Chapter 2 - Historical Development of Atomic Theory

Question 1

What did Dalton discover?

Answer 1

Atoms combine in simple numerical ratios to form compounds

Question 2

What did Avogadro discover?

Answer 2

Equal volumes of gas at equal temperatures and pressures contain the same number of molecules

Question 3

What did Thomson discover?

Answer 3

JJ Thomson came up with the plum pudding model (stable thing with “blueberries” in it), which says that an atom has electrons surrounded by a soup of positive charge. He demonstrated that atoms are actually composed of aggregates of charged particles. Prior to his work, it was believed that atoms were the fundamental building blocks of matter. He was able to determine the charge to mass ratio of the electron.

Question 4

What did Rutherford discover?

Answer 4

Rutherford overturned Thomson’s model in 1911 with his well-known gold foil experiment in which he demonstrated that the atom has a tiny, heavy nucleus. He also came up with the mass of the electron.

The Rutherford model is that the atom is made up of a central charge (this is the modern atomic nucleus, though Rutherford did not use the term “nucleus” in his paper) surrounded by a cloud of (presumably) orbiting electrons.

Question 5

What did Canizzaro discover?

Answer 5

In 1860, he came up with a consistent set of atomic weights. Because each molecule is made up of atoms, the sum of the atomic weights can then give us the molecular weight.

Question 6

What did Mendeleev + Meyer discover?

Answer 6

They came up with the periodic table. They organized the periodic table by similar properties.

Question 7

Period

Answer 7

rows in periodic table

Question 8

Group

Answer 8

columns, also known as “families” in the periodic table

Question 9

IUPAC vs US groups?

Answer 9

IUPAC: groups numbered 1–18
US: groups IA–VIII main group, “B” group are transition metals

Question 10

Balmer

Answer 10

Balmer showed that energies of light emitted by the hydrogen atom are given by the equation:
E = R ( 1/2^2 - 1/n^2) = R ( 1/nl^2 - 1/nh^2)

where R = Rydberg constant = 1.097E7 m-1 = 2.179E-18 J

Question 11

Energy related to wavelength, frequency, wavenumber

Answer 11

E = hv = hc/λ

Question 12

de Broglie

Answer 12

λ = h/mv

Question 13

Heisenberg Uncertainty principle

Answer 13

there is a relationship between the inherent uncertainties in the location and momentum of an electron.

∆x∆p ≥ h/4π

Therefore, since we cannot know the position exactly, we cannot use ORBITS but must use ORBITALS, regions to descibre the probable location of electrons.

Question 14

Electron density

Answer 14

the probability of finding an electron at a particular point in space

Question 15

What are some problems with the Bohr model of the atom?

Answer 15

• works well only when one electron is involved
• spherical –> elliptical orbitals introduced to fit Bohr’s data
• did not take into the wave-like nature of the electron (de Broglie)
• Heisenberg uncertainty principle means that we cannot know the position of the electron exactly. Therefore, we cannot use orbits, but must use orbitals instead to describe the probable location of electrons.

Question 16

Schrodinger Equation

Answer 16

• tells us the wave properties of an electron, its mass, position, total energy, and potential energy
• based on the wavefunction ψ, which describes the wave properties of a given electron in a particular orbital

Hψ = Eψ

Question 17

What is the relationship between atomic orbitals and ψ?

Answer 17

Atomic orbitals are described by a unique ψ, so there is no limit to the number of solutions for the Schrödinger equation.

Question 18

Bohr’s Quantum Theory of the Atom

Answer 18

Negatively charged electrons in atoms moves in stable circular orbits around the positively charge nucleus with no absorption or emission of energy.

The energy emitted/absorbed by electron can be found by:

E = R ( 1/nl^2 - 1/nh^2)

Question 19

µ in the definition of the R used in the Bohr theory?

Answer 19

Reduced mass of the electron/nucleus combination.

