ch 3 - atomic and molecular structure

Question 1

four main points regarding atoms

Answer 1

1. all matter is composed of atoms
   - atoms are typically indivisible & indestructible
2. atoms of a specific element are identical in mass and properties
3. compounds are formed by whole number ratios of two or more atoms
4. a chemical reaction is a rearrangement of atoms

Question 2

electron mass compared to proton

Answer 2

1/1836 of a proton,
- mass is considered negligible
- proton has a mass of 1 atomic mass unit (amu)

Question 3

Bohr’s model for describing the electronic structure of an atom

Answer 3

described electrons as following along defined circular orbits around the nucleus due to centripetal force
- from the attraction of a negatively charged electron to the positively charged protons of an atom

in the outdated Bohr model,
- electrons were thought to occupy distinct circular orbits around the nucleus (rings)

outdated and replaced by the modern quantum theory

Question 4

modern quantum theory

Answer 4

instead of orbiting around the nucleus in a defined circular pathway (Bohr’s model), electrons are actually localized in a “cloud of electrons” around the nucleus
- these regions of space are called orbitals
- Heisenberg uncertainty principle describes that it is impossible to perfectly find both the momentum and the location of a electron in an atom (i.e. can only find one or the other)

Question 5

Heisenberg uncertainty principle

Answer 5

describes that it is impossible to perfectly find both the (1) momentum and (2) location of an electron in an atom
- can only find out one or the other

Question 6

modern atomic theory

Answer 6

uses four quantum numbers that help describe the electrons of an atom:
1. principle quantum number (n)
2. azimuthal quantum number (l)
3. magnetic quantum number (ml)
4. spin quantum number (ms)

Pauli exclusion principle: states that no two electrons in an atom can have the exact same set of four quantum numbers

Question 7

modern atomic theory: principle quantum number (n)

Answer 7

represents the main energy level occupied by electrons
- is a positive integer number, equal to or greater than 1
- at n=1, an electron is closest to the nucleus, and with each successive electron shell, electrons get farther and farther away from the nucleus
- the maximum number of electrons that an electron shell can hold is given by the formula: 2n²

Question 8

modern atomic theory: azimuthal quantum number (l)

Answer 8

describes the shape of the subshells or the orbital shape within each principle energy level
- the possible values of the azimuthal quantum number are all between zero and the value of the principle quantum number minus 1
- e.g. a principle quantum number of 3 would have potential azimuthal quantum numbers of 0, 1, and 2

the subshells of the azimuthal quantum number carry a letter designation such that:
subshell l=0 is “s”
- can hold 2 electrons
subshell l=1 is “p”
- can hold 6 electrons
subshell l=2 is “d”
- can hold 10 electrons
subshell l=3 is “f’”
- can hold 14 electrons

to tie this together, and electron with values: n=3 and l=0 would be in the 3s subshell

Question 9

modern atomic theory: magnetic quantum number (ml)

Answer 9

describes the orientation of orbitals in space
- the magnetic quantum number ranges between the negative and positive magnitude of the azimuthal quantum number
e.g. p subshell with l=1
- would have ml values of -1, 0, and 1

Question 10

modern atomic theory: spin quantum number (ms)

Answer 10

describes the angular momentum of an electron
- denoted as either +1/2 or -1/2
- electrons in the same orbital must have antiparallel spins

Question 11

Pauli exclusion principle

Answer 11

applies to modern atomic theory
- states that no two electrons in an atom can have the exact same set of four quantum numbers

Question 12

electron configuration of a specific atom

Answer 12

tells us the number of electrons in each energy level and the organization of how the subshells are filled
1. first number describes the principle energy level followed by
2. a letter that describes the subsshell
3. a superscript that tells you the number of electrons in that specific subshell

e.g. 3s²
- this tells us that in the first (s) subshell of the third principle energy level, there are two electrons
- we also assume the subshells before 3s have already been filled, that is subshells 1s, 2s, and 2p

when writing the electron configuration notation, we tend to go from right to left, and top to bottom such that the order is:
- 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f
- this is because according to the Aufbau principle, subshells tend to get filled from lower energy to higher energy

