Unit 6, 7 and 8

Unit 6 Ionic Bonding Highlights

• Symbol- a one, two, of three letter designation for an element (Ex. Sodium is Na)

• Subscript- a number written to the right and below a symbol that tells the number of atoms present in a molecule (Ex. An oxygen molecule has two oxygen atoms, O₂)

• Formula- symbols and subscripts used to represent the composition of a substance (can be ionic or covalent (Ex. H₂O, KCl)

• Chemical Bonds- forces that hold atoms together

  • protons of one atom and electrons of another are attracted

• Ionic bonds- between metals and nonmetals, they transfer electrons, metals are losers (lose electrons) and nonmetals gain electrons 

  • Form an ionic compound (a salt), ions stick together like magnets

  • Mostly crystalline solids due to high boiling and melting points 

  • Conduct electricity when in the liquid or dissolved state (needs mobile ions)

• Compound- a substance composed of 2 or more elements chemically combined in definite proportions

  • Ex. MgCl₂ is a compound but Br₂ is NOT

  • Compounds can be broken down by chemical means (unlike pure elements that cannot)

• Polyatomic Ion- a group of covalently bonded atoms possessing a charge, they take part in ionic bonding (Ex. NH₄⁺, CO₃^-2)

  • A compound can have both covalent and ionic bonding if it has a polyatomic ion Ex. BaCO₃, Ca(OH)₂, NH₄F

• Writing Ionic Formulas: Criss-Cross Method

  • Ex.  Mg⁺²Cl^-1 🡪 MgCl₂

  • Ex.  Ca⁺²NO₃^-1 🡪 Ca(NO₃)₂   

    • (Polyatomic ions need parentheses if there is more than one)

  • Ex.  Be⁺²O^-2 🡪 Be₂O₂ 🡪 BeO

    • (Always reduce to most reduced form)

• Naming Ionic Compounds:

  • Metal named first, change ending of nonmetal to “-ide”

    • Ex. NaF 🡪 Sodium Fluoride

  • Transition Metals with more than one possible charge need a Roman Numeral to tell the charge of the metal atom

    • Ex.  MnI₂ 🡪 Manganese (II) iodide 

  • You can find the charge of the metal by using Reverse Criss-Cross if subscripts are present OR by checking the charge of the anion

    • Ex.  Fe₂O₃ 🡪 Fe⁺³O^-2 🡪 Iron (III) oxide

    • Ex.  CuSO₄ 🡪 Cu⁺²(SO₄)^-2 🡪 Copper (II) sulfate

    • Ex.  AuF 🡪 Au⁺¹F^-1 🡪 Gold (I) fluoride

• Lewis Dot Diagram- place number of valence electrons around the element’s symbol, always placing two on top first, then one on each side as you draw them

  • Ex. 

• Lewis Diagrams- every atom needs its octet

  • Ionic- metals give all electrons to nonmetal, forms ions, write brackets and charges, nonmetals should have 8 electrons, metals should have none

    • Ex.:

Unit 7 Highlights: Covalent Bonding

• Covalent bonds- between nonmetals, share valence electrons to give all atoms octets

  • Form a molecule

  • Poor conductors of heat and electricity 

  • Mostly gases, liquids and softer solids because of lower boiling and melting points 

  • Single Bond 🡪 2 e^-, Double Bond 🡪 4 e^-, Triple Bond 🡪 6 e^-

• Lewis Dot Diagram- place number of valence electrons around the element’s symbol, always placing two on top first, then one on each side as you draw them

*Except Carbon, one dot each side

• Lewis Diagrams- every atom needs its octet

  • Covalent- connect lone dots so as to complete octets, redraw shared electrons as lines to represent bonds

    • Ex.:
      

• Polarity- a measure of electron distribution within a molecule (covalent only because they share e^-), is determined by the difference in electronegativity between atoms

  • Polar- uneven distribution of electrons, different electronegativities (different elements), do not share electrons equally

    • One atom is more negative than the other because it pulls the electrons towards it (Ex. HCl, H₂O, PCl₃, NH₃)

    • WATER IS POLAR!!!!!!

  • Nonpolar- even distribution of electrons, same electronegativity (same element), share electrons equally (Ex. H₂, O₂)

    • Exception: A molecule can have polar bonds but still be a nonpolar molecule if it is symmetrical 

    • SNAP- Symmetrical Nonpolar, Asymmetrical Polar

    • 2 Common Cases of Symmetrical Molecules

      • CO₂ & CH₄ (and others like it CF₄, CI₄) 

• VSEPR (Valence Shell Electron Pair Repulsion)

  • Molecular shape depends on the pairs of valence electrons around the central atom which repel each other

    • Electron pairs will arrange themselves to be as far apart from each other as possible

    • Repulsive unit: a single bond, double bond, triple bond or lone pair - will try to get as far away from other repulsive units as possible

• Naming Covalent Compounds

  • First element stays the same, second ends in “-ide”

  • Use prefixes for both to tell the number present (never use “mono-” for the first)

  • Ex.  CBr₄ 🡪 Carbon tetrabromide

  • Ex.  N₂O₃ 🡪 Dinitrogen trioxide

  • Ex.  CO 🡪 Carbon monoxide

Unit 8 Highlights

• Physical Change- a change in state (Gas, liquid, solid) or physical appearance (the substance stays the same from beginning to end)

  • Ex.   CH₄(g) 🡪 CH₄(l)

• Chemical Change- then the identity of the products is different from the identity of the reactants (a chemical reaction, the substance changes)

  • Ex.    2H₂(g) + C(s) 🡪 CH₄(g)

  • Signs a Reaction has occurred

    • Color change

    • Bubbles form (gas released)

    • Temperature change

    • Odor (smell)

    • Formation of a precipitate (insoluble solid)
      

• Reactant- a substance that enters a reaction, on the left side of the equation

• Product- a substance formed in a reaction, on the right side of the equation

  • Reactants 🡪 Products

• Coefficient- a number written in front of a formula to tell how many substances or molecules there are (Ex. 2NH₃ means two NH₃ molecules)

• Five main types of chemical Reactions:

  • Synthesis (Combination) Reaction- A + B → AB 

    • (Ex. 2H₂ + O₂ → 2H₂O)

  • Decomposition Reaction- AB → A + B

    • (Ex. CaCO₃ → CaO + CO₂)

  • Single Replacement Reaction- A + BC → B + AC 

    • (Ex. Zn + CuI₂ → Cu + ZnI₂)

  • Double Replacement Reaction- AB + CD → AD + CB 

    • (Ex. KBr + MgO → K₂O + MgBr₂)

  • Combustion Reaction - C_xH_y + O₂ → H₂O + CO₂

    • (Ex. CH₄ + 2O₂ → 2H₂O + CO₂)

• Law of Conservation of Mass: states that matter AND energy are both conserved in a chemical reaction (charge is also conserved).

  • This implies that reactions must be balanced 

  • To balance a reaction one must modify coefficients to make sure that the number of atoms of each element on both sides of the reaction are equal 

  • Ex. ___Na + ___Cl₂ 🡪 ___NaCl

     2Na  +  Cl₂   🡪  2NaCl

(2 Na atoms and 2 Cl atoms on both sides)

• Ex. ___AlBr₃ + ___K₂SO₄ → ___KBr + ___Al₂(SO₄)₃

2AlBr₃ + 3K₂SO₄ → 6KBr + Al₂(SO₄)₃ 

(2 Al atoms, 6 Br atoms, 6 K atoms, 3 SO₄ ions)