Properties of bonding Mrs. Kay

Properties of Ionic bonding Determined by their crystalline structures (how the crystals form) Solid at room temperature (no movement) High melting points = strong bonds Very hard and brittle

Molten compounds conduct electricity Solid structure does not conduct electricity because of rigidity and no ionic movement for electricity to pass through.

Strength Depends on the radii and charges on the ions Increasing metallic charge= stronger bonds = highest melting points Which would produce stronger bonds Na+, Ca+2, or Al+3? Na+ < Ca+2 < Al+3

Solubility Most are soluble (can dissolve into ions) in polar solvents (ex: water, ammonia) These solutions do conduct electricity because of mobile ions (electrolytes)

Hydration The process which polar solvent molecules interact with ions in the crystal lattice and cause the ionic crystal to dissolve, releasing ions into solution Water surrounds the ion (ion-dipole interactions)

Properties of simple covalent molecules Covalent molecules exist as s, l, or g Usually soft Evaporate easier than ionic Low melting and boiling points

Do not conduct electricity in liquid or solid state Not soluble in polar solvents, but may be soluble in nonpolar solvents (CCl4 or gasoline) Napthalene (smells like moth balls)

Note: Molecules: any electrically neutral group of atoms that are bonded tightly together to be considered one particle. Ex: Cl2, NH3, H2O Ionic compounds are not molecules!!! NaCl is not one molecule but a crystal lattice structure with attractive forces holding them together.

Metallic Bonding Name 4 Characteristics of a Metallic Bond. What is a Metallic Bond? - A metallic bond occurs in metals. A metal consists of positive ions surrounded by a “sea” of mobile electrons. Good conductors of heat and electricity Great strength Malleable and Ductile Luster This shows what a metallic bond might look like.

Metallic bond Occurs between atoms with low electronegativities Metal atoms pack close together in 3-D, like oranges in a box. Close-packed lattice formation

Many metals have an unfilled outer orbital In an effort to be energy stable, their outer electrons become delocalised amongst all atoms No electron belongs to one atom They move around throughout the piece of metal. Metallic bonds are not ions, but nuclei with moving electrons

Physical Properties Conductivity Delocalised electrons are free to move so when a potential difference is applied they can carry the current along Mobile electrons also mean they can transfer heat well Their interaction with light makes them shiny (lustre)

Malleability The electrons are attracted the nuclei and are moving around constantly. The layers of the metal atoms can easily slide past each other without the need to break the bonds in the metal Gold is extremely malleable that 1 gram can be hammered into a sheet that is only 230 atoms thick (70 nm)

Melting points Related to the energy required to deform (MP) or break (BP) the metallic bond BP requires the cations and its electrons to break away from the others so BP are very high. The greater the amount of valence electrons, the stronger the metallic bond. Gallium can melt in your hand at 29.8 oC, but it boils at 2400 oC!

Alloys Alloying one metal with other metal(s) or non metal(s) often enhances its properties Steel is stronger than pure iron because the carbon prevents the delocalised electrons to move so readily. If too much carbon is added then the metal is brittle. They are generally less malleable and ductile Some alloys are made by melting and mixing two or more metals Bronze = copper and zinc Steel = iron and carbon (usually)

Network Covalent Molecules: Allotropes of carbon elements can exist in two or more different forms because the element's atoms are bonded together in a different manner Carbon has 3 allotrophes Diamond Graphite Fullerenes (C60) Nanotubes Buckminster Fullerene

Diamonds carbon atoms are bonded together in a tetrahedral lattice arrangement (3D framework) Giant covalent structure Very strong, so they require a lot of energy to break them M.P is 3820 K Does NOT conduct electricity 4x harder than any other natural mineral

Graphite has a sheet like structure where the atoms all lie in a plane and are only weakly bonded to the sheets above and below. (2D framework) Much softer, conducts electricity. (delocalised electron) The C-C bonds are still quite strong. Each carbon bonded to 3 other carbon.

Fullerene C60 consists of 60 carbon atoms bonded in the nearly spherical configuration C60 is highly electronegative, meaning that it readily forms compounds Low solubility, low conductivity (greater than diamond, but much lower than graphite) Buckminster Fullerene (photo) made up of hexagon and pentagon carbon formations Also includes nanotubes (cylindrical) Made up of hexagons of carbon