First Midterm Deck

Question 1

Explain the basics of waves, including travelling waves and standing waves?
(6.1)

Answer 1

• All waves have: peak, troughs, wavelength and amplitude
• Wavelength is the distance between successive waves, frequency is the number of waves passing a given point
• standing waves are fixed at both ends (have only integer # of half-wavelengths

Question 2

Describe the wave nature of light? Use appropriate equations to calculate related light-properties such as frequency, wavelength and energy?
(6.1)

Answer 2

• Electromagnetic radiation= self-propagating transverse waves
• refer to back of book for formulas
• two waves added together= constructive interference, two opposite waves added together is destructive interference

Question 3

Distinguish between line and continuous emission spectra?

6.1

Answer 3

Line spectra: light emitted from excited gaseous elements

Continuous spectra: radiation from white light (ex. lightbulb) separated into different components

Question 4

Describe the particle nature of light?

6.2

Answer 4

Einstein proposed this, light could have photons

Photoelectric effect such that photons go in and electrons come out (packets of quantum energy)

Question 5

Describe the Bohr model of the hydrogen atom

6.2

Answer 5

A nucleus of protons and neutrons being circled by an electron in a specific energy shell (n=1)
-a ground state= the lowest energy state

Question 6

Use the Rydberg equation to calculate energies of light emitted or absorbed by hydrogen atoms
(6.2)

Answer 6

Refer to page 12 in handbook.

Must memorize the equation where R=1.097e7 m^-1, nf

Question 7

What are the three postulates of Bohr’s model?

6.2

Answer 7

1) only orbits of specific radii are permitted for electrons in an atom
2) An electron in a permitted orbit has a specific energy
3) Energy is only emitted or absorbed by an electron as it moves from one allowed state to another

Question 8

What is photoemission?

6.2

Answer 8

• energy is quantized so photoemission occurs as electron moves from a higher orbit to a lower orbit (light released in line spectra)

Question 9

Extend the concept of wave-particle duality that was observed in electromagnetic radiation to matter as well (6.3)

Answer 9

Louis de Broglie said that the characteristic wavelength or of any other particle depends on mass and velocity
lambda=h/mv
Matter waves: applicable to all but not realistic to ordinary sized objects.

Question 10

Understand general idea of the quantum mechanical description of electrons in an atom and what defines the distribution of probability to find an electron in a particular part of space
(6.3)

Answer 10

Uncertainty Principle:
Heinsenberg: cannot determine the exact position, direction of motion and speed simultaneously (dual nature of matter limits understanding, the more you know of one the less of others)

Question 11

List traits of the four quantum numbers that form the basis for completely specifying the state of an electron in an atom
(6.3)

Answer 11

• orbital (wavefunctions) describes probability density (electron)
• principle energy number (n)
  -angular quantum number (l)
• magnetic quantum number (ml)
• electron spin (ms)
  ((atomic orbital region is where an electron most likely lives)

Question 12

Describe the traits of the four quantum numbers

6.3

Answer 12

-Principle quantum number (n): specifies electron shell
- Angular quantum number (l): subshell, shape of orbital
l=0=s, 1=p, 2=d, 3=f
-Magnetic quantum number (ml): orientation of orbital -l

Question 13

Derive the predicted ground-state electron configurations

6.4

Answer 13

-electrons fill orbitals in increasing energy s<p></p>

Question 14

Identify and explain two exceptions to the predicted electron configurations for the atoms and ions
(6.4)

Answer 14

Cr: [Ar] 3d^5 4s^1 NOT [Ar] 3d^9 4s^2
Cu: [Ar] 3d^10 4s^1 NOT [Ar] 3d^9 4s^2

• due to special stability of half-filled and filled sub configurations

Question 15

Explain in detail the Aufbau principle

6.4

Answer 15

• lower energy orbitals fill with electrons first
• any orbital can hold up to two electrons
• if 2+ degenerate orbitals are available one electron goes into each until they are filled
• particularly stable configuration is one where set of p or d orbitals is either filled or half filled

Question 16

Relate the electron configurations to element classifications in the periodic table
(6.4)

Answer 16

s,p,d,f 
s-block alkali and alkaline earth metals
p-block main group elements
d- block transition metals
f-block lanthanides and actinides

Question 17

Describe the observed trends in atomic size, ionization energy, and electron affinity of the elements
(6.5)

Answer 17

Increasing Zeff and attraction for elements across the groups to the right,
increasing ionization energy up the periods from bottom to top

Question 18

What is a covalent radius?

6.5

Answer 18

1/2 d between nuclei of 2 identical atoms held by covalent bond (also the bonding radius)

Question 19

What is the effective nuclear charge?

6.5

Answer 19

charge experienced by an electron on a many-electron atom

Question 20

Explain the formation of cations, anions and ionic compounds

7.1

Answer 20

• atoms with low ionization energies lose their valence electron[s] to form cations
• atoms that have a high electron affinity gain electron[s] to form anions

Question 21

Predict the charge of common metallic and nonmetallic elements, and write their electron configurations

Answer 21

• metallic elements low ionization potentials (lose electrons easily)– left in period of near bottom of group
• nonmetallic elements have high electron affinity’s and readily gain electrons– upper-right corner of periodic table

Question 22

Describe the formation of covalent bonds

7.2

Answer 22

valence electrons in one atoms are attracted to the nucleus of another and vice versa

Question 23

Define electronegativity and assess the polarity of covalent bonds
(7.2)

Answer 23

• Electrons shared in homonuclear bonds both have same effective nuclear charge
• heteronuclear bonds electrons are not shared equally
• if electronegativity is similar 0.4 or less a covalent bond is non-polar
• if atoms has x =0.4+ a covalent bond is polar
• atoms x=1.8+ bond is considered ionic

Question 24

Write Lewis symbols for neutral atoms and ions

7.3

Answer 24

-drawing that shows valence electrons as dots around atom

Question 25

Draw Lewis structures depicting the bonding in simple molecules
(7.3) (procedure)

Answer 25

1) sum up valence electrons
2) arrange atoms with most electropositive in center
3) add lone pairs (octet)
4) add remaining electrons to central atom as lone pairs
5) rearrange electrons in lone pairs to make multiple bonds with central atom
6) ensure to account for anions vs. cations
7) molecules that are radicals all atoms except one are octet
8) molecules involving group 13 can be electron deficient
9) elements in row 3+ can be represented in Lewis structures with more than 8 electrons

Question 26

Compute formal charges for atoms in any Lewis structure (formula)
(7.4)

Answer 26

(FC)= #VE - #LPE -#bonds

Question 27

Use formal charges to identify the most reasonable Lewis structure for a given molecule
(7.4)

Answer 27

1) structure where formal charges are zero is preferable
2) if not zero, smallest formal charge is preferred
3) Lewis structure preferable with adjacent formal charges or opposite sign

Question 28

Explain the concept of resonance and draw Lewis structures representing resonance forms for a given molecule
(7.4)

Answer 28

Resonance structures are two forms of a molecule where the chemical connectivity is the same but the electrons are distributed differently around the structure.