CHEM Test 4 (5.7-5.9, 7, 8.2-8.12)

12. Which of the following element and ground state electron configuration (either full or noble gas configuration) pairs below are incorrect? Select all that apply. A. Tc: [Ar]5s24d5 B. Mg: [Ne]3s2 C. Li: 1s22s1 D. Mn: 1s22s22p63s23p64s24d5 E. Pb: [Xe]6s24f 145d106p2

AD

7. Which of the following sets of quantum numbers below are valid and may represent any of the three orbitals above? Select all that apply (L=1) A. (6, 1, -1, -½) E. (4, 1, 2, ½) B. (0, 1, 1, ½) F. (5, 2, 2, -½) C. (4, 2, 1, ½) G. (6, 2, 1, ½) D. (5, 1, 0, -½) H. (3, 1, 3, -½)

AD

Consider a set of hypothetical elements in the table below and the number of valence electrons each has. Element Symbol # of valence electrons X 4 Z 5 Which of the hypothetical compounds formed from the elements above would you expect to be free radicals? Select all that apply. A. XZ B. X2Z2 C. XZ2 D. X3Z E. There is not enough information to determine this

AD

Ions: Electron Configurations

• We will now consider ions. Recall that ions are formed when electrons are added or taken away • As a rule, when we take away electrons, we take them away from the valence electrons (not the core electrons) • Furthermore, we take away electrons from the subshell with the highest n value • After we take all of those, we then start taking away electrons with the highest n and l value • Alternatively, when we add electrons, we add to the partially filled orbital having the lowest value of n

The Localized Bonding Electron Model

• We will use the localized electron model to illustrate bonding between molecules • In general, the model states that the electrons in an atom are localized either on one atom (a lone pair) or between atoms(bonding pair) • We will see this model in subsequent sections (Lewis structures, VSEPR) • In the real world, we say electron density is delocalized, but this model is a good foundation to chemical bonding

Electron Spin and the Pauli Principle

• We've already talked about m s, and we established electrons can have two different spins (-½ or +½) • We now need to talk about the Pauli exclusion principle, which states that no two electrons can have the same set of four quantum numbers (n, l, m l, m s) • Finally, we need to establish that since only two values of m s are allowed, orbitals can only hold two electrons

Why is the electron affinity of nitrogen unfavorable, but the electron affinity of phosphorus is favorable?

Adding a new electron to nitrogen requires pairing up two electrons in a 2p orbital (like O) without the benefit of an added proton. The repulsion between these two electrons negates attraction of the new electron to the nucleus. In phosphorus, the same electron configuration applies, but the paired electrons are in a 3p orbital, which has a much larger volume (size) than 2p, which decreases how strongly the two paired electrons repel one another. The energy balance in this case still favors addition of the electron.

Which element in the third period of the periodic table has three valence electrons? Write the chemical symbol (e.g. He).

Al

Rank the compounds given below in order of increasing lattice energy released upon formation (i.e. the least negative, or weakest, associated lattice energy given first). SrSe, MgO, CaS A. SrSe < MgO < CaS B. SrSe < CaS < MgO C. MgO < CaS < SrSe D. MgO < SrSe < CaS E. CaS < MgO < SrSe F. CaS < SrSe < MgO

B

Which of the following statements is true? A. Outer electrons efficiently shield one another from nuclear charge. B. Core electrons effectively shield outer electrons from nuclear charge. C. Valence electrons are the most difficult of all electrons to remove. D. Core electrons have the lowest ionization energies of all electrons. E. Valence electrons in the outermost shell of all elements have the highest ionization energy.

B

According to the theory of wave-particle duality, under what conditions will a particle have the longest wavelength? A. A particle with large mass and high velocity. B. A particle with small mass and high velocity. C. A particle with small mass and low velocity. D. A particle with large mass and low velocity.

C

Which of the following compounds contain BOTH ionic AND covalent bonds? (choose all that apply) a. NaF b. HClO4 c. NaClO4 d. NO2 e. CaCl2 f. LiN3

CF

Which of the following elements has the smallest atomic radius? A. N B. Na C. Al D. F E. Rb

D

The identity of the atom or ion with the electron configuration 1s22s22p63s23p64s13d5 could be: A. chromium in an excited state. B. manganese in an excited state. C. a vanadium cation. D. chromium. E. iron.

D - exception, not excited state

Which of the options provided below represents the first electron affinity of sulfur? A. S+ (g) + e- → S (g) B. S (g) + 2e- → S2- (g) C. S (g) → S- (g) + e D. S (g) → S+ (g) + e E. S (g) + e- → S- (g)

E

Which of the following subshells below may contain an orbital that represent the radial probability distribution function above based on the number of nodes present? Select all that apply A. 2s E. 6s B. 4d F. 5s C. 7f G. 3p D. 2p H. 7d

FH

Explain why a C=N bond is weaker than a C=O bond.

Major reason: Charge separation increases bond strength via ionic component. O is more electonegative than N, so greater ionic character and greater bond strength in the C=O bond. Minor but acceptable reasoning: N has a larger atomic radius than O, so the C=N bond is longer (and therefore weaker) than the C=O bond.

Graham's Law of Effusion:

Rate of effusion for gas 1 / Rate of effusion for gas 2 = Square root M2 / Square root M1 - smaller the molar mass, the faster the rate of effusion

Why is the second ionization energy of magnesium twice as large as the first?

Removing the first electron eliminates an electron-electron repulsion for all of the remaining electrons. The next electron that is removed will require more energy than the first.

Why is the ionization energy of O slightly less than that of N?

The electron that is removed from O begins in a 2p orbital where it is paired with another electron. The repulsion energy of pairing two electrons raises the energy of the O electron relative to the N electron, which is also in 2p but unpaired. Since the electron begins at a higher energy, less energy is required to remove the electron from the O atom.

Why is the electron affinity for group 2A so unfavorable?

The new electron added to Be or Mg must add to the higher energy 2p level (relative to the 2s). Zeff will be smaller, and combined with the repulsions between the new electron and the other electrons, the overall energy balance disfavors adding a new electron to Be or Mg.

