Ch.9 Flashcards
Lewis theory
A simple model of chemical bonding using diagrams that represent bonds between atoms as lines or pairs of dots. In this theory, atoms bond together to obtain stable octets (eight valence electrons)
Lewis electron dot structures
A drawing that represents chemical bonds between atoms as shared or transferred electrons; the valence electrons of atoms are represented as dots
-Only valence electrons are included. For main group elements, the number of valence electrons equals the group number or group number minus 10. The nucleus and Core electrons are represented by the symbol of the element, and the valence electrons are represented by dots. 
-A pair of bonding electrons (or a bond pair) between two Atoms is a single covalent bond and as shown by a line. A pair of non-bonding electrons on one Atom is a lone pair and is shown by two dots.
-Electrons are distributed so that Atoms acquire a stable electron configuration, usually an octet (8 electrons) for most Atoms but (two electrons for hydrogen atoms). Multiple bonds (double or triple bonds) may need to be formed.
Ionic bond
A chemical bond formed between two oppositely charged ions, generally a metallic cation and a non-metallic anion, that are attached to one another by electrostatic forces.
-metal + nonmetal
-1 to 1 attraction between cation and anion.
-Form an extended structure not molecules.
-Full (Completely) electron a transfer.
-Held together by attractive electrostatic forces
-Even though the formation of ions from the elemental metals and nonmetals requires a large input of energy, a great deal of energy is released when the Ions combine to form the ionic solid.
-The lattice energy is the energy released when the gaseous ions combine to form the ionic solid, and this is proportional to the electrostatic energy.
🔺H = -770 KJ/mol (exothermic) if released this amount when creating, need this amount to break them apart.
-when you go down a group size becomes larger, electrons are going to be further apart (more distance) = lower attraction from the nucleus.
-greater magnitude of changes = greater lattice energy = higher melting point.
Na(metal) —e-—> Cl(nonmetal)
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V
Na+ (Cation) e- transfer Cl- (anion)
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V
Ionic Solid (which has lattice E)
Physical properties of ionic solids
-hard and rigid: due to high attractive bonds between ions and he’s extend indefinitely
-brittle: ions of like charges are brought closer together, creating repulsive forces
-when melting ion compounds you’re breaking bonds (adding energy) causing a current
-Poor electrical conductors in the solid, but good conductors when molten or dissolved in solution
Covalent bond
Chemical bond in which 2 atoms share electrons that interact with the nuclei of both atoms, lowering the potential energy of each through electrostatic interactions
-nonmetal + nonmetal
-They share electrons, not exchange.
-Bonded atoms form a molecule.
-Covalent compounds typically exist as liquids or gases with low melting points or boiling points. This is harder to explain, because we need to consider the distance between the strong covalent bonds within the molecules in the weak forces between molecules. 
-intra = Between Atoms
-inter = Between molecules 
-Higher energy = unstable
-lower (more negative) energy = more stable 
Metallic bonding
The type of bonding that occurs in metal crystals, in which metal Atoms donate their electrons to an electron sea, delocalized over the entire Crystal lattice. 
-metal + metal
-Electrons are highly delocalized and shared by all Atoms
-Bonded atoms are held together by a “sea of electrons” in an extended structure
-They come together and give up all electrons and “throw them into a pool”. This means they become more of a metal cation. They are surrounded by the “mass” of electrons
-Metallic compounds are typically malleable (bond easily) and ductile (drawn into wires), have a moderately high melting points and boiling points, and conduct heat and electricity well.
Chemical bond
The sharing or transfer of electrons to attain stable electron configuration is for bonding Atoms
Octet rule
The tendency for most bonded atoms to process or share eight electrons in their outer shell to obtain stable electron configurations and lower the potential energy
Bonding pair
A pair of electrons shared between two Atoms
Lone pair
A pair of electrons associated with only one Atom
Non-bonding electrons
Lone-pair electrons
Double bond and triple bonds
The bond that forms when 2 pairs of electrons are shared between atoms.
-In general, is shorter and stronger than a single bond
The bond that forms when 3 electron pairs are shared between two Atoms.
-Are even shorter and stronger than double bonds. 
Summarizing the method for drawing Lewis structures
- Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule. 
- Write the correct skeletal structure for the molecule drawing a bond between each set of bonding Atoms.
- Distribute the remaining unaccounted for electrons, in pairs, among the Atoms, giving octets to as many atoms as possible
- If any non-hydrogen Atoms lack an octet, form double or triple bonds as necessary to give them octets. 
Or another way to explain
- Determine the total number of valence electrons.
-Add or subtract electrons if the molecule is an anion or a cation. - Determine how Atoms are connected.
-Identify the central Atom(s) and Terminal Atoms. Central Atom is usually the least electronegative Atom. 
-H and F Are always terminal Atoms, and form only one bond - Draw a skeletal structure by joining Atoms with single bonds
-Subtract two electrons for each single bond. - Distribute the remaining electrons in pairs
-First, complete octets Around the terminal Atoms (except H)
-Then, distribute the remaining electrons, if any, around the central Atom(s)
-A common mistake is to put too many electrons (in the mad rush to give every atom an octet), forgetting that you must work with exactly the total number of valence electrons that are available (step 1) - To few electrons?
-Convert lone pairs from Terminal Atoms to form multiple bonds with Central Atoms (To attain octets) 
Lattice energy
The energy associated with forming a crystalline lattice from gaseous ions
Bond energy
The energy required to break one mole of the bond in a gas phase.
-The bond energy(BE) of a bond A—B is the energy required to overcome the attraction of the two Atoms A and B.
Bond breaking:
- A—B(g) —> A(g) + B(g)
🔺H = +BE (endothermic)
Bond making:
A(g) + B(g) —> A—B(g)
🔺H = -BE (exothermic)
-shorter bond = stronger bond
-longer bond = weaker bond
-more bonds = more energy
-besides a single covalent bond, stronger multiple bonds can also form. The bond strength is related to its length and energy (depends on if there’s one bond, two bond, three bond, etc.)
-need to look at the distance = atomic radii to figure answers out
-in a chemical reaction, atoms are rearranged when bonds are broken in reactant molecules (absorbing energy) and made in product molecules (releasing energy).
-The net energy (or enthalpy) change results from the difference in bond energies.
-if + endothermic
-if - exothermic
🔺Hrxn = sum BE(bonds broken) - sum BE(bonds formed)
Homolytic Bond breakage
Bond breakage in which electrons in the covalent bond are divided evenly between the Atoms that make up the bond
