Class 2 How Atoms combine with similar and dissimilar atoms Periodic table Electron configurations Energies Bond types Bond types and properties. Fall 08

Columns represent groups of similar properties, Group 1A, alkali metals; 8, inert gases; 7B halides

Atomic Model Nucleus – Protons +ve charge Neutrons no charge Orbitals – Electrons –ve charge Atoms – protons and electrons balance each other, so no net charge. Ions – Positive or negative charge imbalance Na+ or Cl- Quantum mechanical model Replaced this one

Electron States Each Quantum Number Fills up first- 1, 2, 3,4,5,6 etc Chlorine atom, z=17, 1s2,2s2,2p6,3s2,3p5, electrons = 2+2+6+2+5=17 So 1S2,then 2S2,2p6, then 3S2,3P6,3D10,then 4S etc – there are some Discrepancies though.

Note 4s before 3d Note 4s1 3d5 Note 4s1 3d10

Electron Filling of Orbitals 1s 2s 2p 3s 3p 3d 4s 4p 4d Z = 46 1s 2s 2p 3s 3p 3d 4s 4p 36 1s 2s 2p 3s 3p 3d 4s 30 1s 2s 2p 3s 3p 4s 20 1s 2s 2p 3s 12 1s 2s 2p 10 1s 2s 4 1s 2 On Ionisation, 4s levels empties first.

Filling of Shells Electrons shells like to be full, half full or empty. Full shells – inert status, for example helium Full or empty, between elements by ionic or covalent bonding. Metallic bonding – sharing of electrons.

Periodic Table Rows – number of shells 1,2,3,4 etc Column – number of electrons in outer shell Columns have like properties as number of electrons in outer shell same. Atomic number – number of electrons and so protons in atom. Different number of electrons changes properties and produces elements

Bonding Types Between like atoms – metallic or covalent Between different atoms – metallic, ionic or covalent. Primary – ionic, covalent and metallic. strong Secondary – hydrogen, van der Waals. weak.

Ionisation energies for the hydrogen atom 4s level just below 3d Ionisation energies for the hydrogen atom

Ionic Bonding Na, z=11, # of electrons and protons, 1s2,2s2,2p6,3s1 : 2+2+6+1=11 Note 3s1, lose this electron, full orbitals then ; alkali metals Cl, z=17, 1s2,2s2,2p6,3s2,3p5 : 2+2+6+2+5=17 Note 3p5, gain one electron, full 3p6 orbital, halide ions Na+ ion due to loosing one electron, Cl- due to gaining an electron Ionic bonding between elements with different electronegativities. Elements on right of periodic table highly electronegative, and want to Gain electrons, elements on left electropositive, lose an electron. Directional so brittle type materials. Ordered structure, move one ion Lose order.Crystalline structure from this bonding.

Covalent Bonding Hydrogen 1s1, Carbon 1s2,2s2,2p2, so if hydrogen can share one electron And carbon can share four electrons, both elements will have full orbitals. Four hydrogens share with one carbon in covalent bonds. Molecule of methane gas formed – different elements or same can exhibit Covalent bonds eg carbon.Elements with half full outer orbital.Can be crystalline

Metallic Bonding Ion cores protons And neutrons Electrons shared in a gas cloud as outer orbitals not filled. Lots of empty energy states, such as 3d etc. Can be single element, eg gold, or mixtures such as brass, or aluminum alloys. No directionality to bonds, so ductile and conductive. Crystalline usually, but not ordered.

Metallic Conductivity Electrons in Electrons out Metals are conductive as the electrons are not specific to an ion Applications – electrical wire, copper circuit boards, thermocouples.

Bond Energy and Properties State 25 C Solid Solids Liquid Gases Primary Primary Primary Secondary Secondary Bond energy controls melting point.

Secondary Bonding Hydrogen – needs presence of hydrogen, single electron effect, non crystalline normally Van Der Waals – dipole type bonding, due to slight charge imbalance with distance. Non crystalline normally.

Secondary Bonding Positive charge here as electron Favors carbon side Water good example of hydrogen bond Between hydrogen and oxygen covalent Between molecules, hydrogen bond.

Homework How does calcium and flourine bond to each other? How does aluminum bond to other aluminum atoms and how does silicon bond to other silicon atoms? Why are metals conductive?