Atoms & 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 Harder to Gain an Electron: Because the attraction for a new electron is weaker, it is harder for the atom to gain an additional electron, making the atom less reactive
• 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 Easier to Lose an Electron: Because the hold on the outer electron is weaker, it is easier for the atom to lose this electron, making the atom more reactive.
• Reactions of group 1 (alkali metals)


  • 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
• 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: Used in balloons and airships since it is much less dense than air, so balloons filled with it float upwards
  • Argon: Used inside light bulbs and as a shield gas during welding due to its inertness
  • 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 test colours for metal ions

  • Li+: Red

  Na+: Orange-yellow

  K+: Lilac

  Ca2+: Orange-red

  Ba2+: Green
• s-block

  • Contains a maximum of 2 valence electrons, placed in the s-subshell.
  • Covers all elements in group 1 and 2.
  • If the second outermost shell contains a d-subshell, it is empty.
• p-block

  • Contains between 3 and 8 valence electrons, placed in the p-subshell. 
  • Covers all elements in groups 13-18.
  • If the second outermost shell contains a d-subshell, it is full.
• d-block

  • Always contains 2 valence electrons, placed in the s-subshell.
  • The d-subshell of the second outermost shell contains 1 to 10 electrons. 
  • The first row of the transition elements is known as the first transition series, with the second row being the second transition series and so on.
• f-blocks

  • Contains the lanthanoids and actinoids.
• Halogens


  A group of non-metal elements that are highly reactive, including chlorine (Cl), bromine (Br), iodine (I), and fluorine (F)
• Alkali Metals


  A group of highly reactive metals that are very reactive and often flammable, including lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr)
• Reaction of Halogens with Alkali Metals

  The reaction between halogens and alkali metals to form a crystalline halide salt. This reaction is highly exothermic and releases heat.