Ch.7

Question 1

Amplitude

Answer 1

The vertical height of a crest (Or depth of a trough) of a wave; A measure of wave intensity.

Determines the light intensity or brightness— The greater the amplitude, the greater the intensity.

Wave length and amplitude are both related to the quantity of energy carried by a wave.

Waves of a greater amplitude (higher waves) or shorter wavelength (more closely spaced, and the steeper, waves) 

Amplitude and wave length can vary independently of one another. A wave can have a large amplitude or small amplitude and a short wave length or a long wavelength.

The Most energetic wave lengths have large amplitude and short wave lengths. 

Question 2

Wave length

Answer 2

The distance between adjacent crusts (Or analogous points) of a wave. 

Measured in units such as meters, micrometers, or nanometers. 

The wavelength of light determines its colour. 

Shorter wavelength light inherently has greater energy than long wave length light.

Increase wavelength = less energy —> least energy change = longest wavelength = low frequency = is minimum wavelength

🔺E = hc/ wavelength <— longest ^ Wavelength = |
|smallest energy <————————- /
Change =
|
V
Low frequency

Greatest 🔺E = shortest wavelength = high frequency = is max wavelength

Question 3

Wave-particle duality

Answer 3

Certain properties of light are best described by thinking of it as a wave, what other properties are best described by thinking of it as a particle.

Question 4

Frequency

Answer 4

For waves, the number of cycles (or complete wave lengths) that pass through a stationary point in one second

Units are (cycle s^-1) or simply s^-1.
Hertz (Hz) Is an equivalent unit of frequency Defined as 1 cycle s^-1.

The frequency of a wave is directly proportional to the speed at which the wave is travelling— The faster the wave, the more crusts will pass a fixed location per unit time.
Frequency is also inversely proportional to the wavelength— the farther apart the crests, the fewer will pass a fixed location per unit time.

V = c/ “lamb duh”

V is Frequency
c is speed of light
“lamb duh” is wavelength

Question 5

Colours and their wavelengths

Answer 5

Red light, with a range of wave lengths centre at about 750 nm, has the longest wavelength of visible light.

Violet light, with a range of wave lengths centre at about 400 nm, has the shortest.

Nano = 10^-9

Question 6

This?

Answer 6

Low frequency light does not eject electrons because no single photon has the minimum energy necessary to dislodge the electron. Increasing the intensity of low frequency light simply increases the number of low energy photons, but does not produce any single photon with sufficient energy.

In contrast, increasing the frequency of a light, even at a low intensity, increases the energy of each photon, allowing the photons to dislodge electrons with no lag time.

As the frequency of light is increased past the threshold frequency, the excess energy of the photon (beyond what is needed to dislodge the electron) is transferred to the electron in the form of kinetic energy. The kinetic energy of the ejected electron, therefore, is the difference between the energy of the photo and the binding energy of the electron. 

Question 7

Emission spectrum

Answer 7

The range of wavelengths emitted by a particular element; used to identify the element

Question 8

emission spectrum and discrete lines and stationary states and transitions

Answer 8

When an atom absorbs energy, an electron in a lower energy level is excited or promoted to a higher energy level.

In this new configuration, however, the Atom is unstable, and the electron quickly falls back or relaxes to a lower energy level. As it does so, it releases a photon of light containing an amount of energy precisely equal to the energy difference between the two energy levels. 

The transitions between levels that Are farther apart in energy produce light that is shorter in wave length, and therefore higher energy, then between energy levels that are closer together. 
Closer they are longer the wavelength.

Question 9

The energy determines the frequency and wavelength of the photon

Answer 9

The formula we can use to determine this is

E = hc/v

Question 10

Complementary properties

Answer 10

Those properties that exclude one another, I.e., the more you know about one, the less you know about the other. For example, the wave nature and particle nature of electrons are complementary.

When you try to observe which hole the electron goes through (associated with the particle nature of the electron) you lose the interference pattern (associated with the wave nature of the electron).

Position and velocity of the electron are complementary properties.
Position and energy are also complementary properties.

