Atomic Structure and Bonding Week 2

 Atomic structure and bonding related to properties of materials  Nuclide notation. Isotopes and relative atomic mass. Ions. Ionic bonding.  Covalent molecular, covalent network and ionic lattices.  Physical properties of chemicals explained through bonding.  Chemical and ionic formulae including group ions.

 1 st electron shell holds 2 electrons  2 nd electron shell holds 8 electrons  3 rd Electron shell holds 8 electrons REMEMBER THE NUCLEUS HAS A POSITIVE CHARGE

Sub- atomic particle Position in atom ChargeMass ProtonIn nucleus 1+1 NeutronIn nucleus No charge 1 ElectronOutside nucleus 1-0

 Is overall Neutral ◦ Because has the same number of positive protons and negative electrons  Different elements have different number of protons  Recognised as an ATOMIC NUMBER

 Unique for every element  Equals the ◦ Number of protons ◦ Number of electrons (for an atom)

 The only two particles which have mass in an atom are ◦ Protons ◦ Neutrons BOTH ARE FOUND IN NUCLEUS

 The only two particles which have mass in an atom are ◦ Protons ◦ Neutrons  MASS NUMBER = ◦ PROTONS + NEUTRONS  NO. OF NEUTRONS = ◦ MASS NUMBER – ATOMIC NUMBER BOTH ARE FOUND IN NUCLEUS

 Number of protons = never changes ◦ Is like the elements DNA, will never change  Number of neutrons = affect mass of atom ◦ Some heavier, some lighter atoms  Number of electrons = affects the mood the element in ◦ Neutral ◦ Positive ◦ Negative

 Note the numbers are written at LHS of symbol  Atomic Number is at the bottom – can be checked using data book if you forget  Mass Number (the bigger number) at the top.

 No. of protons = Atomic Number ◦ = 17  No. of electrons = Atomic Number ◦ = 17  No. of neutrons = Mass No. – Atomic No. ◦ = = 20

 Are atoms with the ◦ Same atomic number ◦ But, Different mass number  i.e. Different number of neutrons

 Most elements exist as a mixture of isotopes

 This is rarely a whole number ◦ WHY?  The relative atomic mass is an average of all the different Isotopes taking into account the proportion of each present.

 Relative Atomic Mass = 36.5  Would expect Average to be = 36  Closer to 37 = Grater proportion of 37 isotope present

 Relative Atomic Mass = 36.0  This tells you that there is exactly 50% of each isotope present as the average lies exactly in the middle.

 Charged particles

 Group 1 – Alkali Metals ◦ Reactive  Group 7 – Halogens ◦ Reactive  Group 8 (0) – Nobel gases ◦ Unreactive

 Atoms held together by bonds  Bonds are formed so elements can achieve a stable electron arrangement ◦ i.e. a full outer shell like the noble gas elements. REMEMBER GROUP 8 ARE UNREACTIVE AND SO NOT BOND

 Non-metals only  Atoms share electrons  Covalent bond held together by ◦ Attraction between the positive nuclei and the negative shared electrons  Molecule = group of atoms held together by COVALENT bonds

 2 elements bonded together  Diatomic Elements ◦ H, N, O, F, Cl, Br, I  Diatomic Compounds ◦ CO, NO

MOLECULARNETWORK  Low Mp and Bp  Because strong covalent bonds not broken  Weak forces BETWEEN molecules broken  High Mp and Bp  Because strong covalent bonds need to be broken and this required a lot of energy.  E.g. Diamond Sand (SiO 2 )

 (usually) soluble in non-aqueous solvents  (usually) insoluble in aqueous solvents(water)  NEVER CONDUCT ELECTRICITY!!!!

Try CH 4 NH 3 HCl CO 2 – very tricky!

 Charged Particle  Metals – lose electrons – form +ve Ions  Non-metals – gain electrons – form –ve Ions MP

 Between metals and non-metals  Bond held together by electrostatic attraction between POSITIVE AND NEGATIVE IONS Crystal Lattice Structure of oppositely charged ions

 Do not conduct when solid as IONS NOT FREE TO MOVE  Conduct when in solution or molten as IONS ARE FREE TO MOVE  Solid at room temperature ◦ Strong forces of attraction that need to be broken  Soluble is aqueous solvent (i.e. water)