Unit 8 Bonding and Nomenclature Bonding Notes

Covalent and Ionic Bonds

TYPES OF CHEMICAL BONDS Chemical bond: a force that holds groups of 2 or more atoms together to function as a unit (molecule). Ionic bond: An atom (metal) transfers electron(s) to another atom (nonmetal), which creates an electrostatic attraction of a positively charged ion and a negatively charged ion. Cation: a positively charged ion which is created when an atom loses one or more e-. Anion: a negatively charged ion which is created when an atom gains one or more e-.

TYPES OF CHEMICAL BONDS Covalent bond: The type of bonding where electrons are shared by two nuclei.   Nonpolar covalent bond: A covalent bond with equal sharing of electrons between the two nuclei. Polar covalent bond: A covalent bond with unequal sharing of electrons between the two nuclei

Dogs Teaching Chemistry

Polar/ Nonpolar Polar Covalent Nonpolar Covalent

ELECTRONEGATIVITY Electronegativity: the ability of an atom in a molecule to attract shared electrons to itself

Electronegativity cont. Generally increases going from left to right, decreases going down the periodic table. Metals relatively low electronegativity Nonmetals  relatively high electronegativity The higher the atom’s electronegativity value, the closer the shared electrons tend to be to that atom when it forms a bond.

ELECTRONEGATIVITY CHART

THE RELATIONSHIP BETWEEN ELECTRONEGATIVITY AND BOND TYPE

Using Electronegativity to Determine the Type of Bond Nonpolar: a covalent bond with equal sharing of electrons between the two nuclei. <0.3 Polar: a covalent bond with unequal sharing of electrons between the two nuclei. 0.3-1.4 Above 1.5 is ionic

Electronegativity Problems H2 H-H H= 2.1 Find the difference: 2.1-2.1= 0 Nonpolar covalent HCl H-Cl H= 2.1, Cl=3.0 Find the difference: 3.0-2.1= 0.9 Polar covalent

Electronegativity Problems con’t NaCl Na – Cl Na = 0.9 Cl = 3.0 Find the difference: 3.0 -0.9 = 2.1 Ionic bond What type of bond would be found in between the following atoms? a. Mg and O b. S and O c. O and O 3.5 -1.2 = 2.3 3.5 - 2.5 = 1.0 3.5 - 3.5 = 0.0 Ionic Polar Non-polar

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What Holds Bonded Atoms Together Atoms bond when their valence electrons interact Remember that atoms with full s and p orbitals are more stable than atoms with only partially full s and p orbitals Atoms tend to bond together so that they have stable s and p orbitals (stable electron configuration) like the noble gases

Valence Electrons Valence electrons are the electrons that can participate in the formation of a chemical bond. These electrons can be found in the highest energy level of an atom. In electron configuration, biggest number and usually s and/or p orbitals.

Examples of valence electrons Lithium- Li Calcium- Ca Silicon- Si 1s22s1 or [He]2s1 Number of valence electrons? [Ar]4s2 [Ne] 3s23p2 1 2 2

Lewis Structures

Lewis Dot Structures Lewis Structure: a representation of a molecule that shows how the valence e- are arranged among the atoms in the molecule. Provide a way to draw molecules and better visualize it. How to draw Lewis Dot structures First write the element symbol Determine the number of valence electrons Add valence electrons around element symbol (one dot represents one valence electron) Add one electron to each side before adding two!!

Lewis Dot structure Practice Draw Lewis Structures for each of the following atoms Li C P Br Kr

Lewis Dot structures for Molecular Compounds The most important requirement for the formation of a stable compound is that the atoms achieve a noble gas e- configuration. Octet rule: 8 electrons are required to fill in the s and the p orbitals in the valence energy level. Exception: Duet rule (2 electrons) for hydrogen. Bonding pairs: e- that are shared between 2 atoms Single bond: one pair of e- are shared between 2 atoms Double bond: two pairs of e- are shared between 2 atoms Triple bond: three pairs of e- are shared between 2 atoms   Lone pairs: e- that are NOT involved in bonding.

Lewis Structures When drawing Lewis structures: Remember to think about bonding sites (terminal atoms on the “outside”) N – A = S electrons needed–electrons available= electrons shared octet or duet valence electrons bonding electrons

Practice drawing Lewis Structures Draw Lewis electron dot diagrams for each molecule listed below. 1. Br2 2. HI 3. NH3 4. PCl3   5. H2O 6. CBr4 7. SeO2 8. N2  

End of Day 2

Ionic Bonds Ionic Bonds form between ions with opposite charges (between a metal with a positive charge and a non-metal with a negative charge) One atom will lose electron(s) and the other will gain electron(s) so that they both have stable electron configurations (8 valence electrons)

Ionic Bonds (con’t) EX. NaCl, sodium chloride Sodium has 1 valence electron, it loses the electron to become a positive ion, Na+ Chlorine has 7 valence electrons, it gains one electron to become a negative ion, Cl- The oppositely charged ions attract each other and form an ionic bond Ionic compounds do not form molecules Ions always exist in the same ratio to one another

Ionic Solids Ionic solids have high melting points, will not conduct electricity as a solid (ions are kept in place), but they will when dissolved in water or when melted (ions are free to move)

Lewis Dot structures for Ionic compounds Write element symbols Transfer the electrons! Add the charge!

Ionic compounds Lewis practice Draw the Lewis Structures for the following ionic compounds 1. MgCl2 2. MgO 3. K3P

Structural Formula Structural Formula: The overall shape of the molecule is represented by elements symbols and a single covalent bond is represented by a straight line or dash. Isomers: substances having the same chemical formula but different structures. Go back to the Lewis Structures for Molecular Compounds and write out the structural formulars