Which of the following quantum numbers describes the size and energy of an orbital?

n

Resonance• Formal charge:

the charge assigned to atoms in a Lewis structure assuming bonding electrons are shared equally between bonded atoms • We calculate formal charge by assuming an atom "owns" all its nonbonding electrons, and one-half of its bonding electrons • Formal charge = number of valence electrons - number of nonbonding electrons - ½ number of bonding electrons

Magnetic spin quantum number (m s):

this number provides the orientation of the electron spin • m_s = -½ or +½

Exceptions to the Octet Rule • Some molecules also have incomplete octets,

where molecules form with atoms that do not have a full octet of electrons• Example: BH 3

Periodic Trends (Electron Affinity) Some exceptions:

• Group 2 elements • More positive than expected • Full subshell • Group 15 elements • More positive than expected • Half-filled subshell

Periodic Trends (Ionization Energy) Some exceptions:

• Group 2 elements • Slightly higher than expected • Full s subshell • Group 13 elements • Slightly lower than expected • Electron in p subshell easier to remove • Group 16 elements • Slightly lower than expected• Electron repulsions

Lobes

• The p orbitals (l = 1) have two lobes, like a dumbbell • And then we have the d orbitals (l = 2) • (We also have f, g, etc. orbitals, but we won't get into those)

A real gas will behave most like an ideal gas under which of the following conditions? A. 1 atm and 273 K B. 1 atm and 298 K C. 2 atm and 273 K D. 10 atm and 298 K

B - low pressure, high temp

Which of the following subshell designations are invalid? Select all that apply. A. 1s B. 1p C. 2s D. 2p E. 2d F. 2f

BEF

As mentioned, Einstein also said that that the kinetic energy of the electrons depends on the frequency of light

• As the frequency of light surpasses the threshold frequency, excess energy transfers to the emitted electrons KE = hv - Φ

Partial Ionic Character

• We will consider the ionic character of various bonds • We can also say ionic character increases with ΔEN • Covalent bonds can also have ionic character

Periodic Trends (Electron Affinity)

• Electron affinity is the energy change associated with adding an electron to a gaseous atom X (g) + e - → X- (g) • Electron affinity is typically a negative number because energy is released • Usually becomes more negative across the period • The energy released is a result of attraction between the incoming electron and the nucleus • This goes back to Zeff • The more energy released, the more favored the electron affinity is • In general, we don't see a change down a group • Atomic radii increases, less attraction to the nucleus• Due to radii increase, there are also less electron—electron repulsions

Electronegativity

• Electronegativity: the ability of an atom to attract electrons to itself in a chemical bond • We will use electronegativity to distinguish between different types of bonds (and their polarities, but that will be later) • We can distinguish between ionic and covalent character based on the difference in electronegativities • In general... • Ionic bonding is (typically) seen between a metal and nonmetal • Covalent bonding is (typically) seen between two nonmetals

Polyelectronic Atoms • When we start to look at multi-electron atoms, things change vs the single electron model

• The model we looked at on the previous slide now becomes invalid • Since there are multiple electrons present, electrons closer to the nucleus shield electrons farther away ("shielding") • This causes electrons farther away to be less attracted to the nucleus, providing a higher energy • On the other hand, electrons closer to the nucleus can have a greater attraction to the nucleus, and as a result are at a lower energy ("penetration")

Consider a hypothetical transition metal with the noble gas electron configuration below. What is the most probable charge that will form from this element? A hypothetical noble gas "Ng" is given in the brackets below. Enter your answer as an integer and charge (e.g. +/- 1, 2, 3, etc.). [Ng] 10s19d10

+1

Consider a hypothetical element with the following ionization energies given below. Based on this information, what is the most probable, common oxidation state of the hypothetical element? Answer with the sign and the integer (e.g. +9 or -1). IE1 = 123 kJ/mol IE2 = 344 kJ/mol IE3 = 566 kJ/mol IE4 = 12,300 kJ/mol

+3

• The electromagnetic spectrum is given below, which illustrates all the wavelengths (m) and types of electromagnetic radiation

- Gamma rays 10^-12 - X-rays 10^-10 - Ultraviolet 10^-8 - Visible 4e-7 (purple) to 7e-7 (red) (400-700 nm) - Infrared 10^-4 - Microwaves 10^-2 - Radio waves 10^2 Wavelength and Energy inversely proportional

Based on the following electron configuration, what ion is likely to form? Write your answer with the charge and integer (e.g. +5). [Kr] 5s24d105p5

-1

What is the complete electron configuration of phosphorus? What is the noble gas electron configuration of phosphorus?

1s^22s^22p^63s^23p^3 [Ne]3s^23p^3

The Lewis symbol of an unknown atom has 4 paired electrons and 2 unpaired electrons. How many bonds is this atom likely to form in a covalent compound? Answer with an integer (e.g. 7)

2

How many unpaired electrons does cobalt have in its ground state? Is it paramagnetic or diamagnetic?

3; paramagnetic

For principal quantum level n = 5, determine the number of allowed subshells (different values of l), and give the designation of each.

5 allowed subshells l=0 (5s) l=1 (5p) l=2 (5d) l=3 (5f) l=4 (5g)

Which of the following ions below is diamagnetic? A. V2+ B. Zn2+ C. Fe3+ D. Zr2+

B

Which of the following contains both ionic and covalent bonds? A. MgBr2 B. COS C. BaSO4 D. SF6 E. None of the above contain both ionic and covalent bonds

C

Which of the following orbitals has the most nodes? A. 1s B. 3d C. 4s D. 2p

C

Periodic Trends (Ionization Energy)

• Ionization energy: the energy required to remove an electron from an atom or ion in the gaseous state X (g) → X+ (g) + e - • The energy to remove the first electron is the first ionization energy, the second electron is the second ionization energy, and so on • This is reported as a positive value because energy is required to remove an electron Li (g) → Li + (g) + e - IE1 = 520 kJ/mol Li + (g) → Li 2+ (g) + e - IE2 = 7300 kJ/mol Li 2+ (g) → Li 3+ (g) + e - IE3 = 11,815 kJ/mol • Ionization energy decreases down a group because n increases • Ionization energy increases across a period because Z eff increases

In general, we can also use electron configurations to predict the charges of different ions

- We do this assuming that the ground state noble gas configuration for an element will be most stable • Finally, we can use electron configurations to determine whether an ion (or atom) exhibits magnetic behavior • Paramagnetic: an atom or ion that contains unpaired electrons and that is attracted to an external magnetic field • Diamagnetic: an atom or ion that in which all electrons are paired and that is not attracted to an external magnetic field

How many single bonds, double bonds, triple bonds, and lone pairs are present in the Lewis structure of HCN? Answer with integers (e.g. 7). I. Single bonds: II. Double bonds: III. Triple bonds: IV. Lone pairs:

1, 0, 1, 1

If one of the orbitals given above is the 2pz orbital, how many values of (a) n, (b) ℓ, (c) mℓ, and (d) ms are possible? Answer using an integer (e.g. 11).

1, 1, 1, 2

How many valence electrons are in the elements below? Answer with an integer in the boxes below (e.g. 11). (a) Potassium: (b) Sulfur:

1, 6

How many core electrons does chlorine have? Answer with an integer (e.g. 7).

10

What is the maximum number of electrons that can be found in the 3d subshell? Answer with an integer in the box below (e.g. 11).

