Chem 101b- Atoms and Molecules

Question 1

which theory did dalton base his theory on?

Answer 1

the law of conservation of mass: in a chemical reaction, mass can’t be created or destroyed, only transferred.

Question 2

give basic overview of dalton’s theory

Answer 2

all matter must be composed of atoms, and atoms of different elements have different mass, altering how the elements are grouped into the periodic table.
Dalton linked the microscopic view of ratio of atoms to macroscopic things e.g. grams of water.

Question 3

give basic overview of Thomson’s theory + his experiment

Answer 3

atoms aren’t indivisible.
his experiment: high current was applied to gas, electrons fly from one electrode to the other- negatively charged particle is released. If a neutral atom can release negatively-charged electrons, there must be a positively-charged one.
It was discovered that electrons have the same mass in any element, and the +ve particle had a larger mass than the -ve electron, and there was a multiplicity of electrons.

Question 4

describe Rutherford and geiger’s experiment

Answer 4

fired alpha particles through thin leaves of gold that are a few atoms thick. (alpha particles= helium atoms with 2+ charge). Most went through, only 1/20,000 were deflected. This shows most of the atom is empty space, the radius is much bigger than just the nucleus, and the nucleus is +ve as it deflected the +ve helium atom. Also that the plum pudding isn’t correct as they all should’ve been deflected or absorbed.

Question 5

draw Rutherford’s model of the atom, and why it isn’t possible in the interest of classical physics

Answer 5

Rutherford’s model shows electrons are “somewhere else”?? This model isn’t possible with classical physics, as an electron didn’t have a defined orbit, and according to physics it should spiral and decay into the nucleus or something while giving off high levels of x-ray radiation. Also shouldn’t be stable, they didn’t know what bound the electron and nucleus together. We weren’t aware of strong nuclear force, quarks and neurons.

Question 6

what are 3 problems with Rutherford’s model?

Answer 6

1. no x-rays emitted
2. no collapse as energy runs out
3. spectra, only certain electron energy states are possible within the atom as the atom can only gain or lose energy in fixed quanta

Question 7

how does the atomic electromagnetic radiation spectrum suggest that emission of radiation of atoms in quantised, therefore meaning there are certain energy levels in the atom?

Answer 7

The continuous spectrum shows white light shone through a prism, when light scatters due to raindrops (rainbow). If light is sent through cloud of hydrogen gas, comes out other end and absorbs very specific frequencies. If a hydrogen atom is excited, it only emits those 4 specific frequencies out of all of them that the sun emits. From the previous model, the emission should look like the continuous spectrum, with no absorption. Suggess the absorption or emission of radiation by atoms is quantized. Means must be certain energy levels in atom.

Question 8

difference between Rutherford and Bohr’s view on electron behaviour

Answer 8

According to Rutherford, electrons could be anywhere and emission is continuous, but Bohr said electrons possess fixed quantised energies, each energy level characterised by quantum number “n”, and electrons orbit at different energy levels.
The 4 series in the Balmer series are equal to the 4 lines in the emission spectra.

Question 9

What do quantum numbers define?

Answer 9

periods in the periodic table, and accounts for the difference in chemistry between eg carbon and lead

Question 10

Einstein did an experiment demonstrating the quantum behaviour of radiation, shining UV light onto a metal surface in a vacuum. What did it prove?

Answer 10

radiation frequency must be above a threshold (specific to metal) for electrons to eject. Increasing the radiation frequency (i.e. decreasing wavelength) above threshold doesn’t lead to more e- being ejected, just the e- having greater kinetic energy. Instead, increasing light intensity increases number of e-.
In certain situations, light is a particle, not a wave.

Question 11

what does the double slit experiment prove?

Answer 11

demonstrates that radiation can behave like a wave: wave-particle duality!! Defraction occurs when a beam of light or electrons is shone through the slits, which explains why quantisation occurs and how bonding works in atoms.

Question 12

what is the de Broglie equation and what does it relate?

