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- Understanding Chemical Bonds: Ionic and Covalent -

Discover the chemistry behind chemical bonding. Learn about ionic bonds resulting from electron transfer and covalent bonds from electron sharing. Understand principles like the Octet Rule and Lewis Dot diagrams. -

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- Understanding Chemical Bonds: Ionic and Covalent -

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  1. Chemical Bonding Chapter 7

  2. Chemical Bonds • Chemical Bond – a link between atoms resulting from the neutral attraction of their nuclei for electrons • force of attraction between atoms

  3. Ionic bond • Ionic Bond – chemical bond resulting from the transfer of electrons from one bonding atom to another • Ex. NaCl • Ionic Bonds form ionic compounds which are composed of ions.

  4. Characteristics of an ionic bond • -high melting points (shows a strong bond) • -brittle • -dissolve in water • -good conductors of electricity in solution

  5. Octet Rule • Octet Rule – atoms tend to gain, lose, or share electrons in order to acquire a full set of valence electrons. • Example of an ionic bond: • Consider NaCl – (table salt) • Na – 1s22s22p63s1 • Cl – 1s22s22p63s23p5

  6. What happens? • Na loses 1e- to form an octet • becomes Na+ ion • called CATION • Cl gains 1e- to form an octet • becomes Cl- ion • called ANION

  7. Using Lewis Dot diagrams • used to show valence electrons • dot represents valence electrons (or x’s and o’s) • random order • Na Cl -> Na+ + Cl -

  8. Formulas • Empirical Formula – chemical formulas which gives the simplest whole number ratio of atoms of elements in a compound. • -cation ALWAYS written first • -total “+” must equal total “-“

  9. Monoatomic ions – made of ONLY 1 element • Must remember the charge of the ion • Cations: Anions: • Group 1: +1 Group 17: -1 • Group 2: +2 Oxide: -2 • Silver: +1 Sulfide: -2 • Zinc: +2 Nitride: -3 • Aluminum: +3 Phosphide:-3

  10. Polyatomic ions • -ions which consist of more than one atom • they act as a whole and carry a net charge • Examples of Polyatomic Ions • Hydroxide - OH- Acetate – C2H3O2- • Hypochlorite – ClO- Sulfate – SO4-2 • Nitrate – NO3- Carbonate – CO3-2 • Bicarbonate – HCO3- Phosphate – PO4-3

  11. Binary Compounds • Binary Ionic Compounds • - contain ions of only 2 elements. • - Need to know the ratio of elements in compound

  12. Writing formulas: Criss-Cross method • write symbols of the elements or polyatomic ions putting the charge as a superscript • put parenthesis around only the polyatomic ion • crisscross the numbers not the charges and represent as subscripts

  13. Writing formulas…. • reduce if possible • eliminate and rewrite any subscript of 1 and any parenthesis whose subscript is 1. • Example: Na and O form • Na2O

  14. Covalent Bonds • Covalent Bond – chemical bond resulting from the sharing of electrons between bonding atoms • Ex. CO2

  15. 7.2 Covalent Bonds • Molecule- a group of atoms grouped together by covalent bonds. • Molecular substance – a substance made up of molecules.

  16. Molecular vs. Empirical • Molecular formula – tells you how many atoms are in a single molecule of the compound. • Ex. C6H12O6 glucose • The empirical formula can be written for molecular formulas. • Ex. CH2O glucose

  17. Formulas • Structural formula – shows which atoms bond in a molecule. • - Lewis Structures • - Based on the Lewis Dot diagrams • - The electrons between the 2 elements are the shared electrons.

  18. Ex: F2 and NH3 • Unshared pair – Electrons needed to satisfy the octet • Note: H is an exception. It can only hold 2 e-s

  19. Single Covalent bond – 2 atoms share exactly one pair of electrons • Double Covalent bond – 2 pairs of shared bonds (4 electrons between 2 atoms) • Ex. H2CO • Triple Bond – there will be 6 dots (electrons) between 2 atoms • Note: Use a line to indicate a bond (or 2 electrons). Use dots to show electrons that are unshared.

  20. Exceptions • Exception to the octet rule • Atoms with less than an octet - Boron • Ex. BF3 • Atoms with more than an octet - Sulfur • Ex. SF4

  21. Properties of covalent bonds • Properties of Covalent Bonds • Bonds can be polar depending on the Electronegativity of the atoms. • If one atom is significantly more electronegative than the other, the electron density will change.

  22. The atom that is more electronegative will pull the electrons closer to it and cause a partial negative on that atom. • **Electronegative difference must be between 0.5 and 1.9 • If the electronegatives are similar, the bond will be nonpolar (electronegative difference less than 0.4)

  23. Ionic bonds Involves transfer of electrons • Forms ions • Usually forms between metal and nonmetal and/or polyatomic atom • substances usually soluble in water • high melting point

  24. solutions are usually good conductors of electricity • compounds are brittle • formulas are expressed as empirical formulas • 9.Electronegativity difference greater than 1.9 • 10.strong bond • 11. criss cross

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