1/µ = 1 / mass of electron + 1/mass of nucleus

Question 20

L

Answer 20

Angular momentum/shape of the orbital. Describes angular dependence and contributes to energy.L = 0,1,2,…, n-1

Question 21

M_L

Answer 21

Orientation of the angular momentum vector in a magnetic field. Describes orientation in space (angular momentum in z direction).m_l = -l, …., -2, -1, 0, 1, 2, …, +l

Question 22

M_s

Answer 22

Spin. Orientation of the electron’s magnetic moment (spin) in a magnetic field, either in direction of the field (+1/2) or opposed to it (-1/2).m_s = -1/2, +1/2

Question 23

What l value describes:1. S2. P3. D4. F5. G

Answer 23

1. 2. 13. 24. 35. 4

Question 24

N

Answer 24

Principal quantum number (n) is primarily responsible for determining the overall energy of an atomic orbitaln = 1,2,3,…

Question 25

Radial functions

Answer 25

the radial part of the wavefunction: electron density at different distances (r) from the nucleus

• determined by n and l
• radial prodbability function = 3πr^2 R^2

Question 26

Radial nodes

Answer 26

radial nodes for an orbital = n - l - 1

when R, the radial function, = 0

(counts a conical node surface such as for a d orbital as two nodes)

Question 27

Angular functions

Answer 27

the angular part, Y(ϴ,ϕ). describes the shape of orbitals and their orientation in space

• determined by l and ml

Question 28

Angular nodes

Answer 28

when Y=0, are planar/conical

Question 29

Nodal surface

Answer 29

A surface with zero electron density

Question 30

Node

Answer 30

A surface where the wavefunction is zero as it changes sign

ψ=0 so either R=0 or Y=0

Question 31

Total number of nodes in any orbital?

Answer 31

n-1

Question 32

Total number of radial nodes for any orbital?

Answer 32

n-l-1

Question 33

Total number of angular nodes for any orbital?

Answer 33

l

Question 34

Aufbau principle

Answer 34

Limitations on the values of quantum numbers –> Aufbau (building up) principle, where the buildup of electrons in atoms results from continuously increasing quantum numbers

1. Electrons placed in give lowest total energy
2. Pauli exclusion principle (each electron has unique set of quantum numbers)
3. Hund’s rule of maximum multiplicity (electrons placed to have maximum total spin)

Question 35

Pauli Exclusion principle

Answer 35

Each electron has a unique set of quantum numbers

Question 36

Hund’s rule

Answer 36

Electrons placed to have maximum total spin.

• comes from the coulombic energy of repulsion (πc) vs the exchange energy (πe)

Question 37

Coulombic energy of repulsion

Answer 37

destabilizing effect caused from electron repulsion between electrons in the same orbital

Question 38

Exchange energy

Answer 38

stabilizing effect of increased number of possible exchanges, a consequence of parallel electron spins in separate orbitals.

Question 39

Klechkowsky’s rule

Answer 39

The order of filling of the orbitals proceeds from the lowest available value for the sum n+l

Question 40

Shielding

Answer 40

each electron can act as a “shield” for electrons farther from the nucleus, reducing the attraction between the nucleus and the more distant electrons

Z* = Z - S

where Z* = effective nuclear charge
S = sheilding constant, determined using Slater’s rules

Question 41

Slater’s rules

Answer 41

rules for determining the shielding constant for a specific electron

Question 42

For transition metal atoms, which electrons will be lost more readily?

Answer 42

ns electrons are lost more readily than (n-1)d electrons. This is because the effective nuclear charge for the ns electrons are smaller than the values for the (n-1)d electrons, so they are held less tightly.

Question 43

For Cu and Cr, what is the exception for shielding?

Answer 43

Half and fully filled d and f shells are preferred, so Cr has a configuration of [Ar]4s13d5 instead of [Ar]4s2d4.

Question 44

Ionization energy

Answer 44

Energy required to remove an electron from a a gaseous atom/ion

• generally increases across a period
• break at Boron due to shielding

Question 45

Electron affinity

Answer 45

Energy required to remove an electron from a negative ion (zeroth ionization energy)

Question 46

What has the lowest electron affinities?

Answer 46

Noble gases (an extra electron beyond the noble gas configuration is easily lost)

Question 47

Anions are ____ than cations with similar numbers of electrons.

Answer 47

Larger (e.g. electron repulsion)

Question 48

Covalent radii/atomic size changes?

Answer 48

Increasing nuclear charge + # electrons –> decrease in atomic size across period