Question 11

Atomic orbitals

Answer 11

From quantum theory, probability distribution maps for the electrons as they exist within Atoms. 

Question 12

Orbital

Answer 12

The probability distribution maps for an electron states in which the electron has well defined energy, but not well defined position. In other words, for each state, we can specify the energy of the electron precisely, but not its location at a given instant. Instead, the electrons position as described in terms of an orbital— a probability distribution map showing where the electron is likely to be found. 

These states are known as energy eigenstates.

Orbitals with a higher values of n have greater (less negative) energies. 
As n increases, the spacing between the energy levels becomes smaller.

Orbitals with the same value of n are said to be in the same principle shell (or level). Orbitals with the same value of n and Sabelle (fancy L) are said to be in the same subshell.

•The number of supplements and any level is equal to n, The principal quantum number. Therefore, n = 1 level has 1 sublevel, n=2 has 2.
•The number of orbitals in any Sublevel is equal to 2L + 1. Therefore, s sublevel (L = 0) has 1 orbital, p sublevel (L =1) has 3, d sublevel (L = 2) had 5. Etc
•The number of orbitals in a level is equal to n^2. n = 1 has 1 orbital, n=2 has 4. n=3 has 9. Etc

Because the magnitude of the overall weight functions falls off (or decreases) more slowly due to the exponential term as n increases, the orbitals increase in size as n increases.

As the energy of an orbital increases(s<p<d<f), the number of non-spiracle nodes also increases. The effect on the shape of the orbital is to increase the number of nodes in the contour diagram. Thus, an S orbital is symmetric, a P orbital is bisected into two symmetrical lobes, And so on with orbitals of higher energy.

Question 13

Quantum numbers

Answer 13

Principal quantum number (n):
An integer that specifies the overall size and energy of an orbital. The higher the quantum number, n, the greater the average distance between the electrons and the nucleus and the higher its energy. The energy is negative because the electrons energy is lowered (made more negative) by it’s interaction with the nucleus.

Angular momentum quantum number (l):
An integer that determines the shape of an orbital.

Magnetic quantum number (m subscript l):
An integer that specifies the orientation of an orbital. 

Spin quantum number ( m sub s): Denotes the electrons spin as either 1/2 (up arrow) or -1/2 (down arrow).

How to identify the four quantum numbers for an element.
1) Find the element on the periodic table.
2) Determine n and L by identifying the period and the block that the element is located in.
3) Determine mL by labelling the block from -1 to +1.
4) Determine ms by determining if the element is in the first or second half of the block.
5) Put all of the information together to list for quantum numbers for an element last valence electron. 

Question 14

Probability density

Answer 14

Probability per unit volume of finding the electron at a point in space as expressed by a three-dimensional plot of the way function squared. 

Hi got density near the nucleus at the centre of the plot indicates a higher probability density for the electron there. As you move away from the nucleus, the probability density increases. 

Question 15

Radical distribution function and nodes (Radical And angular)

Answer 15

Function that represents the total probability of finding an electron within a thin spherical shell at a distance r from the nucleus.

Represents the total probability at a radius r. In contrast to probability density, which has a maximum at the nucleus, the radical distribution function has a value of zero at the nucleus.

A point where the wave function and therefore the probability density and radical distribution function, Will go through zero. The probability of finding the electron add a node is zero.

Radical/Spherical: Spherical region in which there is zero probability of finding an electron. The number of radical nodes is calculated by taking the difference between the principal quantum number, n, and the angular momentum quantum number, L, and then subtracting one from this difference. # of radical nodes = n - L - 1

Angular: a plane or surface where there is zero possibility of finding an electron.

Question 16

Electron configuration

Answer 16

A notation that shows the particular orbitals that are occupied by electrons in an atom.

Electrons generally occupied lowest energy orbitals available. 

To write an electron configuration for an element, first find it’s atomic number from the periodic table this number equals the number of electrons. Then use the order of filling to distribute the electrons in the appropriate orbitals. Remember that each orbital can hold a maximum of two electrons. Consequently,

•The S sublevel has only one orbital and can therefore hold only two electrons
•The P sub level has three orbitals and can hold six electrons.
•The D sub level has five orbitals and can hold 10 electrons.
•The sub level has seven orbitals and can hold 14 electrons. 

Question 17

Ground state

Answer 17

The lowest energy state

Question 18

Electron spin

Answer 18

Fundamental property of electrons; can have a value of + or - 1/2. M subscript s

2 fundamental aspects of electrons spin:
1. Spin, like negative electrical charge, it’s a basic property for all electrons. One electron does not have more or less spend another all electrons had the same amount to spin.
2. The orientation of the electron spin is quantized, with only two possibilities that we can call spin up and spin down.


Question 19

Pauli Exclusion principle

Answer 19

The principal that no 2 electrons in an atom can have the same 4 quantum numbers.

Since two electrons occupying the same orbital have three identical quantum numbers (n, L, m subscript L), they must have different spin quantum numbers (+ or - 1/2). Since there are only two possible spin quantum numbers, the principal implies that each orbital can have a maximum of two electrons, with opposing spins.

Question 20

Degenerate

Answer 20

A term describing two or more electron orbitals with the same value of n that had the same energy.

For example, 3S, 3P, and 3D orbitals (which are all empty for hydrogen in its lowest energy state) all have the same energy, so they are degenerate.

The orbitals within a principle level of a multielectron atom, in contrast, are not degenerate— their energy depends on the value of L. We say that the energies of the sub levels are split. In general, the lower the value of L within a principal level, the lower the energy of the corresponding orbital. Thus, for a given value of n: 
E(s orbital) < E (p) < E (d) < E (f)

Question 21

Coulombs law

Answer 21

The equation that describes the potential energy between two charged particles.

Describes the interactions between charged particles.

1. For like charges, the potential energy (E) is positive and decreases as the particles get farther apart (as r increases.). Since systems tend toward lower potential energy, like charges repel each other (in much the same way that like poles of two magnets repel each other)
2. For opposite charges, the potential energy is negative and becomes more negative as the particles get closer together (as r decreases). Therefore, opposite charges (like opposite poles on a magnet) attract each other.
3. The magnitude of the interaction between charged particles increases as the charges of the particles increase. Consequently, and electron with a charge of 1- is more strongly attracted to a nucleus with a charge of 2+ then it would be to nucleus with a charge of 1+.

Question 22

Ionization energy

Answer 22

Energy required to remove an electron. 

Question 23

Shielding

Answer 23

The effect on an electron of repulsion by electrons in lower energy orbitals that screen it from the full effects of nuclear charge.

Electrons in an orbital closer to the nucleus shield outer electrons more efficiently than if they were in the same orbital. 

Question 24

Penetration

Answer 24

In general, and electron in a 2P orbital has a greater probability of being found closer to the nucleus then an electron in an 2s orbital.

However, due to the bump in the 2S radical distribution function near r=0, which represents the significant probability of finding the 2S electron very close to the nucleus, we say that the electron penetrates the 1S orbital. By penetrating into the 1s orbital, the 2s orbital is not fully shielded by the 1S electrons.

In contrast, the 2P orbital has most of it’s probability in the radical distribution function outside that for the 1s orbital. Electrons in the 2S orbital feel more of the nuclear charge then the electrons in the 2P orbital, resulting in the 2S electrons being lower in energy than the 2P electrons. The 2P electrons are, therefore, easier to remove then the 2S electrons.

• Because of penetration, the sub levels of each principal level or not degenerate for multielectron Atoms.
•In the fourth and fifth principal levels, the effects of the penetration become so important that the 3d and 4s orbitals are very close in energy. The 4d and 5S orbitals are, similarly, very close in energy. 

Question 25

Hund’s rule

Answer 25

The principal stating that for a ground state atom, every degenerate orbital is singly occupied with one electron with the same spin before any orbital is doubly occupied. 

Question 26

Summarizing orbital filling

Answer 26

• Electrons occupied orbitals so as to minimize the energy of the atom; therefore, lower energy or bottles filled before higher energy or bottles. Orbitals fill in the following order: 1s, 2s, 2P, 3s, 3p, (4s or 3d), 4p, (5s or 4d), 5p, etc.

•orbitals can hold no more than two electrons in each. When 2 electrons occupy the same orbital, their spins are opposite. Another way of expressing Pauli exclusion principle (No two electrons in one atom can have the same 4 quantum numbers)

•When orbitals of identical energy are available, electrons first occupied these orbital singly with parallel spins rather than in pairs. Once the orbitals of equal energy are half full, the electrons start to pair. (Hund’s rule) 

Question 27

Core electron configuration

Answer 27

When writing electron configurations for elements beyond neon, or beyond any other noble gas, the electron configuration of the previous noble gas— sometimes called the core electron configuration is often abbreviated by the symbol for the noble gas in square brackets.

Question 28

Electron configuration of cations and anions

Answer 28

For anions, we add the number of electrons indicated by the magnitude of the charge of the anion. For example, the electron configuration for fluorine (F) is 1s^2 2s^2 2p^5 And that of the fluoride ion (F^-) is 1s^2 2s^2 2p^6.

For cations, we subtract the number of electrons indicated by the magnitude of the charge. For example, the electron configuration of lithium (Li) is 1s^2 2s^2 And that of the lithium ion (Li^+) is 1s^2 2s^0. (Or simply 1s^2)

When writing the electron configuration of any ion, including transition metal cations, remove the electrons from the highest energy orbitals first— those with the highest n value.

-Medals tend to lose valence electrons to form cations.
-Nonmetals tend to gain valence electrons to form anions.
-Once d orbitals become full, they become part of the core electrons
-Start from highest orbital and work down level/shell.
-When removing/adding electrons remove/add from highest n value
-

1. Write it for the neutral atom first
   2.Figure out if it’s a metal or nonmetal
2. Figure out group and what charge. 

Question 29

Paramagnetic and diamagnetic

Answer 29

The state of an atom or ion that contains unpaired electrons and is, therefore, attracted by an external magnetic field.

The state of an atom or ion that contains only pair of electrons and is, therefore, slightly repelled by an external magnetic field.



Question 30

Effective nuclear charge

Answer 30

The actual nuclear charge experienced by an electron, defined as the charge of the nucleus plus the charge of the shielding electrons. 

Question 31

Photon and quantum

Answer 31

The smallest possible packet of electromagnetic radiation with an energy equal to hv.

The energy of a photon is inversely proportional to its wave length:

         hc  E = —————
        Wavelength 

-If wavelength value small # —> E # has to be big and Vice versa.

The energy of a photon is directly proportional to its frequency:

E = hv



Question 32

De Broglie relation

Answer 32

The observation that the wave length of a particle is inversely proportional to its momentum.
h
Wavelength = ————
mv
h = Planck’s constant (6.626x10^-34 J•s)
m = mass
v = Velocity

•The faster the electron is moving, the higher its kinetic energy and the shorter it’s wavelength. 

Question 33

Heisenberg’s uncertainty principle

Answer 33

The principal stating that due to wave particle duality, it is fundamentally impossible to precisely determine both the position and velocity of a particle at a given moment in time.

-If we want to know it’s wavelength nature, we can’t know it’s particle nature.
-If we want to know it’s particle nature, we can’t know it’s wavelength nature.

-Can’t know it’s position but can know it’s behavior. Vice versa 

Question 34

Principal level (Principal shell)

Answer 34

The group of orbitals with the same value of n.
For ex:
N = 1

N = 2

N = 3

Question 35

Sublevel (sub shell)

Answer 35

Those orbitals in the same principle level with the same value of n and L. 

-The S sublevel has only one orbital and can therefore hold only two electrons.

-The P sublevel has three orbitals and can hold six electrons.

-The D sublevel has five orbitals and can hold 10 electrons.

-The F sublevel has seven or bottles and can hold 14 electrons. 

Question 36

Quantum number table 

Answer 36

Quantum #: Principal
Symbol: n
Allowed values: 1,2,3,….
Interpretation: governs orbital energy and size

Quantum #: Angular momentum
Symbol: L
Allowed values: 0,…, (n-1)
Interpretation: Governs Orbital shape

Quantum #: Magnetic
Symbol: m subscript L
Allowed values: -1,…,0,…,+1
Interpretation: Governs Orbital orientation

Quantum #: Spin
Symbol: m subscript s
Allowed values: +1/2, -1/2
Interpretation: Governs electron spin.

n L mL orbital Subshell Shell
1 0 0 1s 1s n = 1
2 0 0 2s 2s n = 2
2 1 -1 2px 2p n = 2
2 1 0 2pz 2p n = 2
2 1 +1 2py 2p n = 2
3 0 0 3s 3s n = 3
| 1 -1 3px 3p n = 3
| 1 0 3pz 3p n = 3
| 1 +1 3py 3p n = 3
| 2 -2 3dxy 3p n = 3
| 1 -1 3dxz 3d n = 3
| 1 0. 3dz^2 3d n = 3
| 1 +1 3dyz 3d n = 3
3 2 +2 3dx^2y^2 3d n = 3

Question 37

Radical distribution function

Answer 37

A function that represents the total probability of finding an electron within a thin spherical shell at a distance r from the nucleus.

-Known as orbitals. 

Question 38

Phase

Answer 38

The sign of the amplitude of a wave; can be positive or negative.

-The face of a wave determines how it interacts with another wave. 

Question 39

Audbau principal

Answer 39

The principal that indicates the pattern of orbital filling in an atom.

-Electrons should fully occupied the lowest energy orbitals available before any higher energy orbitals are occupied.
1s < 2s < 2p < 3s < 3p < 4s < 3d < ….

Question 40

Valence electron

Answer 40

The electrons important in chemical bonding. For main group elements, these electrons are in the outermost principal energy level.

Question 41

Key equations and relationships

Answer 41

Relationship between frequency (v), wavelength, and speed of light (c).
(The energy of a photon is inversely proportional to its wavelength)

           c (3.00x10^8 m s^-1) V (Hz) = ————
           Wavelength (nm)
                          c Wavelength = ———
                          v c = (wavelength) (v)

Energy of a photon can also be expressed in terms of wavelength
|
V
The relationship between energy (E), frequency (v), wave length, and Planck’s constant (h).
(The energy of a photon is directly proportional to its frequency)

E = hv
hc hc
E = ——— Also = ———
Wavelength v
hc
Wavelength = ———
E

Energy of an electron in an orbital with a quantum number n In a hydrogen atom.

En = -2.18x10^-18 J (1/n^2) (n = 1,2,3…)

1/wavelength = R (1/nf^2 - 1/ni^2)
H can also be a subscript of R (Rh)

Change in energy that occurs for an electron in an atom when it undergoes a transition between levels n initial and n final. (Z is atomic number)

🔺E = -2-18x10^-18 J (1/nf^2 - 1/ni^2)

If nf < ni, energy diff, 🔺E, is negative b/c atom emits energy as electron relaxes from higher energy level to a lower energy level.
If nf > ni, energy diff, 🔺E, is positive as atom absorbs energy as electron is excited from a lower energy level to a higher energy level.

🔺E = -2-18x10^-18 J(z^2/nf^2 - z^2/ni^2)

De Broglies relation: Relationship between wave length, mass (m), and velocity (v) of a particle.
(wavelength of an electron of mass (m) moving at velocity (v). Velocity of moving electron is related to its wavelength— knowing one is equivalent to knowing the other).

                         h Wavelength = ———
                        mv 1/wavelength = mc (^ De Broglie’s wavelength) 

                          h  Wavelength = ———
                         p
                          h             h  Wavelength = ——— = ———
                          p           mu 

(mass (m) of an object times its velocity (v) is it momentum (p). Therefore wavelength of electron is inversely proportional to its momentum). (u = speed of particle)

Heisenberg’s uncertainty principle: relationship between a particles uncertainty in position (🔺x) and uncertainty in velocity (🔺v).

