Some Things For Final Flashcards
Bond order provides measurement of stability of a bond. How?
The higher, the bond order, the more stable the bond
Electronegative differences
Overall dipole moment = polar
EN diff less than 0.4 = nonpolar. Covalent bond (e^- shared equally)
EN diff below 2.0 but above 0.4 = polar. Covalent bond (e^- usually shared unequally, has dipole + & -)
EN diff greater than 2.0 = ionic bond (e transferred)
If there is a net dipole moment = polar (not symmetrical)
If net diaper moment is 0 = nonpolar
As frequency of EN radiation __, wave length__
Increases, decreases
As energy of EN radiation __, frequency __
Increases, increases
As frequency of EN radiation __, energy __
Decreases, decreases
As energy of EN Radiation __, wave length __
Increases, decreases
As wave length of EN radiation __, energy __
Increases, decreases
What causes atomic radius to increase while moving down the periodic table
The number of electron shells increases. Each successive energy level is further away from the nucleus = radius increases.
Atomic radius increases due to addition of electrons in the next energy level
Atomic radius decreases, due to increase number of protons. Positively charged protons in the nucleus pull charged electrons closer to the nucleus, causing the radius to decrease.
Metallic character, increase or decrease or stay same down the periodic table?
Increase moving down. Due to the distance, electrons are away from the nucleus. Further the distance the easier it is to remove an electron.
Easier to remove an electron = more metallic character an element has.
Ionic Bonding:
metals + nonmetals
• electrons are transferred completely from one atom to another
• cations and anions, arranged to form, not molecules, but an extended structure
• The cations and anions are held by attractive electrostatic forces (given by Coulomb’s law) E (no symbol for) q-q2/d, where q1 and q2 are the ion charges and d is the separation between them
— smaller attraction = larger distance = less energy?
— higher the charges are, the attraction is larger and larger the distance
Lattice Energy
• 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 these 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.
— lattice energy —> infinite number of anions/cations bonded in a very specific pattern 
Physical Properties of Ionic Solids + trend
• hard and rigid: due to high attractive bonds between ions and these extend indefinitely
• brittle: ions and like charges are brought closer together, creating repulsive forces.
• poor electrical conductors in the solid, but good conductors when molten or dissolved in solution:
— when melting an ion compound, you’re breaking bonds (adding energy) causing a current.
• 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 Charges = greater lattice energy = greater melting point.
• less energy required —> lower melting point
Covalent Bonding:
nonmetals + nonmetals
• electrons are shared between atoms (localized in the space between them)
• higher energy —> unstable
• lower(more negative) energy —> more stable
• bonded atoms form a molecule
• Covalent compounds typically exist as liquids or gases with low mp or bp. This is harder to explain, because we need to consider the distinction between the strong covalent bonds within molecules and the weak forces between molecules.
Metallic Bonding:
metals + metals
• electrons are highly delocalized and shared by all atoms
• bonded atoms are held together by a “sea of electrons” in an extended structure
• Metallic compounds are typically malleable and ductile, have moderately high mp and bp, and conduct heat and electricity well.
Bond Energy
• 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)
• Besides a single covalent bond, stronger multiple bonds can also form. The bond strength is related to its length and energy.
• Shorter bond = stronger bond
• longer bond = weaker bond
• more bonds = more energy
• somethings that has more energy = needs more energy to break the bonds = means it is not a good fuel
•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
Electronegativity
• Most bonds are intermediate, having both ionic and covalent character. What determines the extent of sharing of electrons?
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V
• Electronegativity (EN or x): The relative ability of an atom, bonded within a molecule, to attract shared electrons to itself.
— electron affinity ( the energy change that happens when an atom gains electrons. For a specific Atom)
• In his approach, Pauling noted that the bond energy in H–F was greater than the average of H–H and F– F, and attributed the extra stability to an electrostatic (ionic) contribution arising from the difference in electronegativity.
• In a polar covalent bond, the more EN atom takes a greater share of the bonding e– and acquires a partial negative charge (d–), whereas the less EN atom acquires a partial positive charge (d+).
Trends
• as you go down groups EN decreases
• as you go across period EN increases
• F > O > Cl, N > Br > I,C,S > H
• more polar = larger EN difference
• as you go across a period Z effective increases, atomic radius decreases, nucleus gets closer together —> smaller atomic radius.
•EN inversely related to atomic size
• smaller atoms are more electronegative
