C2

• Development of Models - new models of the atom were developed due to new evidence being collected from experiments. Scientists built on each others ideas to create a more accurate picture of the atom.
• Atomic Model Theories:

  • Dalton = All matter is made of atoms. Atoms of elements are different.
  • Thompson = Atom had positive and negative electrons spread out inside, idea of plum pudding model
  • Rutherford = Atom is mostly empty space. It has a positive central nucleus
  • Bohr = Electrons are on fixed energy shells
• Atomic Structure:

  • The centre of the atom is called the nucleus
  • This is made of positively charged protons and neutral neutrons
  • The electrons are negatively charged
  • In a neutral atom there are equal numbers of electrons and protons
• Isotopes - have the same number of protons and electrons but have a different number of neutrons
• Atomic Number - number of protons is equal to the number of electrons

  Mass Number - number of protons and neutrons in the nucleus
• Arrangement:

  • The periodic table is arranged in rows (periods) and columns (groups)
  • They are placed in order of their atomic number
  • Elements in the same group have the same number of electrons in their outer shell
• Mendeleev:

  • Proposed the first periodic table
  • He grouped elements with similar properties
  • Left gaps for undiscovered elements, new elements support his decision
• Metals:

  • Most elements are metals
  • Found on the left side of the periodic table
  • Properties = reactive, high melting point/boiling point, dense, good conductors of heat and electricity
• Non metals:

  • Found on the right of the periodic table
  • Properties = dull, low melting and boiling points, low density, poor conductors of heat and electricity
  • Most are gas at room temperature
• Group 1 Elements (Alkali Metals)

  • Very reactive with water, chlorine and moist air
  • Produce hydrogen and hydroxide when reacted with water
  • Lose their outer electron very easily
  • Reactivity increases as you go down the group
  • e.g sodium + water → hydrogen + sodium hydroxide
• Group 7 Elements (Halogens)

  • Reactivity decreases as you go down the group
  • React with group 1 elements
  • e.g Potassium + Fluoride → Potassium Fluoride
  • Fluorine - yellow gas
  • Chlorine - green gas
  • Bromine - orange/brown liquid
  • Iodine - grey solid
• Displacement reactions:

  • when a more reactive halogen will remove a less reactive halogen from its compound
  • e.g Sodium iodide + chlorine → sodium chloride + iodine
• Group 0/8 (Nobel Gases):

  • Elements have full outer shells of electrons
  • Unreactive/inhert
  • All gas at room temperature due to very low melting and boiling points
• Electron Configuration:

  • Electrons contained within energy levels (shells)
  • Inner most shells can hold a maximum of 2 electrons
  • Outer shells can hold a maximum of 8 electrons
  • e.g Sodium has the configuration 2,8,1
• Group 1 Elements:

  ↓ Lithium

  ↓ Sodium

  ↓ Potassium

  ↓ Rubidium
• Group 7 Elements:

  ↓ Fluorine

  ↓ Chlorine

  ↓ Bromine

  ↓ Iodine
• Group 0/8:

  ↓ Helium

  ↓ Neon

  ↓ Argon

  ↓ Krypton
• Ionic Compounds:

  • Metal and non-metal elements combine to form ionic compounds
  • Electrons transfer from one element to the other
  • Positive or negative ions form
  • Metals in group 1 lose their outer electron (now it has a complete outer shell)
  • Metals in group 7 gain the electron (now it has a complete outer shell)
• Models of Ionic Bonding:

  • Dot and cross diagrams are used to model ionic bonding and the transfer of electrons
• Properties of Ionic Compounds:

  • High melting and boiling point
  • Strong electromagnetic forces of attraction between positive and negative ions
  • Poor conductors when solid (no free space for ions to flow)
  • Good conductors when molten/dissolved in solution (charged ions can flow)