Question 13

Aufbau principle

Answer 13

subshells tend to get filled from lower energy to higher energy
- when writing the electron configuration notation, we tend to go from right to left, and top to bottom such that the order is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f

Question 14

Q: what is the element whose ground state electron configuration is [Ne]3s²3p³

Answer 14

phosphorous
- due to the shorthand form, we can skip all the way to neon and start from there
- the ext part of the configuration is the 3s subshell which we see is completely filled (2 electrons)
- we then move on to the next (and last) subshell in the configuration, which is the 3p subshell
- we notice that there are only three electrons in this subshell, so this corresponds to the element phosphorous

note: whenever a noble gas is written in front of an electron configuration, this is used as a shortcut to avoid writing out the electron configuration of everything before that gas

Question 15

Q: a) what is the ground state electron configuration for Cr?
b) what is the ground state electron configuration for Cr2+?

Answer 15

a) 1s², 2s², 2p⁶, 3s², 3p⁶, 4s¹, 3d⁵
b)

a) we start by locating Cr on the periodic table, which is in the 3d subshell
- starting from the 1s² subshell, we pass through each subsequent subshell until we reach Cr, and we obtain the following configuration: 1s², 2s², 2p⁶, 3s², 3p⁶, 4s², 3d⁴

however, there is an exception, according to Hund’s rule, within a given subshell, orbitals are filled such that we have the maximum number of half-filled orbitals
- to satisfy this, an electron from the 2s subshell will go to the 3d subshell, such that we have a half filled 4s orbital and a half filled 3d orbital (note d orbital can hold 10 electrons), so the proper ground state electron configuration would be: 1s², 2s², 2p⁶, 3s², 3p⁶, 4s¹, 3d⁵

b) now to find the electron configuration of Cr²⁺, we must simply remove 2 electrons
- we dothis by removing electrons, first from the highest energy level (n value) which is the 4s subshell, then from the 3d subshell, we get:
1s², 2s², 2p⁶, 3s², 3p⁶, 4s⁰, 3d⁴

note: other special cases that you should know for the DAT include: copper (Cu), silver (Ag), gold (Au), and molybdenum (Mo)

Question 16

covalent bonds vs. ionic bonds on melting and boiling ponts

Answer 16

covalent bonds are much weaker than ionic bonds
- expect covalent compounds to generally have lower melting and boiling points

Question 17

Q: what is the bond length of a C-H bond, if we know that the bond length of a hydrogen molecule, H2, is 74pm and a C-C bond is 154pm?

Answer 17

114pm
- to find the bond length of a C-H bond, we must first find the atomic radii of the two atoms involved
- we can do this by dividing the given bond lengths by two because in both cases the two atoms bound together are identical:
* atomic radius of carbon: 154/2 = 77pm
* atomic radius of hydrogen: 74/2 = 37pm
- then we simply add these atomic radii to find the bond length of the C-H bond
77pm + 37pm = 114pm

Question 18

how does the transfer of electrons work in ionic bonds

Answer 18

complete transfer of electrons from the less electronegative atom to the more electronegative atom
- metal elements do not have a very high affinity for electrons and so they have low ionization energy
- the cations (atom that lost electrons) and anions (atom that gained electrons) that result form ionic bonding are then held together by the electrostatic attraction of their charges

Question 19

intermolecular forces

Answer 19

the weak electrostatic interactions between atoms and compounds
- intermolecular forces are significantly weaker than both ionic and covalent bonds

three types that you should know
1. london dispersion forces (Van der Waals force)
2. dipole-dipole interactions
3. hydrogen bonds

Question 20

intermolecular forces: london dispersion forces

Answer 20

also called Van der Waals forces, the weakest interactions
- occurs in all compounds
- non-polar compounds only experience London dispersion forces

dispersion forces arise when the electron density between two atoms becomes unequally distributed for a very brief moment
- resulting in instantaneous dipole moments in a molecule that doesn’t have dipoles
- these short-lived dipoles induce short-lived dipoles in other neighboring molecules

general rule of thumb
- the number of dispersion forces increases as the molecule gets larger