10

How many single bonds, double bonds, triple bonds, and lone pairs are present in the Lewis structure of SiH2O? Answer with integers (e.g. 7). I. Single bonds: II. Double bonds: III. Triple bonds: IV. Lone pairs:

2, 1, 0, 2

What is the maximum number of orbitals that can be found in the 2p subshell? Answer with an integer in the box below (e.g. 11).

3

How many unpaired electrons are in the following ions below? Answer with an integer in the boxes below (e.g. 11). (a) Fe2+: (b) Mn3+:

4, 4

Write the Lewis structure for the organic compound CH3OCH3. How many total bonding pairs are in the compound? How many total lone pairs are in the compound? Answer with integers (e.g. 7). I. Bonding pairs: II. Lone pairs:

8, 2

. Which of the following elements has the largest atomic radius? A. As B. O C. Br D. S

A

An illustration of an orbital is provided below. What is/are the possible value(s) for mℓ for the orbital below? Single sphere A. 0 B. -1, 0, +1 C. -2, -1, 0, +1, +2 D. 0, +1, +2 E. -2, 0, +1

A

Consider the following ions: S2- , Cl- , K+, and Ca2+. Which of these ions will have the largest ionic radii? A. S2- B. ClC. K+ D. Ca2+ E. Both K+ and Ca2+ because they have the same n value

A

The effusion rate of carbon monoxide was measured at 70 °C. Afterwards, the effusion rate of a number of gases were collected at the same temperature. Which of the following gases provided below had the closest rate of effusion to carbon monoxide? A. N2 B. Cl2 C. F2 D. Ne E. He F. H2

A

What is the noble gas core electron configuration for bismuth (Bi)? A. [Xe]6s24f145d106p3 B. [Xe]6s25f145d106p3 C. [Xe]6s26f146d106p3 D. [Xe]6s24f145d106p3 E. [Xe]6s24f155d96p3

A

Which of the following elements has the smallest atomic radius? A. Mg B. Ca C. Sr D. Ba E. All of these elements have the same atomic radii because they are all group 2 elements

A

Which of the following options represent the Cd4+ ion in the ground state? A. [Kr]4d8 B. [Kr]5s24d6 C. [Kr]5s25d6 D. [Kr]5d8 E. [Xe]4d8

A

Which of the following Lewis structures would violate the octet rule (in their best structure)? Select all that apply. A. NO B. SO2 C. BCl3 D. CF4 E. ClO3 -

ABCE

In which of the molecules below is the carbon-carbon distance the shortest? A. H2C=CH2 B. HC≡CH C. H3C-CH3 D. H2C=C=CH2 E. H3C-CH2-CH3

B

Which of the following represents an excited state electron configuration of selenium (Se)? Select all that apply. A. [Ar] 4s23d104p4 B. [Ar] 4s23d94p5 C. [Ar] 4s23d104p36s1 D. [Ar] 4s23d104p5 E. [Ar] 4s23d104p3

BC

Many reactions done in the laboratory are air-sensitive and require inert gases such as N2 or Ar to prevent unwanted reactivity. Although only one gas is typically chosen in any given scenario, consider a situation in which a scientist introduces 2.00 mol of N2 and 2.00 mol of Ar into a closed container. If a pinhole leak is introduced at a constant temperature, which of the following statements are true after a period of time? Select all that apply. A. The partial pressure of both gases will increase B. The partial pressure of both gases will decrease C. The partial pressure of both gases will remain the same D. The partial pressure of N2 will be higher than the partial pressure of Ar E. The partial pressure of Ar will be higher than the partial pressure of N2

BE

Which of the following options are valid quantum numbers for the orbital illustrated below? Select all that apply. A. (1, 1, 0, ½) B. (2, 1, -1, -½) C. (4, 2, -1, ½) D. (2, 1, 2, ½) E. (4, 1, -1, 0) F. (2, 1, 1, ½)

BF

If the wavelength of the wave is doubled, what will happen to the speed of light (c)?

C will not change

Which of the following represents an excited state electron configuration of Po? Select all that apply. A. [Xe]6s24f145d106p3 B. [Xe]6s24f145d96p4 C. [Xe]6s24f135d106p5 D. [Xe]6s14f145d106p47p1 E. [Xe]6s24f145d106p5

CD

. Based on periodic trends of electronegativity, which of the following bonds demonstrate the highest ionic character? A. Ca-O B. Mg-F C. Ba-P D. Cs-F E. Al-S

D

Which of the following bonds given below would be the longest? A. O-O B. O-S C. O-Se D. O-Te E. All of the bonds have the same length because they all have single bonds

D

Consider a number of gases provided below. Which of the following will diffuse the fastest at 85 °C? A. SO2 B. NO2 C. CO2 D. Kr E. Xe F. Ar

F

Which of the following options below will have the largest de Broglie wavelength when moving at a constant velocity? lighter or heavier?

Lighter

Rank the atoms N, O, Si, and P in order of increasing electronegativity. Format your answer with a greater than sign (e.g. X < Y < Z).

Si < P < N < O

Why is the ionization energy of B slightly less than that of Be?

The electron that is ionized from B begins in the 2p level, which is higher in energy than the 2s level Be. Since the electron begins at a higher energy, less energy is required to remove the electron from the B atom.

Why is the third ionization energy of magnesium five times as large as the second?

The third electron removed comes from a lower energy orbital that was previously a core orbital (2p vs. 3s). The 2p orbital electrons are nearer the nucleus and experience a larger Zeff (because they are shielded only by 1s and, to a smaller extent, 2s), so are much more tightly held than the 3s electrons that were previously ionized. Removing the third electron will require much more energy

How would the pairwise interaction of ions in KF compare to that in CaO?

