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204 Class > Chapter 2 - Atomic Structure and Interatomic Bonding > Flashcards

Chapter 2 - Atomic Structure and Interatomic Bonding Flashcards

1
Q

How can you find the Atomic number?

A

of protons in the nucleus

(or periodic table’s BIG counting number)

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2
Q

How can you find Atomic mass?

A

sum of masses of protons and neutrons in the nucleus

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3
Q

What is atomic weight? What is atomic mass?

A

Atomic weight → average mass of an element icluding the percentages of their isotopes.

atomic mass → mass of just the primary isotope.

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4
Q

Electrons exibit properties of what two things?

A

particles, and waves.

in particular, momentum and interference

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5
Q

Electrons occupy ___ in atoms

A

energy states.

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6
Q

If both a higher and a lower energy state are unoccupied by electrons, which of the two states would a new electron go to?

A

The lowest state.

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7
Q

What is the general equation for energy states?

A

nlx

where:
n = principal electron shell of the atom
l = describes the shape of orbital
x = number of electrons in shell

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8
Q

What are the different azimuthal quantum numbers (l)? (put them into order)

A

s,p,d, and f

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9
Q

how many possibilities are there for a s-orbital?

A

1

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10
Q

How many posibilties are there for a p orbital?

A

3

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11
Q

How many possibilities are there for the d orbital?

A

5

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12
Q

How many possibilites are there for the f orbital?

A

7

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13
Q

How many electrons can the s orbital hold?

A

2

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14
Q

How many electrons can the p orbital hold?

A

6

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15
Q

How many electrons can the d orbital hold?

A

10

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16
Q

How many electrons can the f orbital hold?

A

14

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17
Q

which azimuthal quantum numbers (l) will be in each principal quantum number (n) up to 4?

A
  1. s
  2. s,p
  3. s,p,d
  4. s,p,d,f

5,6,7…n → would be all be the same as 4

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18
Q

what are Valence electrons?

A

electrons in unflilled primary shells

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19
Q

Filled electron shells are more stable or unstable?

Do they require more or less energy to gain/lose electrons

A

are more stable and require more energy to lose electrons.

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20
Q

Using Aufbau’s princible state which sub shells would fill first up to 5 d

A

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d

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21
Q

Which side & corner has the lowest electronegativities?

Is the top, or bottom of that side the lowest?

A

left side, bottom corner.

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22
Q

Which side & corner has the highest electronegativities?

Is the top, or bottom of that side the highest?

A

Right side, top corner.

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23
Q

List the three types of primary bonding from highest to lowest bonding energies.

A
  1. Ionic
  2. Covalent
  3. Metallic

covalent can be split into smaller sub-sections

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24
Q

List the two types of secondary bonding from highest to lowest bonding energies.

A
  1. Hydrogen
  2. Van der Waals

van der waals can be split into smaller sub-sections

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25
Q

Why do atoms want to bond together?

A

They want to obtain a stable electron structure.

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26
Q

How does an Ionic bond form?

To form, elements need a large difference in their ____

To have greater bond straingth you need _____

A

When a metal donates an electron, and a non-metal receives an electron

electronegativity

↑ the difference in electronegativity

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27
Q

For Lewis structure which side of the element name do you start?

Which way should you add electrons?

A

Right

Clockwise

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28
Q

How does a Covalent bond form?

To form, elements need a small difference in their ____

Bond straingth increases with ___

A

When 2 non-metals (including things on the staircase) share electrons in their s and p orbitals

electronegativity

the ↑ difference in electronegativity, becoming more ionic

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29
Q

Define Bond Hybridization.

A

The mixing (or combining) of two or more atomic orbitals which results in more orbital overlap during bonding.

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30
Q

What is the point of Bond Hybridization?

A

It allows atoms to have a larger amount of valence electrons, increasing covalent bonding capabilities.

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31
Q

Describe the process of sp3 orbitals in carbon.

A
  1. promotion of electron - the down spin 2s electron gets promoted to 2p up spin in a empty 2p shell
  2. sp3 Hybridization - the 2s up spin electron orbital moves up and the rest of the 2p orbitals move down creating a 2sp3 orbital
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32
Q

Describe the process of sp2 orbitals in carbon.

A
  1. promotion of electron - the down spin 2s electron gets promoted to 2p up spin in a empty shell
  2. sp3 Hybridization - the 2s up spin electron moves up and all but one electron orbital in 2p moves down creating a 2sp2 orbital
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33
Q

Describe the bonds in sp3 and name one example.