Question 21

intermolecular forces: dipole-dipole interactions

Answer 21

found in both polar solid and polar liquid compounds (but not gases)
- essentially the rearrangement of the partially charged end of one molecule such that it moves closer to the opposite partially charged end of another molecule

similar to London dispersion forces in that they both form dipole-dipole interactions
- except these last much longer than Van der Waals dipoles

Question 22

intermolecular forces: hydrogen bonds

Answer 22

strong dipole-dipole interactions that occur when hydrogen is attached to a very electronegative atom (e.g. oxygen, nitrogen, or fluorine)
- as a result, the electronegative atom takes most of the electron density (becomes partially negative) and this leaves hydrogen with a partial positive charge
- the partially positive hydrogen is able to have electrostatic interactions with the partially negative N, O, or F atoms of neighboring compounds

Question 23

important rules for drawing Lewis dot diagrams

Answer 23

1. arrange the atoms that the least electronegative atom is the central atom that the other atoms are bound to
   - hydrogen and the halogens are almost never a central atom
2. sum up the valence electrons of all the atoms in the compound and draw in bonds
3. assign electron pairs to the atoms attached to the central atom to complete their octets
4. any electrons left over can go on the central atom
   - must form double o triple bonds until octet is satisfied too
5. calculate formal charge using following equation
   formal charge = (# of valence e-) – (# of nonbonding e-) – (# of bonds)

Question 24

Lewis dot structure: exceptions to the octet rule

Answer 24

cannot obtain an octet:
- hydrogen, helium, lithium, beryllium, boron

can exceed the octet rule:
- elements in the third period and below, because they have more than 8 valence electrons due to the presence of d subshell orbitals

Question 25

azimuthal quantum number: subshell and orbital shape in each energy level

Answer 25

l = 0, 1, 2, 3 corresponds to the s, p, d, and f subshells respectively
1. s subshell is a single orbital with a spherical shape
2. p subshell has 3 orbitals with a barbell shape, and lie on the x, y, and z-axes
3. the d and f subshells have 5 and 7 orbitals, respectively
- much more complex than what is expected on the DAT

Question 26

bonding between two atoms results in what to their atomic orbitals

Answer 26

results in the overlap of their atomic orbitals
- leads to the formation of more complex molecular orbitals

Question 27

do sigma and pi bonds allow for free rotation at the axis?

Answer 27

sigma bonds do allow for free rotation at the axis
- pi bonds do not

Question 28

bond strength depends on what factors?

Answer 28

1. bond order
   - single bond has a bond order of one, double bond = 2, triple bond = 3
   - higher bond order equals a decrease in bond length and increases bond strength
2. atomic radii
   - bond strength increases as we go higher up a column
   - because when you go down a column, atomic radii increases and increases the bond length, decreasing bond strength
   - smaller atoms have more overlap of their orbitals than larger atoms with greater atomic radii
   e.g. H-O stronger than H-S, which is stronger than H-Se
3. polarity
   - greater the difference in electronegativity between two atoms usually equates in a stronger bond
   - because a large difference in electronegativity results in a polar bond, leading to the formation of partial positive and partial negative charges
   - these partial charges have an electrostatic attraction which increases the strength of the bond
4. lone pairs
   - bonds that have more lone pairs on their atoms will generally have lower bond strength as the repulsion between the lone pairs weakens the covalent bond

Question 29

Q: order the following bonds from strongest to weakest
H-Cl, H-F, H-I, H-Br

Answer 29

H-F > H-Cl > H-Br > H-I
- all four varying atoms lie in the same column, therefore based off of atomic radii we get the above order
- note: H-I is the strongest acid

Question 30

Q: order the following bonds from strongest to weakest
C-C, C-N, C-F

Answer 30

C-F > C-N > C-C
- this is a simple comparison of the electronegativity differences between the atoms
- we know that there will be the smallest electronegativity difference in the C-C bond and the greatest difference in the C-F bond