Unit 2: Molecular and ionic compound structure and properties Flashcards
Chemical Bonds
Attraction between the nucleus of one atom and the electron of another
*There’s a balance between attraction (+, -) and repulsion (+, +)(-, -)
Bond Energy
Energy required to break the bond
Bond Length
Distance between atoms
Aspects of a single bond
- Fewest shared electrons
- Weaker attraction between nucleus of one atom and the bonding elements
- Weaker + Longer bond
Aspects of a Triple Bond
- Most shared electrons
- Stronger attraction between nucleus of one atom and the bonding electrons
- Stronger + shorter bond
Ionic Bonds
Electrostatic attraction between a cation (positively charged) and anion (negatively charged)
Mg(OH)2 -> Mg^+2 + 2OH-, Lattice Energy: 2900
Sr(OH)2 -> Sr^+2 + 2OH-, Lattice Energy: 2300
Explain why in terms of periodic properties and Coulomb’s law that it takes less energy for Sr(OH)2 to separate its ions in terms of periodic property and Coulomb’s law.
Periodic property: The Sr^2+ ion is larger than the MG^2+ because it has more occupied energy shells; the bond length is longer and weaker
Coulomb’s Law: Since the distance between Sr^2+ and OH- is longer than the distance between Mg^2+ and OH-, the attractive forces in Sr(OH)2 are weaker, and its lattice energy is smaller
Electronegativity
The ability of an atom in a molecule to attract shared electrons to itself
Ionic bonds (electron definition)
Transferring electrons (usually involves a metal and nonmetal)
*think of it as: if the two elements differ in electronegativity, it is more ionic, and the more electronegative atom has more ability to attract or transfer shared electrons
Covalent bonds
Sharing electrons (usually two nonmetals)
*think of it as: if two elements have similar electronegativity, covalent electrons share electrons more equally
Nonpolar (covalent bond)
Electrons shared equally
Polar (covalent bond)
Electrons shared unequally
Metallic bonds
Electrons not associated with a atom or molecule (delocalized electrons)
Positive metal ions surrounded by a sea of mobile valence electrons
Bond Polarity
Difference in electronegativity values of two elements
δ− (negative partial charge)
Atoms with high electronegativity
δ+ (positive partial charge)
Atoms with low electronegativity
Dipole arrow points…
Towards the more electronegative atom
A student claims that F2 is more polar than H2 because fluorine has a higher electronegativity than hydrogen. Do you agree with this claim?
No; Polarity is the difference in electronegativity, not the absolute value of the electronegativity. Fluorine and hydrogen have the same difference in electronegativity (0), so they have the same polarity.
In order to conduct electricity, a substance must have…
- Charged particles
- particles that are free to move
Alloys
Combining two or more metallic elements
Substitutional allow
- Atoms of similar radii
- One atom substitutes for another atom in the lattice
Interstitial alloy
- Atoms of different radii
- Smaller atom fills the spaces between larger atoms
- Usually the alloy is stronger than the base metal
Octet rule
All atoms end up with 8 electrons around them
Expanded Octet
Atoms in periods 3-7 can bond with other atoms in such a way they end up with more than 8 electrons in their octets
They can do this because they have d-orbitals in their outer shells that can accept electrons
Formal charge
The hypothetical charge the atom would have if all of the atoms had the same electronegativity
Calculated to identify the most stable/likely structure
Formal charge formula
(# of valence electrons assigned to the neutral atom) - (# of electrons assigned to the atom in the structure)
**bonds count as one!!!
The potential energy of valence electrons ___ as they approach the nucleus of another atom
Decreases
Bond energy is ___ during the formation of a bond
Released
Potential energy and bond energy has an __ relationship
inverse
As the # of bonds between 2 atoms increases, the bond length ___ and the bond energy ___
Decreases; Increases
Bond Order
The number of bonds between two atoms
Resonance Structures
A set of two or more Lewis Structures; occurs when a molecule has at least one double bond
VSEPR Theory (Valence Shell Electron Pair Repulsion)
Charge clouds repel each other due to Coulombic repulsion- terminal atoms move as far away from each other as much as possible
2 Charge clouds, 2 bonds, 0 lone pairs
Linear, 180 angle
3 charge clouds, 3 bonds, 0 lone pairs
Trigonal Planar, 120 angle
3 charge clouds, 2 bonds, 1 lone pairs
Bent, 120 angle
4 charge clouds, 4 bonds, 0 lone pairs
Tetrahedral, 109.5 angle
4 charge clouds, 3 bonds, 1 lone pairs
Trigonal Pyramidal, 107.3 angle
4 charge clouds, 2 bonds, 2 lone pairs
Bent, 104.5 angle
5 charge clouds, 5 bonds, 0 lone pairs
Trigonal Bipyramidal, 90 and 120 angle
5 charge clouds, 4 bonds, 1 lone pairs
Seesaw, 90 and 120 angle
5 charge clouds, 3 bonds, 2 lone pairs
T-Shaped, 90 angle
5 charge clouds, 2 bonds, 3 lone pairs
Linear, 180 angle
6 charge clouds, 6 bonds, 0 lone pairs
Octahedral, 90 angle
6 charge clouds, 5 bonds, 1 lone pairs
Square pyramidal, 90 angle
Rules of best Lewis diagram structure:
- Will have the minimum number of nonzero formal charges
- If nonzero formal charges must remain, the negative charge must be assigned to the most electronegative atom
- The sum of all the individual formal charges must add up to the charge of the chemical species
6 charge clouds, 4 bonds, 2 lone pairs
Square Planar. 90 angle
Molecular Geometry rules:
- Molecules composed of only nonpolar bonds are nonpolar, regardless of shape
- Molecules in which the central atom is symmetrically surrounded by identical atoms are nonpolar, even if the bonds are polar
- Molecules with asymmetrical shapes that contain any polar bonds are polar overall
Symmetrical shapes:
Linear (3 atoms), trigonal planar, tetrahedral, trigonal bipyramidal, square planar, octahedral