Weaker interaction due to smaller charges

Effusion and Diffusion

While we are in this section, we need to keep in mind the velocity of gas molecules and the relation to molar mass • Diffusion: mixing of gases; more specifically, the migration of gas molecules as a result of random molecular motion • Effusion: the escape of gas molecules from their container through a small orifice or pinhole into a chamber under vacuum • The smaller the molar mass, the faster the rate of effusion

Arrange each of the following sets in order from MOST to LEAST electronegative using the periodic trends only. . a. F, Br, Cl b. As, Se, Br c. N, P, Si

a. F > Cl > Br b. Br > Se > As c. N > P > Si

Arrange each of the following sets in order from LARGEST to SMALLEST atomic radius. a. Ge, Pb, Sn b. Sn, Te, Sr c. S2-, Se2-, O2- d. Xe, Cs+, I

a. Pb > Sn > Ge b. Sr > Sn > Te c. Se^2- > S^2 > O^2- d. I^- > Xe > Cs^+

de Broglie

de Broglie came up with the following equation to quantify the relationship between wavelength, mass, and velocity λ = h/mv mass x frequency

What are the possible values of l and ml for a d-orbital?

l = 2 for ANY d-orbital in ANY principal shell ml = -2, -1, 0 , +1, and +2

Principal quantum number (n):

relates the size and energy of the orbital • This is represented as an integer (e.g. 1, 2, 3, 4, etc.). It must be a positive, whole number (not including zero) • As we increase this value, we increase the size and energy of the orbital

Magnetic quantum number (m l):

this number provides the orientation of the orbital where the electron is located • Ranges from -l to +l. It can be negative, positive, or zero• m l = -l...-1, 0, 1...+l

Angular momentum quantum number (l):

this number provides us the shape of the orbital • This is represented as an integer. Must be a positive, whole number (it can be zero). More specifically, this value cannot be greater than n-1 • l = 0, 1, .... n-1 • This number also tells us the subshell where the electron is located • What is the orbital designation for n = 4, l = 1? 4p • What is the orbital designation for n = 2, l = 4? invalid

Periodic Trends (Atomic Radius)

• Atomic radii increases down a group because n increases • Atomic radii decreases across a period because of an increase in effective nuclear charge (Z eff)

Polarizability

• Polarizability: the ease of distorting the electron density in an atom or ion • For atoms, as you increase the size (atomic radii), you increase the polarizability • For ions, we see the same trend. The larger the ion size, the greater the polarizability • All of these trends also tie back to Z eff

Electromagnetic Radiation

- energy that transmits energy through space in a wavelike behavior • There are multiple components to waves we need to look at... • Wavelength (λ): distance between two consecutive peaks in a wave • Amplitude: height or depth of a crest or trough • Frequency (v): number of waves(cycles) per second that pass a given point in space • Speed (c): all forms of electromagnetic radiation move at the speed of light (3.00 x 10 8 m/s); speedof light is constant under vacuum • We have an equation where we can convert between frequency and wavelength c = λv Wavelength and frequency inversely proportional

van der Waals equation

- to understand real gases further (P + (n^2 a) /V2)(V − nb) = nRT • The van der Waals equation builds off of the ideal gas equation, and makes modifications for the behavior of real gases • The van der Waals constants, "a" and "b", can be found in Table 5.3 from your book - Increase of correction factor decreases how ideally gas behaves. - we can determine how ideally a gas behaves • Recall that "a" and "b" are correction factors. So, the smaller the correction factor, the more ideally the gas behaves • More specifically, "a" corrects for intermolecular forces, and "b" corrects for particle size • Furthermore, we can say that "a" corrects for pressure, and "b" corrects for volume

How many core electrons are in the elements below? Answer with an integer in the boxes below (e.g. 11). (a) Nitrogen: (b) Aluminum:

2, 10

Which of the following reactions represents the lattice energy in the formation of solid barium selenide? A. Ba2+ (g) + Se2- (g) → BaSe (s) B. Ba2+ (aq) + Se2- (aq) → BaSe (s) C. Ba2+ (g) + 2 Se- (g) → BaSe2 (s) D. Ba2+ (aq) + 2 Se- (aq) → BaSe2 (s)

A

Based on the previous question, which resonance contributing structure is the most significant contributor to the true resonance hybrid structure of cyanate? Which atom bears the largest proportion of negative charge? Which bond is shortest and strongest?

A is most important contributor (C is negligible because of large formal charge on N and multiple atoms bearing formal charge unnecessarily). O bears greatest negative charge, CN bond is shortest and strongest

Use the image below to answer each of the following questions. A. Single sphere B. Double C. 4 I. Which of these orbitals are valid when n = 2? More than one may apply. II. Which of these orbitals are valid when ℓ = 1? More than one may apply. III. Which of these orbitals has the largest number of possible values for mℓ? IV. If each of these orbitals has the same n, which is highest in energy in an iron atom?

AB B C C

Which of the following are an invalid set of quantum numbers? Select all that apply. A. (0, 0, -1, ½) B. (1, 1, 0, -½) C. (1, 0, 0, -½) D. (2, 1, -1, ½) E. (2, 0, 0, -½) F. (3, 2, -3, ½) G. (3, -3, 2, ½)

ABFG

It is observed that effective nuclear charge (Zeff) increases across a row. Which of the following options best explain this trend? Select all that apply. A. Electrons don't effectively shield other electrons in the same principle energy level B. P electrons more effectively shield s electrons C. The atomic number increases across a row D. Larger atoms have a greater effective nuclear charge E. Electron affinity decreases across a row

AC

Which of the following statements are false? Select all that apply. A. The addition of an electron to nitrogen is more negative (i.e. more favorable) than carbon because nitrogen will be closer to a stable noble gas configuration B. The electron affinity of arsenic is more positive (i.e. least favorable) than germanium because of the subsequent electron repulsions in adding an electron to arsenic's halffilled 4p subshell C. Chlorine's more negative (i.e. more favorable) electron affinity (-347.8 kJ/mol) compared to fluorine's (-327.8 kJ/mol) can be explained by chlorine's larger atomic radius leading to less electron repulsions D. While effective nuclear charge (Zeff) increases across a row, this trend does not correlate to electron affinity because electron-electron repulsions are the only determining factor in the energy released upon electron addition E. All of the above are true

AD

What is true of resonance structures? Select all that apply. A. Different resonance contributing structures differ by placement of electrons only. B. Molecules that exhibit resonance rapidly switch between different resonance contributing structures. C. Different resonance contributing structures differ by placement of electrons and atoms D. The resonance hybrid structure of a molecule is an average of its different resonance contributing structures. E. Resonance hybrid structures contain delocalized electrons.

ADE

. Which of the following statements regarding the Bohr model of the atom are true? Select all that apply. A. Electrons occupy specific orbits that are at fixed distances from the nucleus. B. Electrons become excited by emitting photons. C. Electrons in the ground state may absorb a photon of any wavelength to become excited to a new energy level. D. The emission spectrum of hydrogen contains a continuum of colors. E. Emission spectra have distinct lines because energy levels are quantized

AE

Which statements are true based on the Bohr model of the atom? Select all that apply A. Atomic emission spectra are due to electrons losing energy and changing energy levels B. The energy of transition for n = 3 to n = 4 would be the same as n = 5 to n = 6 C. The first energy level is set at zero energy D. Electrons lose energy as they travel around the nucleus E. Each energy level has a specific energy value

AE

Why is the electron affinity for group 8A so unfavorable?

An added electron will have to add to a new energy level and will be shielded by all of the other electrons (which are now "core" relative to the new electron). Zeff will be essentially zero for the new electron. Combined with the repulsion energies experienced by the other electrons in the atom, it will be unfavorable to add a new electron to a noble gas.

