GHS Enriched Chemistry

Chapter 6 – Chemical Bonding

Section 6-3Ionic Bonding and Ionic Compounds

Please use your textbook to thoroughly answer the following questions.

1. Describe how sodium and chlorine bind to form sodium chloride.

  • Sodium gives up one electron and becomes a cation.
  • Chlorine accepts an electron and becomes an anion.
  • The oppositely charged atoms are attracted to each other, and become NaCl.

2. What is an ionic compound?

  • Compound composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal.
  • In contrast to molecular compounds, ionic compounds are not composed of independent, neutral units that be isolated and examined.
  • Instead, they are found as crystals.

3. What is a crystal?

  • A three-dimensional network of positive and negative ions mutually attracted to one another.

4. What does the chemical formula of an ionic compound represent?

  • The simplest ratio of the compound’s combined ions that gives electrical neutrality.

Main Idea – Ionic bonds form from attractions between positive and negative ions.

5. Using electron-dot notations, please show how sodium chloride and calcium fluoride are produced.

6. Why do ionic compounds form a crystal lattice?

  • This arrangement minimizes their potential energy and is more stable.

Main Idea – Differences in attraction strength give ionic and molecular compounds different properties.

7. Describe the following properties of ionic compounds. When appropriate, please describe how property is a result of the structure of ionic compounds.

  • The melting point, boiling point, and hardness of a compound depend on how strongly its basic units are attracted to each other.
  • melting point –
  • Have higher melting points than covalent compounds due to strong attractive forces between ions.
  • boiling point –
  • Have higher boiling points than covalent compounds due to strong attractive forces between ions.
  • hardness –
  • Ionic compounds are hard but brittle.
  • Even a slight shift of ions relative to another causes a large buildup of repulsive forces.
  • These forces make it difficult for one layer to move relative to another.
  • If they do shift, the repulsive forces make the layers part completely, which is what we perceive as being brittle.
  • conductivity –
  • In the solid state, they do not conduct electricity.
  • In molten state, and when dissolved in water, they do conduct electricity, because the ions are free to move and carry electrical current.

Main Idea – Multiple atoms can bond covalently to form a single ion.

8. Define polyatomic ion – A charged group of covalently bonded atoms.

Section 6-4 Metallic Bonding

1. What is the conductivity of metals the result of? (Find this in the intro section.)

  • The conductivity of metals is due to the highly mobile valence electrons of that atoms that make up the metal.
  • These mobile electrons are not possible in covalent or ionic bonds.

Main Idea – Metal electrons move freely in empty, overlapping orbitals.

2. Within a metal, the vacant orbitals in the outer energy levels overlap. What does this allow the outer electrons to do?

  • This overlapping of orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal.

3. What does it mean if we say electrons are delocalized?

  • Delocalized electrons do not belong to any one atom but move freely about the metal’s network of empty atomic orbitals.
  • These mobile electrons form a sea of electrons around the metal atoms.

4. Define metallic bonding – The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.

5. Explain what causes a metal’s luster.

  • Metals are both strong absorbers and reflectors of light.
  • Because they contain many orbitals separated by extremely small energy differences, metals can absorb a wide range of light frequencies.
  • This causes electrons to get excited, and then immediately fall back down to lower levels, emitting energy in the form of light.
  • This re-radiated light is responsible for the metallic appearance or luster of metal surfaces.

6. Define the following:

  • malleability – the ability of a substance to be hammered or beaten into thin sheets.
  • ductility – the ability of a substance to be drawn, or pulled, or extruded through a small opening to produce a wire.

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