Chapter 1: Atomic Structure and the Periodic Table

• Elements


  Pures substances that cannot be seperated chemically
• How to write standard atomic notation?

  Below:

  A)

  
• Periodic table is organised by increasing...

  Atomic number (# of protons)
• Properties of Subatomic Particles:

  Symbol, Location, Charge, Mass

  Below:

  
• What is a stable octet?

  • A full valence shell
  • Can be formed by gaining/losing electrons
• Result of losing/gaining electrons

  Ions

  • Losing: Cation
  • Gaining: Anion
• Naming Ions

  Positive ions keep the same name as the atom

  • sodium > sodium ion

  Negative ions change their ending to "-ide"

  • fluorine > fluoride
• Iso-Electronisity

  • As metals lose electrons/non-mentals gain electrons, they adopt the same electron arrangement as the nearest noble gas
  • An ion that has the same # of electrons as another element is isoelectronic
• Multivalent ions

  • Most transition metals can form multiple possible ions
  • These are called multivalent
• Polyatomic ions

  • Ions that consist of more than one atoms
  • Act like a single ion
  • The ionic charge of a polyatomic ion is shared over the entire ion
• What is the result of the change in proton/electron/neutron?

  Proton: different element

  Electron: formation of ion

  Neutron: formation of isotope
• Isotopes

  
  

  • Atoms of an element that have the same # of protons, but different masses and # of neutrons
  • Very similar chemical properties
• Radioisotopes

  • Unstable isotopes that release energy and subatomic particles to "calm down"
  • These decay and produce nuclei that emit nuclear radiation
  • Dangerous to living cells
• Half Life


  The time it takes of one half of a sample to decay
• Isotopic Abundance


  • Some isotopes are more common than others
  • Atomic Mass = P1M1 + P2M2 + P3M3, etc.
• Difference between carbon-14 and carbon-12

  • Carbon-14 has an atomic mass of 14u and 8 neutrons
  • Carbon-12 has an atomic mass of 12u and 6 neutrons
• Periodic Law

  When the elements are organised into a table according to atomic number, the chemical and physical properties of the elements repeat in a regular, periodic parttern
• Patterns of the period/group in a periodic table

  Period: # of electron shells

  Group: # of valence electrons
• Description of hydrogen

  • Unique element
  • Many physical properties of non-metals
  • Many chemical properties of metals
• Description of alkali metals

  • Reactive, soft, silver-coloured
  • Solid at room temperature
  • Exhibit metallic properties
• Description of alkaline earth metals

  • Light, reactive
  • Solid at room temperature
  • Exhibit metallice properties
• Description of halogens

  • May be solids, liquids, or gases at room temperature
  • Non-lustrous, non-conductors
  • Extremely reactive
• Description of noble gases

  • Gases at room temperature
  • Low melting/boiling point
  • Extremely unreactive
• Transition metals

  • Exhibit a range of chemical and physical properties
  • Form multivalent bonds
  • High melting points
  • Strong, hard metals
• Lanthanides/Actinides

  Lanthanides: Elements 57-71

  Actinides: Elements 89-103
• Atomic Radius Trends

  • As you go down a group, atomic radius increases
  • As you go across a period, the atomic radius decreases
• Why does the atomic radius increase as you go down each group?

  • The number of electron shells increase and are further away from the nucleus
• Why does the atomic radius decrease as you go across a period?

  More protons lead to more attraction of the electrons towards the protons in the nucleus (protons weigh more than electrons).
• How to tell if ions are smaller or larger than their element?

  Cations are smaller (lost their valence shell & force of nucleus is shared among less electrons)

  Anions are larger (force of repulsion amongst electrons increase)
• Ionisation Energy

  The energy required to remove an electron from a neutral atom in the gaseous state

  • 1st ionisation energy = removal of most loosely held electron
  • 2nd ionisation energy = removal of the next most loosely held electron, etc.
• Ionisation Energy Trends

  • As you go down a group, ionisation energy decreases
  • As you go across a period, ionisation energy increases
• Why does ionisation energy decerase as you go down a group?

  They atomic size increases and electrons are further away from the nucleus
• Why does ionisation energy increase as you go across a period?

  The atomic size decreases and more protons are pulling the electrons closer which makes it harder to remove.
• Electron Affinity


  The change in energy that occurs when an electron is added to a neutral atom in the gaseous state

  • Likelihood of gaining an electron
• Electron Affinity Trends

  • As you go down a group, electron affinity decreases
  • As you go across a period, electron affinity increases
• Why does electron affinity decrease as you go down a group?

  The # of shells increase and electrons are further away from the nucleus
• Why does electron affinity increase as you go across a period?

  The electrons are closer to the nucleus so the pull of the protons is greater
• Reactivity Trends for metals & Why

  Increases as you go down a group and as you move right to left

  • Metals tend to give away electrons
  • Larger metals can give away electrons easier
• Reactivity Trends for non-metals & Why

  Increase as you up a group and as you move left to right

  • Non-metals tend to take in electrons
  • Smaller non-metals pull in electrons easily since the positively charged nucleus is very close to the shells