Atomic Structure

• Elements:

  • contain only one type of atom –

  meaning they cannot be broken

  down into simpler substances.

  • have their own symbols and are

  listed in the periodic table

  • are either metals or non-metals
• Compounds:

  • are substances made from two or

  more different elements

  chemically joined (i.e. bonded)

  together.

  • have different properties from the

  elements from which they are

  made

  • are difficult to break back down

  into their elements
• Mixtures:

  • Are substances that are *not*

  chemically joined together

  • Can be easily separated by a

  range of techniques, including:

  ✓ Filtration

  ✓ Evaporation

  ✓ Crystallisation

  ✓ Distillation

  ✓ Fractional distillation

  ✓ Separating funnel

  ✓ chromatography
• John Dalton

  (early 1800s)

  • Used experiments to suggest substances were made up of tiny spheres called atoms, which were the fundamental building blocks of nature.

  • He also suggested that chemical elements each had their own atoms, which differed from others by mass.

  • Discovered the electron by applying high voltages to gasses at low pressure.
• J.J Thomspon - Suggested the ‘plum pudding’ model of atoms – tiny negatively charged electrons embedded in a cloud of positive charge• As atoms are neutral,

  the number electrons

  and positive charge

  must be equal.
• Ernest Rutherford

  (1909)

  Based his suggestions on the Gold Foil / Alpha particle

  Experiment conducted by Geiger and Marsden...

  The positively charged alpha particles were shot at gold

  foil and expected to go straight through,

  but actually scattered (i.e. deflected).

  Suggested that the positive charge

  (protons) are found concentrated in a

  central part of the atom, its nucleus.

  This is the nuclear model of atoms.
• Niel Bohr 1914 - Suggested electrons orbit the

  nucleus at set distances (i.e.

  energy levels)

  Bohr’s theoretical calculations

  agreed with experimental

  observations.
• James Chadwick

  (1932)

  Discovered the

  neutron. This

  supported

  Rutherford’s proposal.
• Atoms are made of protons, neutrons and

  electrons. Protons and neutrons are found in the nucleus and electrons are found

  in energy levels (i.e. shells) around the nucleus.
• neutrons have a relative mass of 1

  
• • Elements are pure substances that contain only one type of

  atom.

  • Atoms have no charge because each contains an equal

  number of protons and electrons.

  ✓ Atoms of the same element have the same number of

  protons in its nucleus (i.e. their atomic number).

  ✓ Are very small, about 0.1 nanometer (1 x 10-10 m)
• An atom’s atomic and mass number can be found

  on the Periodic Table.

  
• • Isotopes are atoms that have the same number of protons (ie.

  atomic number), but a different number of neutrons (i.e. mass

  number). An atom’s mass number = number of protons +

  number of neutrons.
• isotopes also have ✓ Have similar chemical properties (similar reactions)

  ✓ Have different physical properties (i.e. density)
• • Ions are charged particles that result due to there being more

  or less electrons to protons.

  ✓ Negative ions are formed when electrons are gained

  ✓ Positive ions are formed when electrons are lost
• Electrons in an atom are arranged in energy levels

  (aka: shells). The lowest energy level must be

  filled first.

  The number of electrons in the outer shell of

  an atom determines how it reacts.

  
• The Periodic Table is so called because

  of the regularly repeating patterns in

  the properties of the elements.
• John Dalton (1808)

  • Suggested ordering elements

  by atomic weight...

  • ...this led to incomplete

  versions or placed elements in

  inappropriate groups.

  • Isotopes helped to explain

  why using atomic weights to

  order elements didn’t work.
• John Newlands (1864)

  • Suggested ordering elements

  by atomic weight...

  • ...this led to incomplete

  versions or placed elements in

  inappropriate groups.

  • Isotopes helped to explain

  why using atomic weights to

  order elements didn’t work.

  • introduced ‘law of octaves’ when he

  noticed that the properties of every

  eighth element seemed similar.

  • didn’t take into account that new

  elements were being discovered and

  his pattern broke down after

  calcium...

  • ...because of this, his table was not

  accepted.
• Dimitri Mendeleev (1869)

  • ordered elements by atomic

  number and their properties, but...

  • ...crucially left gaps for unknown

  elements, which when discovered

  matched his predictions.

  • This table was accepted by the

  scientific community.
• The Periodic Table lists all the chemical elements,

  organised into groups (columns) and periods (rows).

  
• Elements found in

  a group have the

  same number of

  electrons in the

  outermost shell.

  Elements found in

  a period have the

  same number of shells of electrons.
• Elements are arranged in order of atomic (proton) number.

  Metals are found on the left and bottom of the Periodic Table, while non-metals are found on the right and top. The number of electrons in the outermost shell (highest energy level) of an atom determines its chemical

  properties.
• • Metals tend to lose electrons, forming positive ions.

  • Non-metals tend to gain electrons, forming negative ions.

  • Noble gases (elements in Group 0) have unreactive because of their very stable electron arrangements

  (i.e. they have a full outer shell of electrons).
• You can explain trends in reactivity in

  terms of the attraction between

  electrons in the outermost shell and the

  nucleus.

  
• The electrostatic attraction between outer electrons and the nucleus depends on:

  • The distance between the outermost electrons and the nucleus

  (i.e. atomic radius). The larger the distance, the weaker the attraction.

  • The number of occupied inner shells (energy levels) of electrons.

  These create a ‘shielding’ effect, weakening the electrostatic attraction.

  • The number of protons in the nucleus (i.e. nuclear charge or ‘size’ of positive charge)
• Group 1: Alkali Metals

  • Melting point / boiling point decrease down the group.

  • All react with water to produce a metal hydroxide solution (an alkali) and hydrogen gas.

  2Li + 2H2O → 2LiOH + H2

  • Have one electron in their outermost shell, so...

  • ...easily lose one electron to form 1+ ions and make ionic compounds...

  • ...These compounds are usually white and dissolve in water, showing a colourless solution (i.e. NaCl, table

  salt, in water)
• electron structure affects reactivity

  
• Group 7: Halogens

  • Melting points / boiling points increase down the group.

  • Are poor conductors of heat and electricity

  • Are all toxic and have coloured vapours

  • Exist as diatomic (ie. 2-atom) molecules, e.g. F2, Cl2, etc.

  • Have 7 electrons in their outermost shell, so...

  • ...can gain one electron to form 1- ions and make ionic compounds with metals

  • Can also form covalent compounds by sharing electrons with other non-metals

  • A more reactive halogen can displace a less reactive halogen from a solution of one of its salts.
• The reactivity of halogen decease down the group. This is because the atoms of each element get larger going down the group and this means the outer shell gets further away from the positive nucleus and is shielded by more electrons. The further the the outer shell is from the positive attraction of the nucleus, the harder it is to attract another electron to complete the outer shell. This is why the reactivity of the halogen group decreases going down group 7