Answer 12

Relates particle and wave behaviour.
Says that the wavelength associated with a particle is related to its momentum.
Combining E=hν (relating to wave behaviour) and E=mc2 (relating to particle behaviour) gives hν=mc2. We know that ν=c/λ (substitute in c/λ for ν) so hc/λ=mc2 (divide through by c) then h/λ=mc.
We also know that p = mc (substitute in p for mc, p = momentum)
So: h / λ = p

Question 13

what is the equation relating speed, wavelength and frequency?

Answer 13

speed= wavelength x frequency (so, light = wavelength x frequency)

Question 14

draw constructive and destructive wave interference, and describe how they can be in phase

Answer 14

constructive= in phase (can be either positive or negative)
destructive= not in phase

Question 15

at a specific distance from the nucleus (r), what does it mean if the electron wave overlaps destructively with the route around the nucleus?

Answer 15

there cannot be an electron orbit at that distance, as the frequency of the wave of the electron needs to match with the radius of the atom in the equation (2 x pi x r= n x wavelength). only certain wavelengths are allowed, i.e. quantisation of energy. Explains why atomic spectra show lines instead of continuous output.

Question 16

what is the relationship between electron position in an atom, and electron momentum (Heisenberg uncertainty principle)?

Answer 16

the more precisely the electron position is determined (making delta smaller), the less precisely the electron momentum is known in this instance (increasing error in momentum), and vice versa. So we can’t really know the exact location as electrons are delocalised, but we can know the probabilities in different areas of electrons existing there.

Question 17

what is bohr theory in terms of electron momentum, and what is the issue with him?

Answer 17

fixed orbits should allow electron momentum and position to be determined. However, in reality, electrons can be anywhere due to their wave nature, and new maths is needed to consider it.,

Question 18

Wacefunction for Ana electron can be calculated with the Schrodinger equation, what does it involve?

Answer 18

The big bit is a kinetic energy term: how wave function changes with distance from the nucleus.
Epe: total potential energy (repulsions and attractions) associated with wave function.
E: total energy associated with wave function.
Wavelength (squared) is proportional to the probability of finding an electron within a small volume of space.

Question 19

describe cartesian and spherical coordinates

Answer 19

Question 20

which 3 quantum numbers emerge from the Schrodinger equation, and what do they mean for the atom?

Answer 20

n, l, ml. Together, the characterise an atomic orbital (a region of space defined by a wave function).

Question 21

what are radial and angular wave functions?

Answer 21

R: depends only on distance (r) from nucleus
Y: depends only on orientation (angles)

Question 22

what does quantum number “n” tell you? (principle)

Answer 22

the energy of an atomic orbital, comes from energy level shells e.g. n=1, n=2.

Question 23

what does the quantum number “l” tell you? (orbital)

Answer 23

shape of an atomic orbital e.g. dumbbell, spherical. same “n” and “l”= same sub-shell e..g 2p.

Question 24

what does the quantum number “ml” tell us? (magnetic)

Answer 24

orientation of an atomic orbital, kinda defines the number of orbital present.

Question 25

what does it mean when we say all 3 possible p-orbitals are degenerate?

Answer 25

they all have the same energy states

Question 26

what is radial wave function?

Answer 26

defines what happens to wavelength at a specific distance (r), basically what happens to wave function with distance from the nucleus.

Question 27

describe radial wave function at n=1

Answer 27

wave function is at maximum at the nucleus, and radial wave function reaches 0 only when distance from nucleus = infinite

Question 28

describe radial wave function when n > 1

Answer 28

wave function can be + and - and 0. if radial wave function is 0 at a point value r, this is a radial node (wave function changes sign). Number of nodes for when distance is over 0= n-l-1.

Question 29

describe the 2 types of nodes that can occur in atomic orbitals

Answer 29

radial: spherical node surfaces, defined by "n". angular: node planes or cones. defined by "l". Therefore, s-orbitals don't have any angular nodes. All s orbitals are spherical. (Neither θ nor φ appear in Y for a hydrogen 1s orbital, ie it is the same in all directions from the nucleus.)

Question 30

how does the intensity of orbital energy affect the number of nodes?

