Use LEFT and RIGHT arrow keys to navigate between flashcards;
Use UP and DOWN arrow keys to flip the card;
H to show hint;
A reads text to speech;
65 Cards in this Set
- Front
- Back
1. Perform quantitative calculations based on the relationship between wavelength, energy, and the speed of light.
|
1. Perform quantitative calculations based on the relationship between wavelength, energy, and the speed of light.
|
|
Wavelength |
λ (Lambda) Think chinesse frat squinty line |
|
Speed of light |
"c" is speed of light (constant): c = 2.998 x 10^8 m/s |
|
Frequency |
"ν" is frquency normally in Hz |
|
2. Define wavelength, frequency, and energy of a photon. |
2. Define wavelength, frequency, and energy of a photon. |
|
Wavelength |
λ (Lambda) is the distance between two peaks of the wave!!! *typically given in nm* |
|
Frequency |
Frequency is the number of peaks that pass by a given point per second! |
|
Energy of a photon |
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) |
|
How are wavelength and frequency related? |
High wavelength = Low frequency (Inversely proportional) |
|
3. Identify, and rank the different types of light radiation. |
3. Identify, and rank the different types of light radiation. |
|
Rank types of light radiation highest wavelength to shortest wavelength. |
Radio waves Microwaves Infrared Radiation Visible Light Ultraviolet Radiation X-Rays Gamma Rays Think Rabbits Might Infect Vaginas Under Xtra *tiny* G-strings |
|
Rank the VISIBLE LIGHT highest wave length to shortest wavelength! Between what wavelength do these fall? |
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) |
|
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. |
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. |
|
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. |
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!
|
|
Predicting the maximum velocity |
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" |
|
5. Understand the relationship between discrete electron energy levels and atomic absorption and emission spectra. |
More stable = lower energy level. Electron energy level is not continuous. |
|
6. Apply the Rydberg formula to predict the energy of transitions between two n levels in the hydrogen atom. |
6. Apply the Rydberg formula to predict the energy of transitions between two n levels in the hydrogen atom. |
|
Rydberg Formula (Concept) |
This is the difference in two energy levels, or you can use the other formula to solve for the potential energy value. |
|
Rydberg Formula |
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) |
|
Rydberg Formula (for potential energy for individual quantum levels) |
E(subscript)n= -R (1/n)^2 |
|
7. Understand that quantum mechanics is a mathematical model the solutions of which yield wave functions and energies. |
7. Understand that quantum mechanics is a mathematical model the solutions of which yield wave functions and energies. |
|
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. |
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. |
|
Wave function |
Wavefunction = ψ When the wavefunction is squared, this is directly proportional to the probability of finding the electron at a specific coordinate. |
|
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. |
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. |
|
Principal Quantum Number (n) |
n relates to a specfic wave function n = the shell
|
|
Angular Momentum Quantum Number (ℓ) |
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' |
|
Magnetic Quantum Number (mℓ) |
2ℓ+1 will tell you how many possible quantum numbers there are OR mℓ = −ℓ,...,0,...,ℓ |
|
10. State the atomic orbital names based on quantum numbers. |
10. State the atomic orbital names based on quantum numbers. |
|
ℓ = 0, ℓ = 1, ℓ = 2 |
s (sphere), p (bone-shape sort of), d (clover shape) |
|
11. Describe the difference between one-electron systems and multi-electron systems. |
11. Describe the difference between one-electron systems and multi-electron systems. |
|
The spin (ms) An orbital can only hold 2 electrons. |
This is -1/2 or 1/2 doesn't matter which. |
|
12. Apply the Aufbau principle to determine the configuration for any atom or ion. |
12. Apply the Aufbau principle to determine the configuration for any atom or ion. |
|
Aufbau Principal |
1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p
<6s<4f<5d<6p<7s<5f<6d<7p |
|
13. Relate the electronic configuration of an atom of an element to its position on the periodic table. |