Draw the Lewis structure for the ion CH3COO- . How many total possible octetsatisfied (i.e. best) resonance structures can be drawn for this ion? A. There are no resonance structures for this ion (i.e. there is only one way to draw this Lewis structure) B. Two equivalent resonance structures C. Two non-equivalent resonance structures D. Three equivalent resonance structures E. Three non-equivalent resonance structures

B

Which of the following compounds would have the largest lattice energy released upon formation (i.e. the most positive, or strongest, associated lattice energy). A. NaCl B. MgO C. KBr D. SrBr2

B

Which of the options provided below accurately rank the atoms in order of increasing first ionization energies (i.e. the lowest first ionization energy given first)? A. K < Ca < Ga < Ge B. K < Ga < Ca < Ge C. Ca < K < Ga < Ge D. Ca < K < Ge < Ga E. K < Ga < Ca < Ge F. Ga < K < Ca < Ge

B

Which of the following elements given below would have the most positive (i.e. least favorable) electron affinity? A. K B. Ca C. Ga D. Ge

B - Group 2 exception

Which of the following elements given below would have the most positive (i.e. least favorable) electron affinity? A. Na B. Mg C. P D. Si E. Cl

B - Group 2 exception

Which of the options provided below accurately rank the atoms in order of increasing first ionization energies (i.e. the lowest first ionization energy given first)? A. Na < Mg < Al B. Na < Al < Mg C. Mg < Al < Na D. Al < Na < Mg E. Al < Mg < Na

B - Group 2 exception

Consider an unknown gas that is placed in various temperature and pressure conditions provided below. Under which set of conditions will the unknown gas behave most ideally? A. Pressure: 0.1 atm; temperature: 273 K B. Pressure: 0.1 atm; temperature: 350 K C. Pressure: 1 atm; temperature: 273 K (STP) D. Pressure: 1 atm; temperature: 350 K E. The identity of the gas is required to determine what conditions will affect ideal behavior

B - High temp, low pressure

Answer the following questions using the table below of five different hypothetical gases and their van der Waals correction factors Gas / a ((L2×atm)/mol2) / b (L/mol) Gas A 12.391 0.819 Gas B 13.711 0.901 Gas C 1.341 0.100 Gas D 4.981 0.244 Gas E 7.120 0.450 (a) Which gas will behave the most nonideally? (b) Which gas would have the most similar volume to an ideal gas? (c) Which gas would have the least similar pressure to an ideal gas?

B - largest correction factors C - smallest "b" correction factor B - largest "a" correction factor

Complete the following statements with (A) larger than, (B) smaller than, or (C) equal to. You may need to use an option more than once. Write the corresponding letters in the boxes below. I. The first ionization energy of sodium is ______ the first ionization energy of magnesium. II. The second ionization energy of sodium is ______ the second ionization energy of magnesium.

B and A

Qui-Gon Jinn uses a lightsaber that emits green light (l = 517 nm). His padawan, Obi-Wan Kenobi, uses a lightsaber that emits blue light (l = 462 nm). Both lightsabers emit the same number of photons per second. Based on this information, complete the following statements with (A) smaller than, (B) larger than, or (C) equal to. Only answer with the capital letter of your choice. I. The frequency of the light emitted from Obi-Wan's lightsaber is ______ that of Qui-Gon's lightsaber. II. The energy of the light emitted from Obi-Wan's lightsaber is _________ that of QuiGon's lightsaber. III. Both lightsabers have a wavelength that is ________ that of a lightsaber that emits infrared radiation.

BBA

Which of the following statements are true regarding the van der Waals equation? Select all that apply. A. The van der Waals constant "a" corrects for temperature B. The van der Waals constant "b" corrects for volume C. The smaller the van der Waals constant "a", the more ideally the gas behaves D. The larger the van der Waals constant "b", the more ideally the gas behaves E. The R constant 8.314 J/mol · K is the only constant that may be used in the equation

BC

There can be three equivalent best resonance structures of ________. Select all that apply. A. NO2 - B. NO3 - C. SO3 2- D. SO4 2- E. BrO3 -

BCE

Which of the following statements are true? Select all that apply. A. The quantum number ms represents the size of the orbital B. The quantum number mℓ can be zero C. In a one-electron system, the 4s and 4d orbitals are degenerate D. In a multi-electron system, the quantum numbers (4, 0, -1, -½) and (4, 1, 1, -½) are valid and degenerate E. A 4dxz orbital is degenerate with a 4dxy orbital in a multi-electron system

BCE

Watson and Crick were awarded the Nobel Prize in 1962 for the discovery of DNA, but it was Rosalind Franklin's research group who collected "Photo 51", an x-ray diffraction image that was critical in this discovery. The diffraction image clearly illustrated diffraction spots throughout the pattern (shown below) which was a result of x-rays diffracting off of the DNA structure. Based on this information, which of the following statements is/are true? Select all that apply. A. The spots seen were a result of destructive interference B. The spots seen were a result of constructive interference C. The x-rays used in this experiment had a higher wavelength than infrared or microwave radiation D. The x-rays used in this experiment had a lower wavelength than infrared or microwave radiation E. The x-rays used in this experiment had higher energy than infrared or microwave radiation F. The x-rays used in this experiment had lower energy than infrared or microwave radiation

BDE

Which of the following pairs would have the greatest ionic character? a. BaS or MgS b. BaS or BaO

BaS (larger cation polarizes S less strongly BaO (smaller anion is less polarizable)

Barium is predicted to have a higher polarizability than strontium, calcium, or magnesium. Which of the following options below best supports this? A. Barium's larger atomic mass increases its polarizability B. Barium's smaller atomic mass increases its polarizability C. Barium's larger atomic radius increases its polarizability D. Barium's smaller atomic radius increases its polarizability

C

Based on periodic trends solely, rank the following elements in order of increasing electronegativity (i.e. the least electronegative element given first). A. Si < In < Cl < P B. Si < P < In < Cl C. In < Si < P < Cl D. In < P < Cl < Si E. P < Si < In < Cl F. P < In < Si < Cl

C

Consider the electron configuration for an element written below. Which of the following statements are true? Select all that apply. 1s22s22p63s23p64s13d104p1 A. The electron configuration belongs to the element Ga B. The electron configuration illustrates an element in its ground state C. The electron configuration illustrates an element in its excited state D. The electron configuration belongs to the element Ge in an excited state E. As written, the element has 18 valence electrons

C

Heisenberg's uncertainty principle states that. A. matter and energy are really the same thing B. it is impossible to know anything with certainty C. it is impossible to know both the exact position and momentum of an electron D. there can only be one uncertain digit in a reported number E. it is impossible to know how many electrons there are in an atom