Answer 30

The higher the energy of the orbitals, the greater the number of nodes. At n=1, no nodes. At n=2, 1 node. At n=3, 2 nodes.

Question 31

draw a radial wave function graph for a 1s orbital

Answer 31

at the nucleus, the value of R(r) for any s-orbital is high, then drops towards the x-axis.

Question 32

what does radial wavefunction squared mean?

Answer 32

radial distribution (probability of finding electron at a given point). If you square a negative, it becomes positive, so no negative values on the graph anymore. All R(r)2 (especially s) have highest values / maxima near nucleus, but not weighted by RADIUS (2D graph – reality 3D) At r = 0: high probability, but only 1 point that the electron can occupy At r = small: few points each with high probability At R = large: many points with lower probability

Question 33

why is radial distribution function more epic than radial distribution?

Answer 33

radial distribution function- 4πr2R(r)2: function of electron density in a region. the maximum represents the most probable distance from the nucleus of finding an electron. Multiplying R(r)2 by 4πr2 has two main effects: Maxima are shifted further from nucleus because 4πr2 increases with r  eg highest probability to find e- in an s orbital is NOT in nucleus! For all atomic orbitals 4πr2R(r)2 is zero at r = 0, i.e. in the nucleus (because R(r)2 is multiped by r). Takes into account "surface area" of orbital i.e. a 3s orbital will have a larger surface area than 1s.

Question 34

what does l mean in terms of nodal planes?

Answer 34

number of orbital angular nodes (nodal planes)

Question 35

what is a boundary surface?

Answer 35

lines/ surfaces are drawn at 95% probability (probability only reaches 0 at r= infinite). where the graph approaches 0, very small probability of finding an electron is extremely close to 0= the surface of the drawn orbital.

Question 36

how are phases of the wavefunctions (+ or -) shown on boundary surfaces?

Answer 36

shown by shading of the boundary surfaces, s-orbitals have the same phase over the whole boundary surface. p and d orbitals contain both positive and negative phases.

Question 37

draw all 5 possible drawings for d-atomic orbitals (due to 5 quantum states)

Answer 37

Question 38

what equation relates mass, speed, Planck's constant and wavelength?

Answer 38

wavelength= h/ (mass x speed)

Question 39

The Schrodinger equation is only solvable for atoms with one electron e.g. H. How do you calculate wavefunction for multi-electron systems?

Answer 39

calculate the wave function of each electron individually and multiply them all together to get the wave function of the atom. In many-electron atoms, energy depends not only on attractive electron-nucleus interactions but also on repulsive electron-electron interactions, which depend on exact positions of all electrons at a given instant, which we cannot specify. The effect of electron-electron repulsion on orbital energies is dealt with by approximate methods.

Question 40

what is the consequence of trying to estimate energies in multi-electron systems?

Answer 40

in a 1 electron atom, energy is only given by n e.g. 2s= 2p, 3s= 3p= 3d. in many-electron atoms, for the same n, the lower the l, the lower the orbital energy so 3s < 3p < 3d.

Question 41

why do atomic orbital energies increase with n?

Answer 41

Electrons are at lower energy (more stable) when they are closer to the positive nucleus. In a multi-electron atom, “inner shell ” electrons are close to nucleus (low n) and “outer shell” electrons the furthest away (highest n).

Question 42

why do atomic orbital energies in many electron atoms increase with l?

Answer 42

shielding + penetration Electron-electron repulsion shields electrons from the full nuclear charge. In general shielding increases with n. BUTshielding also depends on l, eg. an electron in a 2p orbital is more shielded than in a 2s. Electrons experience reduced charge from the nucleus as there are electrons that are closer, shielding the attraction. Reason: Orbital penetration: In 2s orbital there is an area of electron density close to nucleus that does not exist in 2p.

Question 43

If we know relative atomic energies, what does it tell us about the element?

Answer 43

it tells us in what order to fill them and we can predict the electron configuration of elements

Question 44

The difference of energies between subshells (l) can be larger than between principal shells (n). Why do 4s orbitals have a lower energy than 3d orbitals?

Answer 44

4s orbitals have a really close node, according to Aufbau principle

Question 45

what does the ms quantum number mean?

