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Types of bonds

Types of bonds. Ionic-transfer of electrons Covalent-sharing of electrons Metallic-metals lose electrons which ‘float in a sea’ around the cations. Formation of bonds. Atoms will position themselves to achieve lowest possible energy Distance where energy is minimum is bond length

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Types of bonds

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  1. Types of bonds • Ionic-transfer of electrons • Covalent-sharing of electrons • Metallic-metals lose electrons which ‘float in a sea’ around the cations

  2. Formation of bonds • Atoms will position themselves to achieve lowest possible energy • Distance where energy is minimum is bond length • Bond will form if energy of aggregate is lower than that of the separated atom • Example: H2 is more stable than two separate H atoms. e- reside in space between two H atoms where they are attracted equally by both p+ (covalent bond)

  3. Ionic Bonds metal + nonmetal lose e- gain e- cation anion + 

  4. Ionic compound properties Most are crystalline solids b/c ions are close together Held together by electrostatic force (attraction between + and – ions) Conduct electric current in molten state High B.P. and M.P. Brittle Soluble in water

  5. Lattice energyA lattice is a stable, ordered solid three-dimensional array of ions. • Indicates the strength of an ionic bond • Energy released when ionic solid forms its ion. • Energy required to separate a mole of a solid ionic compound into its gaseous ions • Higher lattice energy usually means lower solubility • Higher lattice energy; higher melting point • Larger ion; smaller lattice energy • Higher charge; larger lattice energy • Coulombs law lattice energy = k (Q1Q2) r

  6. Covalent Bonds Atoms share electrons Nonmetal + nonmetal

  7. Covalent compound properties • Low boiling point • Variable solubility • Do not conduct electricity • Variety of states • Held together by overlap of unfilled orbitals

  8. Number of covalent bonds that an atom can form = number of valence electrons Electronegativity used to predict bond type < 0.3 nonpolar covalent bond >0.3 but <1.7 polar covalent bond >1.7 ionic bond (metal + nonmetal) Covalent Ionic 0.0 polarity increases as electronegativity difference increases Numbers are guidelines only

  9. Bond Polarity/ Dipole moment Examples: d- d+ CH4 C = 2.5 H = 2.1 C—H C—H d- d+ NH3 N = 3.0 H = 2.1 N—H N—H d- d+ H2O H = 2.1 O = 3.5 H—O H—O

  10. + + − + + − − − − − − − − − − − − − + + + + − − − − − − − − − − − − − − − + + + + − − − − Metal bondsMetals lose electrons which form a ‘sea’ that floats around the cations. Electron is not associated with any particular cation. Properties of metals directly related to this ‘sea’ • Malleable • ductile • Conduct heat and electricity • Luster (shiny)

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