C

It is observed that the electron affinity of sulfur (-200. kJ/mol) is more favorable than the electron affinity of oxygen (-141 kJ/mol). Which of the following reasonably explains this observation? A. Sulfur is more electronegative than oxygen, so it attracts electrons more strongly B. Sulfur has a smaller ionization energy than oxygen, and electron affinity is opposite of ionization energy C. Sulfur has a larger atomic radius than oxygen so the electrons experience less electron-electron repulsion D. Sulfur has a smaller effective nuclear charge than oxygen, so it is better able to shield added electrons

C

Of the following transitions in the Bohr hydrogen atom, the ________ transition results in the emission of the lowest-energy photon. A. n = 1 → n = 6 B. n = 6 → n = 2 C. n = 6 → n = 3 D. n = 3 → n = 6 E. n = 1 → n = 2

C

The electron configuration [Ar]4s13d10 corresponds to which of the following options below? A. Zn B. Zn in an excited state C. Cu D. Cu in an excited state E. Ni F. Ni in an excited state G. None of the above; the electron configuration given is invalid

C

Which of the following has the smallest atomic radius? A. N0 B. N3+ C. N5+ D. N5- E. N3- F. There is not enough information to determine this

C

Which of the options provided below represents the second ionization of Ca? A. Ca+ (aq) + e- → Ca (aq) D. Ca+ (g) + e- → Ca (g) B. Ca+ (aq) → Ca2+ (aq) + e- E. Ca2+ (aq) → Ca+ (aq) + e C. Ca+ (g) → Ca2+ (g) + e- F. Ca2+ (g) → Ca+ (g) + e-

C

Which of the following elements given below would have the most positive (i.e. least favorable) electron affinity? A. Cl B. Si C. P D. S

C - Group 15 exception

Which of the following bonds given below would be the shortest? Which of the following bonds given below would be the longest? Which of the following bonds given below would be the strongest? Which of the following bonds given below would be the weakest? A. X-X B. X=X C. X≡X D. All of the bonds have the same length

C, A, C, A

Two particles with the same number of protons and neutrons have atomic radii of 110 pm (Particle A) and 280 pm (Particle B). Which of the following statements regarding particles A and B is/are true? Select all that apply. A. Particle A is probably a metal, and Particle B is probably a nonmetal. B. Particle A has more valence electrons than particle B. C. The valence electrons in Particle A experience a larger effective nuclear charge than the valence electrons in Particle B. D. Particle B could be an anion of Particle A. E. Particle B could be a cation of Particle A.

CD

Which of the following statements are true? Select all that apply. A. The reactivity of a hydroxide ion (OH- ) is expected to be the same as a hydroxyl radical (•OH) because their chemical formulas are identical B. The Lewis structure of the nitrogen monoxide free radical is expected to have an unpaired electron on the oxygen atom C. The formal charge on nitrogen in the nitrogen dioxide free radical is +1 D. The methyl radical (•CH3) has a total number of 7 valence electrons

CD

Based on periodic trends solely, would you expect nitrogen or sulfur to be more electronegative? A. Nitrogen would be more electronegative because it is farther up in the periodic table B. Sulfur would be more electronegative because it is farther right in the periodic table C. Sulfur would be more electronegative based on its placement in the periodic table and because it contains more protons than nitrogen D. Based on periodic trends, it is inconclusive whether nitrogen or sulfur is more electronegative

D

Consider a hypothetical atom with the following ionization energies given below. Based on this information, how many valence electrons does this atom likely have? IE1 = 312 kJ/mol IE2 = 418 kJ/mol IE3 = 561 kJ/mol IE4 = 691 kJ/mol IE5 = 10,124 kJ/mol IE6 = 49,459 kJ/mol A. 1 valence electron B. 2 valence electrons C. 3 valence electrons D. 4 valence electrons E. 5 valence electrons F. 6 valence electrons (due to a total of 6 different ionization energies present)

D

Hafnium (Hf) can have multiple oxidation states, such as Hf 2+ and Hf 4+. What are the electron configurations of these ions? A. Hf 2+: [Xe] 6s2 4f12 5d2 Hf 4+: [Xe] 6s2 4f10 5d2 B. Hf 2+: [Xe] 6s2 4f14 Hf 4+: [Xe] 4f14 C. Hf 2+: [Xe] 6s2 4f14 5d4 Hf 4+: [Xe] 6s2 4f14 5d6 D. Hf 2+: [Xe] 4f14 5d2 Hf 4+: [Xe] 4f14 E. Hf 2+: [Xe] 6s2 4f14 5d2 6p2 Hf 4+: [Xe] 6s2 4f14 5d2 6p4

D

In a multi-electron system, an orbital in a 4s subshell is at a lower energy than an orbital in a 3d subshell despite a higher n. Which of the following statements best explain this behavior? A. A 3d orbital experiences less shielding than a 4s orbital B. The 3d subshell is able to hold more electrons than a 4s orbital which subsequently results in a lower energy C. There are more nodes in a 3d orbital than a 4s orbital causing an increase in energy D. A 4s orbital experiences more penetration than a 3d orbital E. The number of possible orientations for the 3d orbitals are higher than 4s orbitals resulting in a higher energy

D

Of the bonds C-N, C=N, and C≡N, the C-N bond is ____________. A. strongest/shortest B. strongest/longest C. weakest/shortest D. weakest/longest E. intermediate in both strength and length

D

Rank the compounds given below in order of increasing lattice energy released upon formation (i.e. the least negative, or weakest, associated lattice energy given first). NaI, NaBr, CaO, SrO A. CaO < SrO < NaI < NaBr B. CaO < SrO < NaBr < NaI C. NaI < NaBr < CaO < SrO D. NaI < NaBr < SrO < CaO E. CaO < NaI < SrO < NaBr F. SrO < NaBr < CaO < NaI

D

The atomic radius of main-group elements generally increases down a group. Which of the following options best explains this behavior? A. Effective nuclear charge increases down a group B. Effective nuclear charge decreases down a group C. Effective nuclear charge zigzags down a group D. The principal quantum number of the valence orbitals increases E. Both effective nuclear charge increases down a group and the principal quantum number of the valence orbitals increases

D

Which of the following does not have the element paired with its correct ground state electron configuration? A. As: [Ar]4s23d104p3 B. Re: [Xe]6s24f145d5 C. S: [Ne]3s23p4 D. Y: [Kr]5s25d1

D

Which of the following is the correct electron configuration for a fluoride ion in the ground state? A. 1s22s22p3 B. 1s22s22p4 C. 1s22s22p5 D. 1s22s22p6 E. 1s22s22p63s1