Answer 45

electron spin, either +1/2 or -1/2. Hence why 2 electrons max can be in one orbital, and they must have opposite spins. A way to think about it is the angular momentums having opposite spins. The spin is either +1/2 or -1/2.

Question 46

what does the aufbau principle entail?

Answer 46

(German = building up): To move from one element to the next, add one electron into the orbital of lowest energy which is available. Basically filling up orbitals based on lowest energy first.

Question 47

describe Pauli's exclusion principle

Answer 47

“No two electrons in an atom can have identical values for all 4 quantum numbers” 2 electrons in same orbital cannot have the same spin state, and no more than 2 electrons can be in the same orbital. not all 4 quantum numbers can be the same, 2 electrons in one orbital can’t have the same spin state (ms number), and no more than 3 electrons can be in one orbital. Goes so far as to say no 2 electrons in one atom can have the same quantum numbers.

Question 48

describe Hund's rule

Answer 48

“The most stable state is the arrangement with the maximum number of electrons with identical (“parallel”) spin” Degenerate orbitals are occupied first with single electrons. most stable arrangement= same spin state, as the whole state has lower energy if all the quantum numbers are the same. Basically, orbitals are occupied first with single electrons because more stable (and these electrons should have the same spin state), then they get filled. This is good because there will be less repulsion.

Question 49

why are electrons removed or added from 4s first before 3d, but 4s appears sooner in an electronic configuration?

Answer 49

4s electrons are farther and more loosely held by the nucleus as compared to 3d

Question 50

do periods in the periodic table share the same principle quantum number "n"?

Answer 50

yes, very important for size, ionisation energy and relativistic energy

Question 51

do groups in the periodic table share the same quantum number "n"?

Answer 51

no, however they all have the same valence shell configuration, defining chemical properties and reactions as only outer electrons are important

Question 52

what is electronegativity determined by ?

Answer 52

how much of the element’s nuclear (positive) charge (Z) electrons still feel after all the shielding of other electrons - this is the effective nuclear charge – Z eff. Schrodinger equation explains the electronegativity thing; in a way, the more electronegative atom drags the probability of finding an electron closer to it.

Question 53

does electronegativity increase or decrease across a period?

Answer 53

Across a period from left to right the electronegativity of elements increases as Z eff increases. Across a period, atoms have a greater nuclear positive charge, i.e. Z increases as the number of protons increases. In quantum terms, the value of n is unchanged across a period, so maximum of radial distribution function stays the same, but Z increases across the period, therefore the valence outer (bonding) electrons feel more of Z.

Question 54

does electronegativity increase/ decrease down a periodic group and describe why?

Answer 54

Down a group from top to bottom electronegativity of elements decreases as Z eff decreases. Down a group, atoms have a greater nuclear positive charge, i.e. Z increases as the number of protons increases but: In quantum terms, the value of n increases down a group, therefore the maximum of the radial distribution function moves further away from nucleus; i.e. with increasing n more and more orbitals with electrons closer to nucleus, so the valence outer electrons (used for bonding) experience more and more shielding.

Question 55

what is the relationship between atomic radius and ionisation energy?

Answer 55

when atomic radius decreases across a period, ionisation energy increases for the same reasons but irregularly. when atomic radius increases down a group as n increases, ionisation energy decreases for the same reason, outer valence electrons removed more easily as Z effect decreases due to shielding.

Question 56

why are isotopes of carbon and hydrogen used for NMR instead of the original elements?

Answer 56

carbon-13 and hydrogen-2 are used as they have spins and require less applied magnetic field.

Question 57

why is deuterium used for some reactions instead of H?

Answer 57

the D makes them slower, for example D2O reacts slower than H2O, sometimes 8 times slower. The force holding the molecule together (bond) is the same (electrons) but the mass of atomic nucleus is different. For example, The 13C / 12C isotope ratio can be used to assess origin of material –e.g. photosynthesis in some plants discriminates against 13CO2 in favour of 12CO2 so that organic matter has, on average, 20% lower 13C content compared with atmospheric carbon dioxide.

