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Properties of bonding

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

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Properties of bonding

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  1. Properties of bonding Mrs. Kay

  2. 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

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

  4. 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

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

  6. 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)

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

  8. 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)

  9. 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.

  10. 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.

  11. 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

  12. 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

  13. 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)

  14. 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)

  15. 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!

  16. 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)

  17. 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

  18. 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

  19. 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.

  20. 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

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