D

Which of the following statements are true for a Ni2+ ion? A. The ion is diamagnetic B. The ion is diamagnetic with one unpaired electron C. The ion is paramagnetic with one unpaired electron D. The ion is paramagnetic with two unpaired electrons E. The ion is paramagnetic with three unpaired electrons F. The ion is paramagnetic with four unpaired electrons

D

Which of the options provided below accurately rank the atoms in order of increasing first ionization energies (i.e. the lowest first ionization energy given first)? A. Br < Se < As B. As < Se < Br C. Br < As < Se D. Se < As < Br E. As < Br < Se

D - Group 16 exception

Consider an experiment in which a low-frequency light is unsuccessful in the emission of electrons from a metal. Based on the principles of the photoelectric effect, which of the following statements below is/are true? Select all that apply. A. Increasing the intensity of the light will eventually cause electron emission B. Decreasing the intensity of the light will eventually cause electron emission C. Increasing the wavelength of the light will eventually cause electron emission D. Decreasing the wavelength of the light will eventually cause electron emission E. Increasing the frequency of the light will eventually cause electron emission F. Decreasing the frequency of the light will eventually cause electron emission

DE

Which of the following are valid, degenerate sets of quantum numbers? Write all answers that meet the criteria in capital letters with no spaces in the answer box (e.g. ABCDE). A. (2, 1, -2, ½) B. (3, 0, -1, -½) C. (3, 0, 0, ½) D. (2, 1, 0, ½) E. (2, 1, -1, -½)

DE

Consider the following ions and atom below. Which of the following statements is/are true? Select all that apply. I - , Cs+, Xe A. All of the atoms and ions have 53 electrons B. All of the atoms and ions are isoelectronic and have the same size C. The I - ion has the smallest atomic size D. The Xe atom has the highest number of protons because it is neutrally charged E. None of the above are true

E

Why does boron have a lower ionization energy than beryllium? A. Beryllium has a full 2s subshell, making it very unstable B. Boron has an electron in an additional sublevel, making the electron harder to remove C. The 2s sublevel is higher in energy than the 2p sublevel D. Boron is to the left of beryllium; ionization energy decreases across a period E. Boron has a single 2p electron, which is easier to remove than one of beryllium's electrons in a full 2s subshell F. The ionization energy of boron is expected to have a higher ionization energy, not lower

E

To quantify each of these levels, we can use the following equation

E = -2.178 x 10 -18 J (Z^2/n^2 ) • We can also look the change of energy when an electron transfers from one level to another level ΔE = -2.178 x 10 -18 J ( 1/(n_f )^2 - 1/(n_i )^2)

Coulomb's Law

which quantifies the energy of interaction between a pair of ions E = (2.31 x 10 -19 J ⋅ nm) ((Q_1Q_2)/r) • You can use this law for like charged or oppositely charged ions • A negative sign indicates an attractive force, and a positive sign indicates a repulsive force

Continuous and Line spectrums

• Continuous spectrum: spectrum containing all of the wavelengths of visible light as a result of white light passing through a prism • Line spectrum: a spectrum showing only discrete wavelengths

Einstein with the Photoelectric effect

• Einstein later concluded via the photoelectric effect that light energy must come in packets; we say these packets are made of photons • He also concluded that the energy of a photon can be described in the equation below where E is energy and h is Planck's constant • h = 6.626 x 10 -34 kg⋅m 2⋅s-1 or h = 6.626 x 10 -34 J⋅s E = hv E = hcλ • This leads us to the dual nature of light, in which we understand that light exhibits both wave and particle like properties

Real Gases

• Equations like the ideal gas equation assume ideal behavior, but real gases do not behave this way • Gases tend to behave ideally at high temperatures and low pressures • Why is this? • In general, it comes down to intermolecular forces • At high temperature, there are less intermolecular forces between gas molecules • At low pressure, particle volume becomes more negligible

Exceptions to the Octet Rule • We will start with odd-electron species, or free radicals

• Free radicals have an unpaired electron in their Lewis structure, making them very reactive • Example: ethyl radical (CH 3CH 2·)

To explain the line spectrum seen with hydrogen, Bohr proposed a model for the hydrogen atom

• He proposed that the electron in a hydrogen atom moves around the nucleus in circular orbits • These orbits are referred to as energy levels, which are the allowed energy states of an atom • The lowest energy level is the ground state • n = 1 for the electron in the hydrogen atom • An electron can go from a higher state to a lower state, or to the ground state. As a result, a photon is emitted, causing energy to be released • Also, an electron can go from the ground state to a higher energy level as a result of absorbing energy, in which we say it is at an excited state

Sizes of Ions

• Ions will differ in size compared to the corresponding neutral atom • In general, we see the following trend with a given element: Anion > neutral atom > cation • The trend with anions, neutral atoms, and cations can be explained using electron configurations • We can also compare the ionic radii of different ions • In general, we see the ionic radii of ions increase as we go down a group • We also need to consider isoelectronic ions (ions with an identical number of electrons) • If you have a set of isoelectronic ions, then look at the number of protons. The higher the number of protons, the smaller the ionic radii - Compare Br-, Kr, and Rb+ Br- > Kr > Rb+

Lattice Energy

• Lattice energy is defined as the change in energy when separated gaseous ions are packed together to form an ionic solid M + (g) + X- (g) → MX (s) • The higher (more negative) the lattice energy, the more energy that's released upon formation • We will look at the general trends of increasing lattice energy including ion size and ion charge • In terms of ion size, an increase in ionic radii results in a decrease in the release of energy (less negative) because the ions cannot get as close to each other • In terms of ion charge, an increase in the magnitude of ionic charge results in an increase in the release of energy (more negative) • Recall Coulomb's Law ... E = (2.31 x 10 -19 J ⋅ nm) ((Q_1Q_2)/r)

Photoelectric effect:

• Light was thought to only behave as a wave, but it was later determined that is also behaved as a particle through the photoelectric effect • Photoelectric effect: when light strikes the surface of certain metals, electrons are ejected • Emission only occurs at certain threshold frequencies (v0) • Once you surpass v0, the number of electrons depends on the intensity • If the v0 is not surpassed, increasing intensity will not have an effect • Furthermore, after passing v0, kinetic energy of electrons depends on the frequency • A common analogy to the photoelectric effect is trying to break a glass window with ping pong balls or baseballs - Breaking window = e- emission - low V0 = 1 ping pong ball - higher intensity (100 ping pong balls) - No glass breaking (no e- emission) either way - High V0 = 1 baseball --> glass breaks (e- emission) - Higher intensity (100 baseballs) --> more glass breaking

