Properties of covalent substances, metals and ionic compounds


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Properties of covalent substances, metals and ionic compounds

  1. 1. Properties of Covalent Substances, Metals, and Ionic Compounds
  2. 2. Molecular substances have a broad range of physical and chemical properties.• Gases are elements such as Noble gases or diatomic nonmetals (N2 , Cl2 , H2 , O2 , F2) or small, nonpolar compounds. Weak intermolecular forces do not let molecules stick to one another to form liquids or solids.• Liquids are molecular compounds with intermediate-strength intermolecular forces. Molecules are slightly polar. They stick together but not in rigid structures.
  3. 3. • Solids are of 4 major types:1) Ionic compounds – positive & negative ions arranged in repeating patterns. SALTS!!!! Held together very strongly in 1 big structure called a crystal lattice. Hard, brittle solids with high melting points. Many are soluble in water because water is polar. NO MOLECULES!Lots more info later! 
  4. 4. 2) Covalent network crystals - no separate molecules. Atoms are bonded to all surrounding atoms with covalent bonds. Very hard & brittle, insoluble in water. Diamond, graphite and quartz are examples.
  5. 5. Metallic solids – layers of metal atoms that all share a “sea of valence electrons”. Valence electrons move FREELY between layers of atoms. These delocalized electrons cause metallic properties, especially electrical conductivity. amounts of energy can be absorbed and Very small released by these electrons:this makes metals…shiny!
  6. 6. Metallic PropertiesBonds are ALWAYS occurs between METAL atoms (TripleDUH!)Malleability – can be hammered into shapesLuster –shiny!Ductility – can be stretched into a wireMetals have a broad range of melting points, goodconductors of heat &electricity.
  7. 7. 4) Covalent molecular compounds – made of individual molecules. (molecules are groups of atoms held together by covalent bonds with specific numbers of atoms and in specific geometric arrangements) Properties vary broadly depending on how POLAR the molecules are. INTERMOLECULAR FORCES hold molecules near one another. The strength of these forces determines the physical properties of the substance.London dispersion <dipole-dipole< hydrogen bond
  8. 8. Very special Properties of1. 104.5o bond angle WATER2. 2 unshared pairs of valence electrons on oxygen atom3. Strongly polar covalent bond between O and H Hydrogen bonding between water molecules
  9. 9. • Surface tension – all particles of a liquid are attracted to one another (COHESION), but particles at the SURFACE of the liquid only have other particles of the liquid underneath them. This produces uneven attractions that pull surface particles closer together than particles within the body of the liquid.
  10. 10. • Capillary action – particles of the liquid are attracted to particles of their container (ADHESION). This causes the liquid to be pulled up into a narrow tube higher than if only gravity were acting. Allows water to rise in plant stems and blood to move into glass capillary tubes.
  11. 11. Why does ice float?When water molecules cool, their movement slows and more hydrogen bonds can form. Molecules are forced into a more rigid pattern that spreads them farther apart than in liquid phase.When molecules are farther apart, there are fewer molecules in a unit of volume and the density is less. Ice is LESS DENSE than water!
  12. 12. Ionic vs. Molecular CompoundsForces BETWEEN ions Forces WITHIN moleculesare very strong. are very strong. ForcesHIGH melting & boiling BETWEEN molecules arepoints; hard, brittle much weaker.solids. LOWER m.p. & b.p.,Solids DO NOT conduct softer solids, liquids orelectricity; fixed charges gases. AQUEOUS solutions Do not conduct conduct; mobilecharges.“aqueous” means“dissolved in water”
  13. 13. Ionic Compound VocabFORMULA UNIT: Simplest collection of atomsin an ionic compoundExpressed as an EMPIRICAL FORMULA:smallest whole number ratio of elementsTHERE ARE NO MOLECULES in an ioniccompound.CRYSTAL STRUCTURE: all ions surroundedby ions of opposite charge.
  14. 14. Examples of Crystal Structures: NaCl and CaF2