periodic table

• Atoms


  Extremely small particles that are basic building blocks of ordinary matter
• Atoms can join together to form

  molecules
• Electron configuration


  Summary of where the electrons are found around a nucleus
• Electron configuration


  • 2,8,8,2 rule shows max amount of valence electron can have up to atomic number 20 (calcium)
  • First electron can contain 2 electrons, second contains 8, third contains 8, fourth shell contain 2
  • Electron configuration can change if the element become an ion
  • For example oxygen (16) = 2,8,6 but to have full valence electron it becomes a +2 anion oxygen (18) = 2,8,8
• Periodic table


  Table of chemical elements
• Atomic radii


  The total distance from the nucleus of an atom to the outermost orbital of its electron
• Atomic radii


  • Increases towards the bottom left, this is because there is more shells of the atom, causing the valence electron to be further away from the nucleus
  • This makes the element more reactive as it has less attraction to nucleus due to long distance
  • Decreases towards the top right, this is because there is less shells and there is full valence shells making the attraction strong and the atom stable
  • This makes the element less reactive as there is a powerful attraction towards the valence electron and the nucleus due to short distance
• Valency


  A measure of an atoms ability to combine with other atoms to create molecules or chemical compounds
• Valency


  • Based on the valence electron(s) and valence shell
  • If the valence shell is full with valence electrons, than it will have less reactivity causing it to have lower chance to combine with other atoms as it is already stable
  • If the valence is not full with valence electrons, than it will have more reactivity causing it to have a higher chance to combine with other atoms, as it wants a full valence shell to be stable instead of its current unstable state
  • The longer the atomic radii, the further the valence shell, the more reactive due to weak attraction of nucleus
  • The shorter the atomic radii, the closer the valence shell, the less reactive due to strong attraction to nucleus
  • Example: One atom of nitrogen combines with three atoms of hydrogen to form ammonia gas. So, the valency of nitrogen is 3.
• 1st ionisation energy


  Energy needed to remove the outermost electron from a neutral atom
• 1st ionisation energy


  • The more higher and more right you go on the periodic table, the higher the ionisation energy will be because the valence electron is more close to the nucleus, meaning the attraction is very strong, needing a lot of energy to separate the valence electron from the atom
  • The more lower and more left you go on the periodic table, the lower the ionisation energy will be because the valence electron is more further away from the nucleus, meaning that attraction is more weak, needing less energy to separate the valence electron from the atom
  • Helium has higher 1st ionisation energy, francium has lower 1st ionisation energy
• Electronegativity


  Describes the tendency of an atom or a functional group to attract electrons towards itself
• Electronegativity


  • High electronegativities tend to acquire electrons in chemical reactions, found in upper right corner of periodic table, anions want to gain electrons to have full valence shell
  • Low electronegativities tend to lose electrons in chemical reactions and are found in lower left corner of periodic table, cations want to lose electrons to have a full valence shell
  • Fluorine has highest electronegativity, francium has lowest electronegativity
• group 1

  • Alkali metals, except hydrogen which is reactive non-metal
  • 1 electron in outermost shell
  • +1 cations, soft, highly reactive, relatively low melting/boiling points
  • Hydrogen, lithium, sodium, potassium, rubidium, caesium, francium
• atoms have 3 subatomic particles

  • protons, electrons, neutrons
• charge of protons, neutrons and electrons

  • Protons are positively charged, electrons are negatively charged, and neutrons are neutral charge
• protons, neutrons and electrons are located in

  • Protons are located in the nucleus with the neutrons, while electrons surround the nucleus in rings or electron cloud
• what makes up atomic weight of atom

  • Protons and neutrons make up the atomic weight of an atom, electrons are 1/1864 the amount of protons
• atomic number=

  • number of protons
• atomic weight=

  • protons + neutrons
• no. of protons=

  • no. of electrons
• ions form when

  • there is a different number of protons or electrons in an element
• isotopes form when

  • there is an increase of neutrons in the nucleus, a huge surplus can lead to the atom being radioactive
• atom is what charged

  • neutrally charged due to protons and electrons cancelling out
• atom is held by

  • electrostatic attraction of positively charged nucleus (protons) and and negatively charged shells (electrons)
• group 2

  • Alkaline earth metals 
  • 2 electrons in outermost shell, +2 cations 
  • Low density, low melting/boiling points (except beryllium), shiny, silvery-white
  • Beryllium, magnesium, calcium, strontium, barium, radium 
• group 3

  • Transition metals
  • Soft, hardness increases with atomic number, silver/grey metals, good conductors of electricity, low melting/boiling points, quickly tarnish in air, react with water   
  • Reactivity influenced by their atomic radius and electronegativity 
  • Scandium, Yttrium, lanthanum, actinium 
• group 4

  • Transition metals 
  • 4 valence electrons 
  • Solids at room temperature, most not electronegative, most form covalent bonds, more metallic as you go down periodic table, high melting points 
  • Titanium, zirconium, hafnium, rutherfordium 
• group 13

  • Boron group
  • 3 valence electron, +3 cations 
  • Hard, brittle, low melting point, good electrical/thermal conductors 
  • Boron, aluminium, gallium, indium, thallium 
• group 14

  • Carbon group containing metalloids and metals  
  • 4 electrons in outermost shell 
  • High melting/boiling points, solids, often form covalent bonds, increasing metallic character down the group 
  • Carbon, silicon, germanium, tin, lead 
• group 15

  • Nitrogen group 
  • 5 valence electrons, -3 anions  
  • Descending down group elements change from nonmetal to metalloids then to metals, reactivity decreases down group, solid at room temp (except nitrogen), density increases as we go down 
  • Nitrogen, phosphorus, arsenic, antimony, bismuth 
• group 16

  • Oxygen group/ chalcogens 
  • 6 valence electrons, -2 anions 
  • Reacts with oxygen to shape dioxides/trioxides, first four are nonmetals, melting/boiling points increase as atomic size decreases 
  • Oxygen, sulphur, selenium, tellurium, polonium 
• group 17

  • Halogens 
  • 7 valence electrons, -1 anions 
  • Soluble in water, low melting/boiling point, as solids they are dull, brittle, poor conductors, highly reactive, high electronegativities, nonmetals, form ionic bonds and covalent bonds 
  • Fluorine, chlorine, bromine, iodine, astatine
• group 18

  • Noble gas 
  • Full valence shell, no ions  
  • Odorless, colourless, non-flammable, gases, low chemical reactivity, extremely stable so unlikely to form chemical bonds, low boiling points
  • Helium, neon, argon, krypton, xenon and radon 
• atomic structure can be used to explain why the changes in


  physical and chemical properties of the elements are periodic
• first ionisation energy affected by 3 factors:

  atoms nuclear charge

  atomic radius

  shielding by inner electrons
• law of conservation of mass

  • No mass is gained or lost in a reaction *in a closed system