Terms and their definitions to keep in mind:
chemical bond; ionic bond; covalent bond; Lewis (dot) structure/symbol; electronegativity; dipole moment; lone pair (of electrons); VSEPR
There are ~100 elements, but millions of compounds with different chemical and physical properties; and chemistry is the tool for the synthesis and study of these compounds.
Example: How many different ways for 5 C atoms and 12 H atoms to bind together? (Hint: Each C can have 4 "bonds", but H can have only one.)
Atoms in compounds are held by "chemical bonds". Depending upon the properties of the atoms, different chemical bonds are formed.
Molecule: an electronically neutral group of tightly bound atoms that acts as one single unit. Its composition can be shown by a chemical formula, e.g., oxygen gas as O2, sugar as C6H12O6, the amino acid glycine as C2H5NO2. <<What about NaCl?>> The chemical formula of a molecules gives no clue about the structure of the molecule.
Review: Valence electrons (VEs): the electrons in the outermost principal energy level of the atoms of an element (which determine the chemistry of the element). Elements in the same group have the same number of VEs, e.g., Mg: [Ne]3s2, Ca: [Ar]4s2, and Sr: [Kr]5s2
Electron dot structures (Lewis dot structure): A "pictorial"
representation of valence electrons
—Dots (which are the valence electrons) are placed around 4 sides of the symbol of an element, following the Hund's rule).
such as :F • with 7 dots/VEs (1s2 2s2 2p5),
•Be• with 2 dots/VEs (1s2 2s2),
and :O with 6 dots/VEs (1s2 2s2 2p4).
Ions: Formed from atoms by gaining or losing electrons
Cation— a "+" charged ion; anion— a "–" charged ion
The easier an atom can form a cation, the easier the electron(s) can be removed from the atom.
Ionization energy of an atom: The energy required to remove an electron from the atom in the gaseous state.
Inert gases have the highest IEs, thus do not form ions easily.
In other words, inert gases are very stable, thus take relatively higher energy to reach excited state and ionized.
Conversely, alkali metals are very active (unstable), thus take relatively less energy to become ionized.
Cations (cf. Tables 3.1 & 3.3)
For example: Na, 1s2 2s2 2p6 3s1 Ne, 1s2 2s2 2p6
(Very active) (Very inert)
Na (Na•) ---> loses VE (from 3s) ---> Na+ + e–
1s2 2s2 2p6 (a Ne-like electron configuration with 8 VEs)
How about Mg, 1s2 2s2 2p6 3s2?
**Transition metals can lose 2 e– from an s orbital, and
more e– from the d orbitals, e.g., Fe2+, Fe3+,
and Fe4+; Cr2+ to Cr6+; etc.
Fe, 1s2 2s2 2p6 3s2 3p6 4s2 3d6;Cr, ...4s2 3d4
For example: Cl, 1s2 2s2 2p6 3s2 3p5 Ar, 1s2 2s2 2p6 3s2 3p6
:Cl • ---> gains 1 e– ---> Cl– ,
1s2 2s2 2p6 3s2 3p6 (an Ar-like electron configuration with 8 VEs)
How about O, 1s2 2s2 2p4?
Ions can be formed from more than one atoms (Table 3.2): NH4+, NO3–, OH–, CN–, CO32–, PO43–, SO42–, etc.
Ionic bonds—formed by the attraction between oppositely charged ions, e.g., Na+ and Cl–, Mg2+ and PO43–, and so on.
—may not have equal number of cations and anions
—Anios and cations in some ionic compounds can be "separated" by water molecules when dissolved in water.
—Can conduct electricity when dissolved in water or melted.
Examples (P. 68–72):
LiF (Li gives 1 e– to 1 F to form Li+ and F–.)
MgF2 (Mg gives 2 e–‘s to 2 F’s, one each, to form Mg2+ and F–.)
AlF3 (Al gives 3 e– to 3 F’s to form Al3+.)
Al2O3 (2 Al’s gives total of 6 e–‘s to 3 O’s to form 2Al3+ and 3O2–.)
Covalent bonds—formed by electron sharing between atoms,
e.g., hydrogen gas, oxygen gas, nitrogen gas, etc.
H• + •H ---> H2 (H:H or H–H)
.. .. .. ..
:F • + • F: ---> :F : F : (F2 or F–F)
• • • • • • • •
**Sharing of one electron pair (single bond) and a few unshared pairs not shown (lone pairs)
:O • + •O: ---> :O :: O : (O::O or O=O)
Sharing of two electron pairs — a double bond
How about N + N ---> N2? <<a "triple bond"!!>>
Heteronuclear molecules can be formed in a similar way
such as, H2O, NH3, CO2, CH4 (methane), and so on.
Coordinate covalent bonds (PP. 83.84)—the shared electron paired
between two atoms/ions is contributed by only one atom,
. .. ..
• C • + •O: ---> :C :: O : (This is not correct! Why not?)
• C • + •O: ---> :C ::: O : (This is correct! Why is it correct and how does it formed?)
H+ + :NH3 ---> H:NH3+ (NH4+)
Some "rules" about ions with respect to the periodic table)
Naming compounds (cf. Tables 3.1, 3.2, and 3.4 and example 3.6)
Representation of covalent bonds: use a "bar", —, to represent a covalent bond, such as H2 as H–H and O2 as O=O, etc.
Sections 3.7 and 3.8 show some step-by-step descriptions about how to draw Lewis structures of both simple and more complicated molecules.
Dipole interaction—interactions between polar molecules,
such as: Hd+—Fd– •••••• Hd+—Fd–
H-bond: a kind of dipole interaction, in which a hydrogen atom
e.g., >Nd–—Hd+ •••• Od–=Cd+< in proteins; the bonds that hold the two helices of DNA together
VSEPR (Valence-Shell Electron-Pair Repulsion Theory) and the shape of molecules: The non-bonding electron pairs in a molecule is taken into consideration for the determination of molecular shapes
What are the possible and correct shapes for the different HnX
molecules? How do H2O and NH3 interact with H+?
HCl H2O NH3 CH4
linear bent trigonal pyramidal tetrahedral (But why?)