Quantum Numbers

• Quantum numbers describe the properties of orbitals in different atoms and molecules • There are four classifications to quantum numbers... • n: principal quantum number • l: orbital angular momentum quantum number • m_l: angular momentum quantum number • m_s: electron spin quantum number • Quantum numbers are written in the following format: (n, l, m l, m s)

We will now switch to Lewis structures for covalent compounds

• Start by writing the general skeletal structure for the molecule • We assign the central atom, which is always the least electronegative atom • All of the other atoms are the terminal atoms (more electronegative atoms) • Hydrogen atoms are always terminal • Calculate the total number of valence electrons • Distribute the valence electrons among the atoms, giving octets to as many atoms as possible • An octet means that an atom is surrounded by 8 electrons • You may also assign pairs of extra electrons as lone pairs (typically to the terminal atoms first) • You may assign double or triple bonds as necessary to give atoms their octets • There are exceptions to the octet rule we will talk about, most importantly is hydrogen and helium • We say that they follow the duet rule • Basically, both atoms may only have two electrons in their valence shells • Your electrons need to be accounted for as bonds or lone pairs(each counts as two electrons) • If you have a polyatomic ion, follow the same rules but put your structure in brackets with the charge

The Aufbau Principle and the Periodic Table • As with most things, there are always exceptions to the rule

• The important exceptions we need to cover include the group 6 elements Cr and Mo, and the group 11 elements Cu, Ag, and Au • In all of these cases, the electron configuration is the same, except one of the s electrons moves up to the d electron subshell Not excited, ground states. Cr : [Ar]4s 1 3d 5 Mo : [Kr]5s 1 4d 5 Cu : [Ar]4s 1 3d 10 Ag : [Kr]5s 1 4d 10 Au : [Xe]6s 1 4f14 5d 10

Hund's Rule

• The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons in a set of degenerate orbitals • Electrons in singly occupied degenerate orbitals must have parallel spins

Resonance • The following rules apply when we are looking at formal charge:

• The sum of all formal charges in a neutral molecule must be zero • The sum of all formal charges in an ion must equal the charge of the ion • Smaller (or zero) formal charges on individual atoms are better than larger ones • When nonzero formal charge cannot be avoided, negative formal charge should reside on the most electronegative atom

When we write electron configurations, we can use a shortcut called the noble gas configuration

• This configuration only shows the valence electrons• These are the electrons in the outermost principal quantum level of an atom • All of the other electrons are classified as core electrons • To write this, we write the core electrons with the nearest noble gas, and then put it in brackets• We then proceed to write the rest of the electron configuration

electrons transition to an excited state

• To be clear, this is different than ions (where electrons have been lost or gained)—the total number of electrons will not change Nitrogen (ground state): [He] 2s^2 2p^3 Nitrogen (excited state): [He] 2s^2 2p^2 3s^1 [He] 2s^2 2p^2 9d^1 as we progress through these subshells, we are doing it in order of increasing energy

Waves can interact with each other in a way called

• Waves can interact with each other in a way called interference • Constructive interference: waves are in phase with each other when they interact • Can lead to the formation of diffraction patterns • When the waves come together, the amplitude increases, and thus the intensity of light increases • Destructive interference: waves are out of phase when they interact

• Resonance occurs when more than one valid Lewis structure can be written for a molecule

• We also need to establish that the atom arrangement is always the same, and only electron density changes (i.e. the electrons are delocalized)• More importantly, we say that the actual electronic structure is not represented by one structure, but rather the average of all the resonance structures • Be sure to realize that the structures don't "alternate" or "switch" from one to another • In many cases, resonance isn't straightforward, and we will find certain structures to be more favorable than others (we determine this with formal charge)

Polyelectronic Atoms

• We must distinguish between the energy in a single electron atom(i.e. hydrogen) and a multi-electron atom • For a single electron atom, subshells with the same principal quantum number (n) are degenerate • That is, they all have the same energy

Exceptions to the Octet Rule • And finally, we have expanded octets

• We only see this in elements of the third row of the periodic table and beyond • These expanded octets can hold more than 8 electrons • Example: SF

• Let's switch to Lewis symbols for ions

• We write these with a few simple steps: • Draw the Lewis symbol of the atom • Either take away electrons (if positive charge) or add electrons (if negative charge) • Put the chemical symbol in brackets and write the charge on the outside • Now we will talk about Lewis structures for both ionic and covalent compounds • We will start with ionic compounds. Recall in an ionic compound there is a complete transfer of electrons • For these compounds, we write the ions using their Lewis symbols like we have covered thus far

Lewis Structures

• We'll start by looking at Lewis symbols for atoms and ions• We do this by counting the number of valence electrons in an atom or ion • Afterwards, we write the symbol, and fill in the electrons around that symbol • As a rule, we start filling in electrons singly first, then we pair them. More specifically, we write the first four electrons unpaired first, then start pairing them

The Aufbau Principle and the Periodic Table

• We've talked about the energy levels of different subshells and orbitals. Now we need to talk about the way electrons fill these orbitals, and the order this occurs in • This brings us to electron configurations • Electron configurations tells us how many electrons are present and the orbitals they fill • The superscripts we use tell us the number of electrons present • Each orbital is filled first before moving on (there are exceptions) • When we write the electron configuration, we are applying the Aufbau principle • Electrons fill subshells of of the lowest available energy first, and then subshells with greater energy

Covalent Bond Energies

• When we talk about different bonding environments, we see a trend with single, double, and triple bonds • With increasing bond order, we see increased strength, but decreased length, and vice versa • What if we have multiple different bonding environments with the same bond order? • In these cases, we compare over all atomic/ionic radii. With increasing atomic/ionic radii, we see increasing bond length

Quantum Mechanical Model of the Atom

• While the Bohr Model was very insightful, it was later determined that for other elements outside of hydrogen it did not work • More specifically, it was determined that electrons do not circle around the nucleus in orbits • Rather, they exist in a wave function that we call orbitals• Heisenberg uncertainty principle • There is a fundamental limitation to just how precisely we can know both the position and velocity of an electron at a given time Δx ⋅ mΔv ≥ h/4π

Lewis Structures for Simple Organic Compounds

• Writing Lewis structures for organic compounds follows the same rules we have covered thus far • In general, you will find carbon is always the central atom

Orbital Shapes and Energies

• You will recall that we use the value l to describe the shape of an orbital, and the letter we use to describe it • We can also count the number of orbitals in a subshell using the equation 2l +1 • Let's start with the s orbitals(l = 0) • These orbitals take on aspherical shape • Nodes: zero probability of finding electron density • The number of nodes indifferent orbitals can be determined using n-l-1

• What is Z eff?

• Zeff is the term for the positive charge felt by electrons • As we go across the row, the number of protons increases, causing Zeff to increase • Electrons increase also, but only core electrons help in shielding, not valence electrons