Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

CHEM104 > Chapter 9 > Flashcards

Chapter 9 Flashcards

1
Q

Lattice Energy

A

energy associated with forming a crystalline lattice of alternating cations and anions from the gaseous ions.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Hess’s Law

A

If a chemical equation can be expressed as the sum of a series of steps, then the change in delta H(rxn) for the overall equation is the sum of the heats of reactions for each step.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Trends in Lattice Energy

A
  • Energy becomes less endothermic (less negative) with increasing ionic radius.
  • Energy becomes more exothermic (more negative) with increasing magnitude of ionic charge.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Bonding Pair

A

Shared Pair of Electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Polar Covalent Bond

A
  • Intermediate between a pure covalent bond and an ionic bond.
  • ΔEN is intermediate (0.4-2.0)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Electronegativity

A

The ability of an atom to attract electrons to itself in a chemical bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Electronegativity characteristics

A
  1. Increases across a period in the periodic table.
  2. Decreases down a column in the periodic table.
  3. Inversely related to atomic size - the larger the atom, the less ability is has to attract electrons to itself in a chemical bond.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Bond Polarity

A
  • The degree of polarity in a chemical bond depends on the electronegativity difference (ΔEN) between the two bonding elements.
  • The greater the electronegativity difference, the more polar the bond.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Covalent Bond

A
  1. Nonpolar
  2. Two elements with identical electronegativities share electrons equally.
  3. ΔEN is small (0-0.4)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Ionic Bond

A
  1. Large electronegativity difference between the two elements in a bond.
  2. Metal and nonmetal.
  3. Electron from metal is almost completely transferred to the nonmetal. ΔEN is large (2.0+)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Dipole Moment (µ)

A

µ=qr

  • Occurs anytime there is a separation of positive and negative charge.
  • The smaller the magnitude of the charge separation & distance the charges are separated by, the smaller the dipole moment.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Resonance Structure

A

One of two or more Lewis structures that have the same skeletal formula (atoms are in same locations), but different electron arrangements.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Resonance Hybrid

A

The actual structure of the molecule is intermediate between the two (or more) resonance structures.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Formal Charge

A

Formal Charge = # valence electrons - (# of nonbonding + 1/2[# of bonding electrons])

The formal charge of an atom in a Lewis Structure is the charge it would have if all bonding electrons were shared equally between the bonded atoms.

The calculated charge for an atom if the effects of electronegativity were completely ignored.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Rules for Formal Charges

A
  1. The sum of all formal charges in a neutral molecule must be zero.
  2. The sum of all formal charges in an ion must equal the charge of the ion.
  3. Small (or zero) formal charges on individual atoms are better than large ones.
  4. When formal charge cannot be avoided, negative formal charge should reside on the most electronegative atom.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Exceptions to the Octet rule

A
  • Molecules/ions with an odd number of elections (7 e- around an atom)
  • nonmetals with incomplete octets
17
Q

Predicting Molecular Geometries

A
  1. Draw a Lewis Structure for the molecule.
  2. Determine the total number of electron groups around the central atom.
  3. Determine the number of bonding groups and the number of lone pairs around the central atom.
18
Q

Molecular Shape and Polarity

A
  1. Draw a Lewis Structure for the molecule.
  2. Determine whether the molecule contains polar bonds
    1. A bond is polar if the two bonding atoms have sufficiently different electronegativities (Fig. 9.8). i
  3. If the molecule contains polar bonds, superimpose a vector pointing towards the more electronegative atom, on each bond.
    1. Make the length on the vector proportional to the electronegativity difference between the bonding atoms.
    2. Determine whether the polar bonds add together to form a net dipole moment.
19
Q

Valence Bond Theory

A
  • Electrons reside in quantum-mechanical orbitals localized on individual atoms. Orbitals can be s, p, d, f or a hybrid combination.
  • A chemical bond results from the overlap of two half-filled orbitals with spin-pairing of the two valence electrons.
  • The shape of the molecule is determined by the geometry of the overlapping orbitals.
20
Q

Valence Bond Theory: Hybridization of Orbitals

A
  • Hybrid Orbitals are still localized on individual atoms, but have different shapes and energies from those of standard atomic orbitals.
  • The number of standard atomic orbitals added together always equals the number of hybrid orbitals formed; total number of orbitals is conserved.
  • The particular combinations of standard atomic orbitals added together determines the shapes and energies of the hybrid orbitals formed.
  • The particular type of hybridization that occurs is the one that yields the lowest overall energy for the molecule.
21
Q

sp3 Hybridization

A

Hybrid orbitals are mixtures of one s orbital and 3 p orbitals

22
Q

sp2 Hybridization

A

Hybridization of one s and two p orbitals results in three sp2 hybrids and one leftover unhybridized p orbital

23
Q

Pi (π) Bond

A

Occurs when p orbitals overlap side by side.

Electron density is above and below the internuclear axis.

24
Q

Sigma (σ) Bond

A

Occurs when orbitals overlap end to end.

25
Q

Boyle’s Law

A
  • Pressure-Volume Law
  • n & T are constant
  • As Pressure increases, Volume decreases
26
Q

Charles’ Law

A
  • Temperature-Volume Law
  • n & P constant
  • As temperature increases, volume decreases
27
Q

Avogadro’s Law

A
  • P & T constant
  • As n increases, volume increases
28
Q

Clausius-Clapeyron Equation

A

Relationship between Vapor Pressure, the Heat of Vaporization, and Temperature.

ln(Pvap) = (-ΔHvap / RT) + ln(β)

ln(P2 / P1) = (-ΔHvap / R)•[(1/T1) - (1/T2)]

29
Q

Bragg’s Law

A

Relationship between light wavelength (λ), angle of reflection (θ), and distance (d) between the atomic layers.

nλ = 2dsin(θ)

30
Q

Lone Pair / Nonbonding Electrons

A

Electron Pair associated with only one atom

CHEM104 (10 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions