Chem Exam #2

Front 1

1. Perform quantitative calculations based on the relationship between wavelength, energy, and the speed of light.

Back 1

1. Perform quantitative calculations based on the relationship between wavelength, energy, and the speed of light.

Front 2

Wavelength

Back 2

λ (Lambda)

Think chinesse frat squinty line

Front 3

Speed of light

Back 3

"c" is speed of light

(constant): c = 2.998 x 10^8 m/s

Front 4

Frequency

Back 4

"ν" is frquency

normally in Hz

Front 5

2. Define wavelength, frequency, and energy of a photon.

Back 5

2. Define wavelength, frequency, and energy of a photon.

Front 6

Wavelength

Back 6

λ (Lambda) is the distance between two peaks of the wave!!!

*typically given in nm*

Front 7

Frequency

Back 7

Frequency is the number of peaks that pass by a given point per second!

Front 8

Energy of a photon

Back 8

Written as: Ephoton=hν (Remember h = Plank's constant = 6.626 x 10^-34 Joules a second ( J s)

Energy of a photo is proportional to the frequency of light (HIGH ENERGY OF A PHOTON = HIGH FREQUENCY)

Front 9

How are wavelength and frequency related?

Back 9

High wavelength = Low frequency

(Inversely proportional)

Front 10

3. Identify, and rank the different types of light radiation.

Back 10

3. Identify, and rank the different types of light radiation.

Front 11

Rank types of light radiation highest wavelength to shortest wavelength.

Back 11

Radio waves

Microwaves

Infrared Radiation

Visible Light

Ultraviolet Radiation

X-Rays

Gamma Rays

Think Rabbits Might Infect Vaginas Under Xtra *tiny* G-strings

Front 12

Rank the VISIBLE LIGHT highest wave length to shortest wavelength!

Between what wavelength do these fall?

Back 12

Red

Orange

Yellow

Green

Blue

Indigo

Violet

ROY G. BIV had a lot of rabbits and hookers wearing xtra tiny g-strings. (700nm to 400nm)

Front 13

4. Describe the photoelectric effect and relate the energy and/or intensity of the photons to the work function and kinetic energy of the ejected electrons.

Back 13

4. Describe the photoelectric effect and relate the energy and/or intensity of the photons to the work function and kinetic energy of the ejected electrons.

Front 14

Describe the photoelectric effect and relate the energy and/or intensity of the photons to the work function and kinetic energy of the ejected electrons.

Back 14

As the light source is increased (made brighter) in intensity the number of electrons ejected increases!

Higher frequency of said light increases the velocity of the ejected electrons!

Front 15

Predicting the maximum velocity

Back 15

Ek=hν−Φ OR Ek=1/2mv^2

- Ek, is the Max kinetic energy of the electron.

- Φ, is the threshold energy called the"work function"

Front 16

5. Understand the relationship between discrete electron energy levels and atomic absorption and emission spectra.

Back 16

More stable = lower energy level.

Electron energy level is not continuous.

Front 17

6. Apply the Rydberg formula to predict the energy of transitions between two n levels in the hydrogen atom.

Back 17

6. Apply the Rydberg formula to predict the energy of transitions between two n levels in the hydrogen atom.

Front 18

Rydberg Formula (Concept)

Back 18

This is the difference in two energy levels, or you can use the other formula to solve for the potential energy value.

Front 19

Rydberg Formula

Back 19

1/λ=R(1/n^2final - 1/n^2initial)

In this equation ΔE(Change in energy)=R(.......)

Where R is Rydberg's constant = 1.097 x 10^7 (1/meters)

Front 20

Rydberg Formula (for potential energy for individual quantum levels)

Back 20

E(subscript)n= -R (1/n)^2

Front 21

7. Understand that quantum mechanics is a mathematical model the solutions of which yield wave functions and energies.

Back 21

7. Understand that quantum mechanics is a mathematical model the solutions of which yield wave functions and energies.

Front 22

8. Understand that the wave function can be used to find a radial distribution function that describes the probability of an electron as a function of distance away from the nucleus.

Back 22

8. Understand that the wave function can be used to find a radial distribution function that describes the probability of an electron as a function of distance away from the nucleus.

Front 23

Wave function

Back 23

Wavefunction = ψ

When the wavefunction is squared, this is directly proportional to the probability of finding the electron at a specific coordinate.

Front 24

9. List, define, and describe the three quantum numbers for the H-atom wave functions and know what possible combinations of quantum numbers are allowed.

Back 24

9. List, define, and describe the three quantum numbers for the H-atom wave functions and know what possible combinations of quantum numbers are allowed.

Front 25

Principal Quantum Number (n)

Back 25

n relates to a specfic wave function

n = the shell

Front 26

Angular Momentum Quantum Number (ℓ)

Back 26

ℓ = the subshell (s,p,d,f,g,h;0,1,2,3,4,5)

ℓ can have any integer starting at 0 going to 'n - 1'

Front 27

Magnetic Quantum Number (mℓ)

Back 27

2ℓ+1 will tell you how many possible quantum numbers there are

OR

mℓ = −ℓ,...,0,...,ℓ

Front 28

10. State the atomic orbital names based on quantum numbers.

Back 28

10. State the atomic orbital names based on quantum numbers.

Front 29

ℓ = 0, ℓ = 1, ℓ = 2

Back 29

s (sphere), p (bone-shape sort of), d (clover shape)

Front 30

11. Describe the difference between one-electron systems and multi-electron systems.

Back 30

11. Describe the difference between one-electron systems and multi-electron systems.

Front 31

The spin (ms)

An orbital can only hold 2 electrons.

Back 31

This is -1/2 or 1/2 doesn't matter which.

