1.2 Atomic Structure & Periodic Table

• Subatomic particles


  Protons, neutrons, electrons

  
• Nucleus


  • Contains most of the mass of an atom and is found in its centre
  • Protons are positively charged and neutrons have no net charge, giving the nucleus an overall positive charge
• Electrons


  • Orbiting the nucleus
  • Smallest subatomic particle
  • Negatively charged

  
• Atoms have no overall charge because the number of protons and electrons are equal
• Atomic number


  The number of protons an atom has dictates what element it is
• Mass number


  The relative mass of the atom, given by the total number of protons and neutrons
• Isotopes


  Atoms of the same element (therefore the same number of protons) can have different numbers of neutrons
• Calculating relative atomic mass


  ((isotope 1 mass x abundance) + (isotope 2 mass x abundance)) ÷ 100
• Carbon has 2 isotopes: carbon-14 with abundance 20% and carbon-12 with abundance 80%. The relative atomic mass of carbon is 12.4
• Periodic table


  • Elements are arranged in order of atomic (proton) number
  • Elements with similar properties are in columns, known as groups
• Group


  Elements in the same group have the same amount of electrons in their outer shell, which gives them similar chemical properties
• Arrangement of elements


  • Metals are found on the left side and centre
  • Non-metals are on the right side
  • Going across the period (row), elements with intermediate properties are found between non-metals and metals
• Electron configurations


  • Electrons are organised in shells around the nucleus
  • Electrons occupy the lowest available energy levels, the shells closest to the central nucleus
  • The lowest energy shells must be filled before adding electrons to other shells
  • The first shell can hold 2 electrons, the next 2 shells can each hold 8 electrons and for the 4th shell you will only ever have to fill up to a final 2 electrons

  
• Electron configuration


  For an atom, the number of protons is equal to the number of electrons so the atomic number tells you how many electrons to put into shells

  
• Electron configurations


  • Be (4 electrons): 2,2
  • P (15 electrons): 2,8,5
  • K (19 electrons): 2,8,8,1
• Relationship between electron configuration and periodic table

  • The group number an atom is in (e.g. group 1) is equal to the number of electrons in its outer shell
  • The period is related to the total number of electron shells occupied

  
• Trends in the periodic table

  • Elements in the same group have similar chemical and physical properties
  • This is best illustrated by groups 1 and 7
• Ion formation of group 1 and group 7

  • Group 1 elements lose one electron to form a +1 ion
  • Group 7 elements gain one electron to form a -1 ion
• Reactivity of group 7

  • Decreases down the group because the number of shells of electrons increases, so down the group the element attracts electrons from other atoms less, so can't react as easily
• Reactivity of group 1

  • Increases down the group because the number of shells of electrons increases, so electrons further from the nucleus are held to it less strongly and are lost more easily, making larger group 1 atoms more reactive
• Reactions of group 1 elements

  • React vigorously with water to create an alkaline solution and hydrogen
  • React with oxygen to create an oxide
  • React with chlorine to form a white precipitate
• Reactions of group 1 elements

  • Lithium: Burns with a strongly red-tinged flame and produces a white solid, fizzes steadily, gradually disappears, white powder is produced and settles on the sides of the container
  • Sodium: Strong orange flame and produces white solid, fizzes rapidly, melts into a ball and disappears quickly, burns with a bright yellow flame, clouds of white powder are produced and settles on the sides of the container
  • Potassium: Large pieces produce lilac flame, smaller ones make solid immediately, ignites with sparks and a lilac flame, disappears very quickly, reaction is even more vigorous than with sodium
• Group 7 - The halogens


  • Similar reactions due to their seven electrons in their outer shell
  • Non-metals-exist as diatomic molecules made of pairs of atoms
  • They react with metals to form ionic compounds in which the halide ion carries a -1 charge
  • They react with non-metals to form covalent compounds, where there is a shared pair of electrons
• Trends in group 7

  • As you go down the group, relative molecular mass, melting point and boiling point all increase
• Reactivity of halogens


  A more reactive halogen (one from higher up group 7) can displace a less reactive one in an aqueous solution of its salt
• Displacement reactions of halogens

  • Chlorine will displace bromine and iodine
  • Bromine will displace iodine but not chlorine
  • Iodine can replace neither chlorine or iodine
• Uses of chlorine


  • Disinfectant and kills bacteria so is used to sterilise drinking water and clean swimming pools
  • Reacts with sodium hydroxide and water to form bleach
  • Used in the manufacture of chemicals including insecticides, PVC (as polymers) and chlorofluorocarbons
• Uses of iodine


  Iodine is an antiseptic so can be used to prevent infection in hospital procedures
• Reactions of halogens with alkali metals

  The halogens all react quickly with alkali metals to form a crystalline halide salt
• Reactions of potassium with halogens


  • 2K + F2 → 2KF (potassium fluoride)
  • 2K + Cl2 → 2KCl (potassium chloride)
  • 2K + Br2 → 2KBr (potassium bromide)
  • 2K + I2 → 2KI (potassium iodide)
• Reactions of halogens with iron wool

  • Fluorine: Cold iron wool reacts almost instantly to form white iron(III) fluoride
  • Chlorine: Reacts vigorously to form an orange-brown precipitate of iron chloride
  • Bromine: Reacts quickly to form a red-brown precipitate of iron bromide. The reaction has to be warmed
  • Iodine: Reacts slowly in iodine vapour to form a grey iron iodide precipitate. The reaction has to be heated strongly
• Group 0 elements


  • Chemically inert compared to other elements
  • Have 8 electrons in their outer shell (except helium, which has 2- but this shell is still full)
  • Their full outer shell makes them unreactive because they are very stable
  • They are monatomic

  
• Uses of group 0 elements

  • Helium: Has a very low density so it is used in balloons and airships
  • Argon: Is very inert and non-flammable so is used inside light bulbs and as a shield gas during welding
  • Neon: Used in advertising signs; it glows when electricity is passed through it and different coloured glows can be created by coating the glass tubing with other chemicals
• Flame tests to identify metal ions

  1. Clean a metal loop by dipping it in hydrochloric acid then holding it in a Bunsen burner blue flame

  2. Dip the loop into the test sample

  3. Hold the loop in a blue Bunsen burner flame and observe the colour
• Flame test colours for metal ions

  • Li+: Red
  • Na+: Orange-yellow
  • K+: Lilac
  • Ca2+: Orange-red
  • Ba2+: Green
• Test with silver nitrate solution to identify halide ions


  1. Add a few drops of dilute nitric acid to the sample to react with any carbonate ions present

  2. Then add a few drops of dilute silver nitrate solution

  3. Observe the colour of any precipitates formed
• Precipitate colours for halide ions

  • Chloride, Cl-: White
  • Bromide, Br-: Cream
  • Iodide, I-: Yellow
• Test for hydrogen gas

  1. Collect some of the gas in an upturned test tube over the reaction mixture

  2. Place a lighted splint into the test tube

  3. A squeaky pop sound means hydrogen gas is present