A

4 single sigma bonds.

Diamond

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34
Q

Describe the bonds in sp2 and what the bonding allows for. In addition, name one example.

A

1 sigma bond, 1 pi
bond which allows double bond formation

graphite

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35
Q

Describe the bonds in sp and what the bonding allows for. In addition, name one example.

A

1 sigma bond, 2 pi
bonds which allows triple bond formation

C2H2 (ethane)

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36
Q

In VSEPR theory, greater repulsion of electron pairs shows…

A

higher energy and less stable electron pairs

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37
Q

rank single, double and triple molecular bonds based on closeness.

A
  1. triple
  2. double
  3. single
38
Q

rank single, double and triple molecular bonds based on straingth.

A
  1. triple
  2. double
  3. single
39
Q

Double and triple bonds have a stronger repulsive force than single bonds. Their relative strength is similar to ____

A

lone pairs.

40
Q

What is metallic bonding?

A

electrons around atoms delocalize and form a “electron cloud”

41
Q

list the types of bonding that is directional and non-directional?

A

Directional bonding → covalent and secondary bonding

Non-directional bonding → ionic and metallic bonding

42
Q

List the boiling points of the different types of primary bonding ( Ionic bonding, Nonpolar covalent, Hydrogen bonding, Metallic bonding, and Polar covalent).

A

Nonpolar covalent < Polar covalent < polar Hydrogen bonding < Metallic bonding < Ionic bonding.

43
Q

List the bonding strength of the different types of primary bonding ( Ionic bonding, Nonpolar covalent, Metallic bonding, and Polar covalent).

largest to smallest

A
  1. Ionic bonding
  2. Polar covalent bonding
  3. Nonpolar covalent bonding
  4. Metallic bonding
44
Q

What are Intermolecular and Intramolecular bonds?

A

Intermolecular – between atoms or groups of atoms
Intramolecular – primary bonds

45
Q

Are secondary bonds relatively weak or strong?

A

Relatively weak.

46
Q

How does secondary bonding arise?

A

from atomic or molecular dipoles.

47
Q

What are the three types of dipoles?

A

Fluctuating Dipoles → dipoles that change due to momentary changes in electron clouds

Permanent dipoles → Dipoles that form between covalent atoms from net positive and negative charges that are respectively associated with the different ends of a molecule

Hydrogen bonds → a special type of permenant dipole

HCl the H is positively charged and chlorine is negatively charged

48
Q

What is the definition of directional bonding?

A

Electrons or shared electrons are localized in a specific area between atoms.

49
Q

What is the definition of non-directional bonding?

A

Electrons cannot be localized in a specific location between atoms

Bonding forces are the same in all directions around the atom

50
Q

What is %IC? What are the variables XA & XB stand for in this formula?

%𝐼𝐶 = {1 − 𝑒𝑥𝑝 [− 1/4 (𝑋𝐴 − 𝑋𝐵)^2]} × 100

A

%IC → % Ionic character

XA & XB → the electronegativities of the two elements participating in the bond

51
Q

List the relative strength of intermolecular forces of attraction.

A

Hydrogen bonding, Dipole-dipole attraction, London dispersion attraction

52
Q

What is another term for london dispersion attraction.

A

Momentary dipoles

53
Q

what is FN, FA, & FR

A

Net force
Force of attraction
Force of repulsion

54
Q

What is the force formula if the particle is at an equilibrium distance from another one?

A

0 = FA + FR

55
Q

What is this “𝐸(r) = int(𝐹)𝑑𝑟” formula saying

A

That energy is the integral of force with respect to separation distance.

56
Q

Particles apply ___ and ___ forces on other particles that come too close or are too far

A

attractive, repulsive

57
Q

How do we know which interatomic graph talks about potential energy or force?

A

For force, the graph has attraction as up, and the graph heads to -∞

For potential energy, the graph has attraction down, and the graph heads to ∞

58
Q

When do particles have the least amount of energy?

A

when at Equilibrium.

NOTE: Energy is never at zero

59
Q

What is the formula EN = EA + ER mean?

A

Net energy is equal to the sum of attraction and repulsive energies.

60
Q

What is the one equation you need to know to find Net energy in an interatomic system algebraically?

A

EN = -A / r + B / r^n

where N will be given in the question, B is a constant, and A is determined algebraically on the formula sheet.