                                                              h 🔺x x m 🔺v >(& line underneath) ———
                                                             4 pi

Question 42

Atomic emission spectrum and Line spectra

Answer 42

•When atoms of a vaporized element absorb energy, they admit light to give a discontinuous line spectrum.
-Discontinuous: shows they only emit or absorb specific energy

•Elements produce characteristic colors, sometimes indivisible region that are used to identify them through flame tests.

•Spectral Lines are produced when an electron moves from one energy level to another. (As it goes back down it produces light)(Relates to electrons being removed)

•If energy could take continuous values, then the light spectrum would look very similar to the visible spectrum. (They’ll produce all visible light)

•The visible Balmer series is what our naked eyes can see.

-Smallest amount of energy transmission = longest wave, low frequency.
(3 —>2) (Small energy, long wavelengths, low frequency)

-High amount of energy needed = short waves, high frequency.
(5 —> 2) (High energy, short wave length, high frequency)

-If electron comes from higher —> Lower energy level = emission
-If electron comes from lower —> Higher energy level = Absorption
-If electron relaxes back down to initial state, will always produce light. 

Question 43

Photoelectric effect

Answer 43

The idea that energy increases with frequency is consistent with the observed increase in radiated energy with increasing frequency. 

•Relating between energy and wave length.

The study of photoelectric effect produced the following observations:

1. For a lower frequency than a threshold frequency, no electrons are you made it regardless of the light intensity. (If V < Vo —> no e^- ejected.)
2. The threshold frequency depends on the metal being studied.
3. For light with frequency higher than the threshold frequency, the number emitted electrons increases with the intensity of light. (If some thing holding electrons to the surface is 10KJ. You’ll need 10KJ+ to eject.)
4. For light with frequency higher than a threshold frequency, the kinetic energy of the admitted electrons increases linearly with the frequency of light. (10KJ Will just detach it, but will need more to make a move. Excess energy goes to kinetic energy)

These observations were consistent with planck’s postulate of quantized energy.
Quantization is a particle like behaviour.

•Suggested to Einstein that light is made of particles (photons = is a quantum)

•Absorption of one photon of sufficient energy ejects one electron.

-if E > Eo Electrons ejected
-E-Eo = KE electron = 1/2 mv^2
•Higher frequency = higher energy.
(K = Kinetic energy. V = Velocity. E = Energy.)

•Increase intensity = increase current.

•Lower the frequency less likely electrons will be injected.

•Number of electrons that gets ejected increases when the intensity increases.

•Longer wavelengths like red don’t allow electrons to be injected. (This starts around green wavelengths) 

Question 44

Shapes of hydrogen orbitals

Answer 44

(It’s best to look at lecture slides pp. 66)

S Orbitals (L = 0)
1s: 1 circle
2s: 2 Circles and has a radial node.

P orbitals (L = 1)
2px: 2 lobes in x axis & has angular node
2py: 2 lobes in y axis
2pz: 2 lobes in z axis

D orbitals (L = 2)
3dxy: 2 lobes in x axis, 2 lobes in y axis
3dxz: 2 lobes in x axis, 2 lobes in z axis
3dyz: 2 lobes in y axis, 2 lobes in z axis
3dx2y2: no idea how to explain last 2.
3dz2:

-Where + and - cross there’s a node.
-To figure out the notes it is n - 1.
-Node refers to how far away (the radius) the electron is from the nucleus. 

Question 45

Many-electron atoms

Answer 45

•Atoms containing more than one electron, the presence of electron electron repulsions causes the energy of the orbitals to vary in a complicated way.