Front 32

12. Apply the Aufbau principle to determine the configuration for any atom or ion.

Back 32

12. Apply the Aufbau principle to determine the configuration for any atom or ion.

Front 33

Aufbau Principal

Back 33

1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p

<6s<4f<5d<6p<7s<5f<6d<7p

Front 34

13. Relate the electronic configuration of an atom of an element to its position on the periodic table.

Back 34

13. Relate the electronic configuration of an atom of an element to its position on the periodic table.

Front 35

Relate the electronic configuration of an atom of an element to its position on the periodic table.

Back 35

The further you go on right on the periodic table the electrons in the shell increase.

Front 36

14. Recognize that there are exceptions to the Aufbau principles and where they are likely to occur on the periodic table and why.

Back 36

14. Recognize that there are exceptions to the Aufbau principles and where they are likely to occur on the periodic table and why.

Front 37

Exceptions to aufbau principle

Back 37

When a subshell is half full or one away from being full it will take an electron away from the s orbital to get to half full state Ex: 4s(1)3p(5) for Cr, when you'd think it was 4s(2)3p(4)

Front 38

15. Apply Hund's rule to determine electron configuration using an orbital diagram (electrons in individual orbitals with spins).

Back 38

15. Apply Hund's rule to determine electron configuration using an orbital diagram (electrons in individual orbitals with spins).

Front 39

Hund's Rule

Back 39

Electrons will fill orbitals one by one then start pairing.

Unpaired are called paramagnetic.

paired are diamagnetic.

Front 40

16. Apply the shell model of multi-electron atoms to describe the concept of core vs. valence electrons.

Back 40

16. Apply the shell model of multi-electron atoms to describe the concept of core vs. valence electrons.

Front 41

Electron affinity & Ionization energy

Back 41

Electron affinity: The energy realeased from an atom upon the addition of an electron to form an anion.

BOTH

INCREASE right to left. DECREASE top to bottom

Front 42

When two species have the same configuration (Ex: F(-) and Ne) What do we call this?

Back 42

isoelectronic

Front 43

17. (Answer included don't skip) Define ionization energy.

Back 43

The energy required to remove an electron from an atom.

Ionization of all elements is postive!

INCREASES right to left. DECREASES top to bottom

Front 44

18. Describe the concept of electronic shielding and effective nuclear charge (Zeff) and their relationship to trends in ionization energy, atomic and ionic radii, and electronegativity.

Back 44

18. Describe the concept of electronic shielding and effective nuclear charge (Zeff) and their relationship to trends in ionization energy, atomic and ionic radii, and electronegativity.

Front 45

Zeff (Effective Nuclear Charge)

Back 45

This is simply the atomic number minus the # of core electrons.

(Aka: The number of valence electrons)

Front 46

19. Identify metals and non-metals and predict types of compounds (ionic/covalent) between different elements.

Back 46

19. Identify metals and non-metals and predict types of compounds (ionic/covalent) between different elements.

Front 47

Covalent Bonds

Back 47

Two non-metals

(They share an electron)

Front 48

Ionic Bonds

Back 48

One metal and one non-metal

One takes an electron, and one gains one (anion + cation)

This creates a solid

Front 49

Ionic bonding

Back 49

Ionization (acquiring or losing an electron) > electron affinity (energy released) > crystallization

Front 50

20. Relate coulombs law and lattice energy to ionic radii, ionic charge, and lattice energy.

Back 50

20. Relate coulombs law and lattice energy to ionic radii, ionic charge, and lattice energy.

Front 51

Ionic radii

Back 51

Increase from top to bottom

From left to right it decreases

Front 52

Ionic charge

Back 52

negatively charged is anion (non-metals)

positively charged is cation (metals)

Front 53

Lattice Energy

Back 53

The energy required to seperate an electron

Front 54

Coulomb's Law

Back 54

Closer elemets attract more than further away ones. (Ex. Mg and O more than Mg and Se)

E (is proportional to) (q(1)q(2))/r ϵ

q(1)&q(2) = are the changes on the ions

r = is the distance between them

ϵ (constant) = 1

Front 55

Relate latice energy to atomic radii

Back 55

High latice energy = hard to seperate = Low ionic radii

and vise-versa

The higher up on the periodic table the lower the ionic radii, therefore they're smaller which means they're lattice energy is higher!!! (Because they're more stongly bonded)

Front 56

21. Name basic binary ionic compounds including polyatomic ions.

Back 56

21. Name basic binary ionic compounds including polyatomic ions.

Front 57

22. Describe the distance dependence of the potential energy of a covalent bond.

Back 57

22. Describe the distance dependence of the potential energy of a covalent bond.

Front 58

23. Name basic covalent compounds containing two elements.

Back 58

23. Name basic covalent compounds containing two elements.

Front 59

24. Draw the best Lewis structure (including any the resonance structures) for a molecule or polyatomic ion.

Back 59

24. Draw the best Lewis structure (including any the resonance structures) for a molecule or polyatomic ion.

Front 60

Lewis structure rules

Back 60

1. Hydrogen forms ONE bond!

2. Oxygen atoms don't bond together!(Except O2)

Front 61

25. Apply formal charges to structures and use them to predict the most likely structure.

Back 61

25. Apply formal charges to structures and use them to predict the most likely structure.

Front 62

26. Predict and explain relative bond strength and lengths in a compound using the Lewis structure.

Back 62

26. Predict and explain relative bond strength and lengths in a compound using the Lewis structure.

Front 63

27. Recognize and apply exceptions to the octet rule.

Back 63

27. Recognize and apply exceptions to the octet rule.

Front 64

28. Rank the polarity of covalent bonds based on relative electronegavities.

Back 64

28. Rank the polarity of covalent bonds based on relative electronegavities.

Front 65

nm to meters = ?

Back 65

nm to meter = 1.0 x 10^-9