13. Relate the electronic configuration of an atom of an element to its position on the periodic table. |
|
Relate the electronic configuration of an atom of an element to its position on the periodic table. |
The further you go on right on the periodic table the electrons in the shell increase. |
|
14. Recognize that there are exceptions to the Aufbau principles and where they are likely to occur on the periodic table and why. |
14. Recognize that there are exceptions to the Aufbau principles and where they are likely to occur on the periodic table and why. |
|
Exceptions to aufbau principle |
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) |
|
15. Apply Hund's rule to determine electron configuration using an orbital diagram (electrons in individual orbitals with spins). |
15. Apply Hund's rule to determine electron configuration using an orbital diagram (electrons in individual orbitals with spins). |
|
Hund's Rule |
Electrons will fill orbitals one by one then start pairing. Unpaired are called paramagnetic. paired are diamagnetic. |
|
16. Apply the shell model of multi-electron atoms to describe the concept of core vs. valence electrons. |
16. Apply the shell model of multi-electron atoms to describe the concept of core vs. valence electrons. |
|
Electron affinity & Ionization energy |
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 |
|
When two species have the same configuration (Ex: F(-) and Ne) What do we call this? |
isoelectronic |
|
17. (Answer included don't skip) Define ionization energy. |
The energy required to remove an electron from an atom. Ionization of all elements is postive! INCREASES right to left. DECREASES top to bottom |
|
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. |
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. |
|
Zeff (Effective Nuclear Charge) |
This is simply the atomic number minus the # of core electrons. (Aka: The number of valence electrons) |
|
19. Identify metals and non-metals and predict types of compounds (ionic/covalent) between different elements. |
19. Identify metals and non-metals and predict types of compounds (ionic/covalent) between different elements. |
|
Covalent Bonds |
Two non-metals (They share an electron) |
|
Ionic Bonds |
One metal and one non-metal One takes an electron, and one gains one (anion + cation) This creates a solid |
|
Ionic bonding |
Ionization (acquiring or losing an electron) > electron affinity (energy released) > crystallization |
|
20. Relate coulombs law and lattice energy to ionic radii, ionic charge, and lattice energy. |
20. Relate coulombs law and lattice energy to ionic radii, ionic charge, and lattice energy. |
|
Ionic radii |
Increase from top to bottom From left to right it decreases
|
|
Ionic charge |
negatively charged is anion (non-metals) positively charged is cation (metals) |
|
Lattice Energy |
The energy required to seperate an electron |
|
Coulomb's Law |
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 |
|
Relate latice energy to atomic radii |
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) |
|
21. Name basic binary ionic compounds including polyatomic ions. |
21. Name basic binary ionic compounds including polyatomic ions. |
|
22. Describe the distance dependence of the potential energy of a covalent bond. |
22. Describe the distance dependence of the potential energy of a covalent bond. |
|
23. Name basic covalent compounds containing two elements. |
23. Name basic covalent compounds containing two elements. |
|
24. Draw the best Lewis structure (including any the resonance structures) for a molecule or polyatomic ion. |
24. Draw the best Lewis structure (including any the resonance structures) for a molecule or polyatomic ion. |
|
Lewis structure rules |
1. Hydrogen forms ONE bond! 2. Oxygen atoms don't bond together!(Except O2)
|
|
25. Apply formal charges to structures and use them to predict the most likely structure. |
25. Apply formal charges to structures and use them to predict the most likely structure. |
|
26. Predict and explain relative bond strength and lengths in a compound using the Lewis structure. |
26. Predict and explain relative bond strength and lengths in a compound using the Lewis structure. |
|
27. Recognize and apply exceptions to the octet rule. |
27. Recognize and apply exceptions to the octet rule. |
|
28. Rank the polarity of covalent bonds based on relative electronegavities. |
28. Rank the polarity of covalent bonds based on relative electronegavities. |
|
nm to meters = ? |
nm to meter = 1.0 x 10^-9 |