•Just refers to how much attraction the electron is feeling from the nucleus. If there is one electron, it’ll feel the full electron charge. If there’s more than one electron, the attraction is going to be split between all electrons. Can say that a single electron feels 100% Zeff. Can also say depending on the ion. It feels that Atoms protons, For example Li Has three protons, So we can say the electrons feels 3^+ of its charge. How much Zeff It feels depends on the Z of the atom, I.e., it’s charge.

-Attraction between electrons and nucleus.

-And repulsion between the electrons.

-An electron in an orbital with lower (more negative) energy is more stable.

-In a one electron Atom: orbitals within a shell (same) are degenerate (same energy).

-in a many electron atom: orbitals within a sub shell (same n and L) are degenerate.

-higher in energy, lower in negative energy.

-orbital gets larger as n increases.

•Orbital energy depends on several factors:
-Effect of nuclear charge.

-Effect of electron electron repulsions.

-Effect of orbital shape


Question 46

Many electron atoms: effect of nuclear charger

Answer 46

(Best way to study this is to look at lecture slides pp.71/72)
•Greater nuclear charge lowers orbital energy.
-Lower and energy = more stable = more negative energy.
-Most amount of attraction = more stable. 

Question 47

Many-electron atoms: effect of electron-electron repulsions

Answer 47

(Best way to study this is to look at lecture slides pp.72-73)

•e-e repulsions increases orbital energy
-Electrons in the same shall lead to a slightly lower Zeff.
-Each electron makes the other electron easier to remove. This subshell is less stable. Less stable = more negative energy.
-e-e repulsions are less attractive due to shielding.

•Electrons in outer orbitals (higher n) are shielded from the full nuclear charge, so they have higher energy. 

Question 48

Many-electron atoms: effect of orbital shape

Answer 48

•Orbitals with good penetration (having electron density close to the nucleus) have energy

Example: Compare the radial probability distribution curves for the 2S and 2P orbitals. Note that the 2S orbital is more penetrating then the 2P orbital, and in turn, the 2S orbital is more shielding then the 2P orbital. 

On graph:
•Maximum point tells us highest probability of finding an electron.
• a node is when the graph dips down to x axis.
•sub modal maximum is the smaller curve on 2s orbital graph. Tells us probability of finding a 2s e^- closer to nucleus than on 2P. Meaning 2s able to penetrate more than 2P e^-‘s.

In general for orbitals in the same shell (or same n):
(More shielding, more penetrating)(lower energy) s < p < d < f (Less shielding, less penetrating) (higher energy)

Question 49

Successive ionization energies

Answer 49

Increase as each electron is removed and are especially large when core electrons are removed.

Mg_(g) —> Mg_(g)⁺ + e^- = IE₁

Mg_(g)⁺ —> Mg_(g)²⁺ + e^- = IE₂

Mg_(g)²⁺ —> Mg_(g)³⁺ + e^- = IE₃

Question 50

Electromagnetic (EM) radiation 

Answer 50

•Energy travel through space via electromagnetic radiation, which consists of waves. 
•The name is derived from the fact that it has an electrical field and a magnetic field the travel perpendicular to each other.
-Electrical field is 90° to the magnetic field.

Question 51

Wave nature of light

Answer 51

•The wave properties of EM radiation or described by three parameters:

-Wave length: distance between two consecutive peaks or troughs in a wave. units of m or nm.

-Frequency: number of waves the pass a given point in space in a second. Units of s^- or hertz (Hz)

-Amplitude: the height of a wave crest or depth of a trough.

•In a vacuum, all EM waves travel at the speed of light: (wavelength)(Frequency) = c where c = 2.998x10^8 m/s. Relates to brightness/intensity.

•Reciprocal relationship between wavelength and frequency.
-Longer wave length, shorter frequency
-Shorter wavelength, Higher frequency (b/c more cycles per second)


Question 52

Paramagnetic and diamagnetic

Answer 52

A property related to electron configuration is magnetic behavior, substances can be classified into:

-Can figure out magnetic behaviour by looking at unpaired electrons.

Para:
•Attracted to magnetic field.
•Contains unpaired electrons.

Dia:
•Not attracted to magnetic field.
•All electrons paired