Question 58

how would 226-Ra change when involved with alpha decay?

Answer 58

ejects an alpha particle (helium nucleus)

Question 59

how would 32-P be affected after beta decay?

Answer 59

the electron comes from decay of neutron to proton and electron

Question 60

what does gamma decay eject?

Answer 60

electromagnetic radiaton

Question 61

why are some isotopes unstable, and how can they become more stable?

Answer 61

Protons do repel but stuck together by neutron glue – The ratio proton to neutron is crucial, so some isotopes decay to more stable ratio (roughly between 1&2 neutron : proton, but lots of complications). e.g. alpha decay changes ratio, and beta decay changes neutron to proton to get towards band of stability. Often an element only has unstable isotopes and created products also unstable leading to complex decay pathways.

Question 62

what are nuclear fission reactions induced by?

Answer 62

neutron bombardments e.g.

Question 63

how do isotopes help in imaging?

Answer 63

Nuclear synthesis of artificial isotopes for imaging SPECT imaging= single-photon emission computed tomography (beta decay) PET imaging= positron emission tomography (beta+ decay)

Question 64

what is a bond?

Answer 64

Something that holds a system (molecule, solid etc.) together – can be ionic or covalent or somewhere in between, or involve other interactions An arrangement of electrons in an orbital that joins 2 or more atoms together (covalent)

Question 65

how are salts held together?

Answer 65

by attractive cation-anion electrostatic interaction, which is NON-DIRECTIONAL- charged spheres packed together to maximise filling of space and cation to anion interaction.

Question 66

what is ionic bonding?

Answer 66

attractive cation-anion electrostatic interaction, that is quite weak individually (hence why salts often dissolve easily in polar solvent as the ions are taken out one-by-one by opposite charges e.g. Na+ ions attracted to O2-, and Cl- attracted to H+) leads to very strong overall lattice energy and high melting point in ionic solids.

Question 67

what is covalent bonding?

Answer 67

a filled orbital shared between 2 or more atoms.

Question 68

what is a molecule? + explain basics of in-phase/ out-of-phase + node stuff

Answer 68

atoms joined by covalent bonds, formed by IN-PHASE overlap of molecular orbitals. When both atoms are in phase= constructive interference= sigma molecular orbital, electron density between the 2 nuclei is high, lowering the energy of the whole system. If not in phase= destructive interference= antibonding molecular orbital (repulsive). There is a node (wavefunction is 0), so no electron density between the nuclei there, so they repel each other.

Question 69

what are the 3 specific features of covalent bonding + draw the formation of a 2p MO and pi bond

Answer 69

1. directional: has a shape, gives molecules their shape 2. strong: to break bond you must move electrons from bonding orbital to another higher energy orbital 3. can be single C-C, double C=C or triple, and lines indicate something about the position of electrons

Question 70

what is the proof of repulsion due to out-of-phase overlap of orbitals?

Answer 70

you can put another electron into an antibonding molecular orbital (sigma star orbital) forming e.g. an H2-, and the bond length will increase (becomes weaker), which is proof of the repulsion.

Question 71

which kind of atoms does an ion-dipole interaction involve?

Answer 71

ion + polar molecule e.g. Na+ and H2O

Question 72

in covalent molecules, which are stronger, the intermolecular bonds or intramolecular bonds, and what does this mean for changes of state?

Answer 72

the intramolecular bonds are strong, intermolecular bonds weak. Therefore, changes of state (solid  liquid gas) occur without breaking the bonds of molecules (unlike ionic).

Question 73

what is a molecular dipole?

Answer 73

Polarised bonds between atoms of different electronegativity gives rise to molecules with permanent dipole moment .

Question 74

draw a diagram explaining how an instantaneous dipole forms, and how dipole-dipole interactions form fixed structures

Answer 74

Question 75

what is lattice energy?

Answer 75

sum of all +/- attractions so overall they’re strong, with high ionisation energy, e.g. NaCl gas could probs kill you and a lot of the energy will come from lattice energy in the gas phrase.

Question 76

describe the van der Waals radius