61
Q

In this formula 𝐴 = 1/(4𝜋𝜀0)(𝑍1|𝑒)(|𝑍2|𝑒) What do the terms mean?

A

A = the value nessesary to find Energy of attraction

𝜀0 = Bond energy (ussally given)

𝑍1,2 = charges of ions / atoms

62
Q

A larger bond energy (𝜀0) would result in an increase or decrease in melting temperature (Tm)

A

increase.

63
Q

A larger bond energy (𝜀0) would result in an increase or decrease in thermal expansion (αl)

A

decrease.

64
Q

List the relative bond energy of Ceramics, Polymers, and Metals from largest to smallest.

A

Ceramics, Metals (can be quite variable though), Polymers.

65
Q

List the relative Thermal expansion of Ceramics, Polymers, and Metals from largest to smallest.

A

Polymers, metals, ceramics

66
Q

List the relative net energies of Ceramics, Polymers, and Metals from largest to smallest.

A

Ceramics, Metals (can be quite variable though), Polymers.

67
Q

List the relative melting temperatures of Ceramics, Polymers, and Metals from largest to smallest.

A

Ceramics, Metals (can be quite variable though), Polymers.

68
Q

What type of bonding is responsible for most physical properties in Polymers?

A

Secondary bonding

69
Q

Describe van der Waals bonding. Where does it matter where they occur?

A

Arise due to atomic or molecular dipoles (some separation of + and – charges in a
particle) where the + of one particle is attracted to the – of another.

it matters between uncharged atoms or molecules (inert gasses) where there are very little other forces happening.

70
Q

Describe hydrogen bonding. Where does it usually occur?

A

Is present when a hydrogen atom is bound to a very electronegative molecule. The + H
is attracted to the other molecules around it because it’s a unshielded proton.

where Hydrogen is covalently bonded with N, O, F,

71
Q

Which type of intermolecular bonding is typically seen in a material with a relatively high or low bonding point?

A

ionic and covalent respectively.

72
Q

Bonds usually form between two atoms such that they are separated by a distance where their net force and energy is

A

zero and minimized respectively.

73
Q

Draw a graph of force and energy vs interatomic separation (r) for two isolated atoms. Include plots
for attractive (FA), repulsive (FR), and net forces (FN), as well as label the equilibrium spacing (r0).

A

Did you do it?

74
Q

Stronger bonds have higher or lower electronegativity for both ionic compounds and covalent.

A

Both covalent and ionic bonds are stronger and have larger electronegativities. (exceptions are bond hybridization and hydrogen bonding.

75
Q

For covalent bonding, what things produce the strongest bond

A

The closer the elements are to each other (smaller size difference), amount of electrons that are shared (amount of bonds)

Larger changes in electronegativity

76
Q

Can you draw a lewis structure? If you can’t here is your reminder to try one.

A

:)

77
Q

The wave-particle duality shows which two properties of electrons

A

Electrons have momentum, and can cause interference patterns

78
Q

What do you have to do to energy to find force?

A

Take the derivative

79
Q

What type(s) of bonding would be expected for steel (an alloy composed of iron
with a very small amount of carbon)?

A

Metallic bonding

Note: this is a exam question, they are looking for most prevalient bond

80
Q

Which sections in the periodic table is the s block?

A

first two rows

81
Q

which section is the p block

A

all the non-metals

82
Q

for the weird changes in shell number (3s,2d) How to figure out what shelled is filled

A

always the lowest energy first

But then it wants to either fully or half fill the higher energy shell if possible

83
Q

Look at the different shapes of the s,p,d, and f orbitals.

A

Did you do it?

84
Q

If copper had 6,8, and 11 electrons in valence. How would they be put into the last 4s and 3d shells?

A

4s1 and 3d5

4s2 and 3d6

4s1 and 3d10

(trying to show how valance electrons want to get higher energy full or half shells)

85
Q

What is stronger, a sigma or a pi bond?

A

sigma

86
Q

List the relative straingths of the Spn structures.

A

sp3, sp, sp2

87
Q

What is the atomic number of H?

A

1

88
Q

What is the atomic number of He?

A

2

89
Q

What is the atomic number of O?

A

8

90
Q

What is the atomic number of C?

A

6

91
Q

What is the atomic number of Fe?

A

26

92
Q

What is the atomic number of Al?

A

13

204 Class (9 decks)

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