed the existence of protons

• Further experiments by Rutherford and others showed that the nucleus can be divided into smaller particles, which each have the same charge as a hydrogen nucleus.

  • These particles were named protons

• About 20 years after scientists had accepted that atoms have nuclei, James Chadwick carried out an experiment which provided evidence for neutral particles in the nucleus which are now called neutrons.

• The discovery of neutrons resulted in a model of the atom which was pretty close to the modern day accepted version, known as the nuclear model.

3.2-Electronic Structure

The fact that electrons occupy ‘shells’ around the nucleus is what causes the whole of chemistry.

Electron shell rules:

• Electrons always occupy shells(sometimes called energy levels).

• The lowest energy levels are always filled first-these are the ones closest to the nucleus

• Only a certain number of electrons are allowed in each shell:

  • 1st shell = 2

  • 2nd shell = 8

  • 3rd shell = 8

• Atoms are much happier when they have full electron shells-like the noble gases in Group 0

• In most atoms, the outer shell is not full and this makes the atom want to react to fill it

Follow the rules to work out electronic structures

• You can easily work out the electronic structures for the first 20 elements of the periodic table(things get a bit more complicated after that)

  • Example: What is the electronic structure of nitrogen?

    • Nitrogen’s atomic number is 7. This means it has 7 protons, so it must have 7 electrons

    • Follow the electron shell rules above. The first shell can only take 2 electrons and the second shell can take a maximum of 8 electrons

    • So the electronic structure for nitrogen must be 2,5

• More examples:

  • Hydrogen = 1

  • Helium = 2

  • Lithium = 3

  • Carbon = 6

  • Neon = 10

  • Calcium = 20

3.3-Development of the Periodic Table

In the early 1800s elements were arranged by atomic mass

• Until quite recently, there were two obvious ways to categorise elements:

  • Their physical and chemical properties

  • Their relative atomic mass

• Remember, scientists had no idea of atomic structure or of protons, neutrons or electrons, so there were no such thing as atomic number to them.

  • It was only in the 20th century after protons and electrons were discovered that it was realised the elements were best arranged in order of atomic number

• Back then, the only thing they could measure was relative atomic mass, and so the known elements were arranged in order of atomic mass.

• When this was done, a periodic pattern was noticed in the properties of the elements.

  • This is where the name ‘periodic table’ comes from

• Early periodic tables were not complete and some elements were placed in the wrong group.

• This is because elements were placed in the order of relative atomic mass and did not take into account their properties.

Dmitri Mendeleev left gaps and predicted new elements

• In 1869, Dmitri Mendeleev overcame some of the problems of early periodic tables by taking 50 known elements and arranging them into his table of elements - with various gaps as shown

• Mendeleev put the elements mainly in order of atomic mass but did switch that order if the properties meant it should be changed.

  • An example of this can be seen with Te and I - iodine actually has a smaller relative atomic mass but is placed after tellurium as it has similar properties to the elements in that group.

• Gaps were left in the table to make sure that elements with similar properties stayed in the same groups.

• Some of these gaps indicated the existence of undiscovered elements and allowed Mendeleev to predict what their properties might be.

• When they were found and they fitted the pattern it helped confirm Mendeleev’s ideas.

  • For example, Mendeleev made really good predictions about the chemical and physical properties of an element he called ekasilicon which we know today as germanium

• The discovery of isotopes in the early 20th century confirmed that Mendeleev was correct to not place elements in a strict order of atomic mass but to also take account of their properties.

• Isotopes of the same elements have different atomic masses but have the same chemical properties so occupy the same position on the periodic table.

3.4-The Modern Periodic Table

The periodic table helps you to see patterns in properties

• There are 100ish elements, which all materials are made of

• In the periodic table the elements are laid out in order of increasing atomic, proton, number.

• Arranging the elements like this means there are repeating patterns in the properties of the elements. (The properties are said to occur periodically, hence the name periodic table)

• If it wasn’t for the periodic table organising everything, you’d have a heck of a job remembering all those properties.

• It’s a handy tool for working out which elements are metals and which are non-metals. Metals are found to the left and non-metals to the right

Formation

• Elements with similar properties form columns

• These vertical columns are called groups

• The group number tells you how many electrons there are in the outer shell.

  • For example, Group 1 elements all have one electron in their outer shell and Group 7 all have seven electrons in their outer shell.

    • The exception to the rule is Group 0, for example Helium has two electrons in its outer shell.

    • This is useful as the way atoms react depends upon the number of of electrons in their outer shell, so all elements in the same group are likely to react in a similar way

• If you know the properties of one elements, you can predict properties of other elements in that group. For example, the Group 1 elements are Li, Na, K, Rb, Cs and Fr. They’re all metals and they react in similar ways

• You can also make predictions about trends in reactivity

  • e.g. In Group 1, the elements react more vigorously as you go down the group, but in Group 7 reactivity decreases as you go down the group

• The rows are called periods. Each period represents another full shell of electrons

3.5-Metals and Non-Metals

Most elements are metals

• Metals are elements which can for positive ions when they react

• They’re towards the bottom and to the left of the periodic table

• Most elements in the periodic table are metals

• Non-metals are at the far right and top of the periodic table

• Non-metals don’t generally form positive ions when they react

The electronic structure of atoms affects how they will react

• Atoms generally react to form a full outer shell

  • They do this via losing, gaining or sharing electrons

• Metals to the left of the periodic table don’t have many electrons to remove

• Metals towards the bottom of the periodic table have outer electrons which are a long way from the nucleus

  • They feel a weaker attraction

  • Both these effects means that not much energy is needed to remove the electrons so it’s feasible for them to either share or gain electrons to get a full outer shell

Metals and non-metals have different physical properties

• All metals have metallic bonding which causes them to have similar basic physical properties

  • They’re strong but can be bent or hammered into different shapes

  • They’re great at conducting heat and electricity

  • They have high boiling and melting points

• As non-metals don’t have metallic bonding, they don’t tend to exhibit the same properties as metals

  • They tend to be dull looking, more brittle, aren’t always solids at room temperature, don’t generally conduct electricity, and often have a lower density

Transition metals can be found between group 2 and group 3

• Transition metals are in the centre of the periodic table

• Transition metals are typical metals and have the properties you would expect of a proper metal

  • They’re good conductors or heat and electricity and they’re very dense, strong and shiny

• Transition metals also have some pretty special properties

  • Transition metals can have more than one iron

    • Copper forms Cu+ and Cu2+ ions

    • Cobalt forms Co2+ and Co3+ ions

  • Transition metal ions are often coloured, and so compounds that contain them are colourful

    • Potassium chromate which yellow and potassium manganate is purple

  • Transition metal compounds often make good catalysts

    • Nickel based catalysts are used in the hydrogenation of alkenes

    • Iron catalyst is used in the haber process for making ammonia