Qualifier Electron Configuration and the Periodic Table______

Name

Part 1: Vocabulary (1 pt each). Choose the correct letter from the choices on the right column which best match the definitions below.

Electron Configuration and the Periodic Table______

Name

Electron Configuration and the Periodic Table______

Name

  1. An electronic configuration that violates the Aufbau Principle
  1. The amount of energy required to remove an electron from an atom in the gaseous phase
  1. The rule that stipulates electrons fill degenerate orbitals individually before pairing up
  1. A structure that shows the valence electrons for a given element
  1. The radius of an atom when it either loses or gains electrons

A)Quantum

B)Electromagnetic Spectrum

C)Photon

D)Atomic Emission

E)Ground State

F)Excited State

G)Orbital

H)Lewis Dot Structure

I)Atomic Radius

J)Ionic Radius

K)Ionization Energy

L)Electronegativity

M)Aufbau Principle

N)Hund’s Rule

Electron Configuration and the Periodic Table______

Name

Electron Configuration and the Periodic Table______

Name

Part II: Multiple Choice (1 pt each). Each question must show either a mathematical calculation or process of elimination explaining why it is not the correct answer for credit.

Electron Configuration and the Periodic Table______

Name

  1. All the elements in Period 3 have a total of 2 electrons in the

All atoms in Period 3 have a completely filled 1st and 2nd energy level, but they do not necessarily have a completely filled 3 energy level.

  1. 2s sublevel – holds 2 electrons – full for each element
  2. 2p sublevel – holds 6 electrons –full for each element
  3. 3s sublevel – holds 2 electrons – NOT full for each element
  4. 3p sublevel – holds 6 electrons – NOT full for each element
  1. Which two characteristics are associated with metals?

Generally, metals have luster, malleability, ductility, and are good conductors. Metallic atoms are generally the largest in the period, have the lowest electronegativity and ionization energy.

  1. Low ionization energy and low electronegativity
  2. Low ionization energy and high electronegativity
  3. High ionization energy and low electronegativity
  4. High ionization energy and high electronegativity
  1. The elements in the periodic table are arranged byatomic number/number of protons
  2. Atomic mass
  3. Atomic number
  4. Number of electrons
  5. Number of neutrons
  1. Which atom has a partially filled d sublevel?Write the orbrital notation for any element in the d sublevel
  2. Mn c. Clnot in the d sublevel
  3. Fnot in the d subleveld. Zn

Mn: [Ar] ______

4s 3d

Zn: [Ar] ______

4s 3d

Manganese is the only element here with any unpaired electrons in the d subshell

  1. Elements that readily gain electrons tend to have

Atoms that gain electrons are nonmetals. Nonmetals are brittle, dull, substances that do not conduct electricity or heat. Nonmetal atoms are generally the smallest in the period, have the highest electronegativity and ionization energy.

  1. High ionization energy and high electronegativity
  2. High ionization energy and low electronegativity
  3. Low ionization energy and high electronegativity
  4. Low ionization energy and low electronegativity
  1. How do the atomic radius and metallic properties of sodium compare to the atomic radius and metallic properties of phosphorus? Sodium and phosphorus are in the same period on the periodic table which means that the radius of sodium is larger than that of phosphorus. Since sodium is larger it will naturally lose electrons more readily making it a more metallic atom.
  2. Sodium has a larger atomic radius and is more metallic
  3. Sodium has a larger atomic radius and is less metallic
  4. Sodium has a smaller atomic radius and is more metallic
  5. Sodium has a smaller atomic radius and is less metallic
  1. Quantum numbers [n, l, m, s], are the solutions to the Schrödinger Wave Equation. The angular momentum quantum number, l,

Also called the azimuthal quantum number, l, describes the shape of the orbital.

  1. Describes the principle level and the size of the orbital, n
  2. Describes the sub-level and the size of the orbital, not a quantum number that does both of these
  3. Describes the principle level and the shape of the orbital, n
  4. Describes the sub-level and the shape of the orbital
  1. A diatomic element with a high ionization energy would most likely be

Diatomic elements = 7,7,7 , nonmetals have high electronegativity, small radius

  1. a nonmetal with a high electronegativity
  2. a nonmetal with a low electronegativity
  3. a metal with a high electronegativity
  4. a metal with a low electronegativity
  1. Which grouping of circles, when considered in order from the top to the bottom, best represents the relative size of the atoms Li, Na, K, and Rb, respectively?

Radius of atoms increases as you go down a group, so Li < Na < K <Rb

  1. As the elements across Period 2 are considered the atomic radius is

Across a period – atomic radius decreases, ionization energy increases, electronegativity increases, nuclear charge is greater and more effective

  1. larger and the nuclear charge is greater
  2. larger and the nuclear charge is smaller
  3. smaller and the nuclear charge is greater
  4. smaller and the nuclear charge is smaller

Electron Configuration and the Periodic Table______

Name

Part III: Completion

Using the information given in the table below, complete the open blocks [6 pts]

Electron Configuration and the Periodic Table______

Name

Electron Transition / Wavelength Emitted / Radiation Type / Color
n = 3 to n = 1 /
  1. 103 nm
/
  1. UV
/ None
  1. n = 6 to n = 2
/ 410 nm /
  1. Visible
/ violet
n = 4 to n = 2 /
  1. 486 nm
/ visible /
  1. blue

  1. How many unpaired electrons in each of the following elements? Write the orbital notation correctly labeling the orbitals containing the electrons. Use lines or boxes for each orbital and arrows for the electrons.

[2 each]

Electron Configuration and the Periodic Table______

Name

  1. N 3

Orbital Notation:

[He] ______

2s 2p

  1. K1

Orbital Notation:

[Ar] ____

4s

Electron Configuration and the Periodic Table______

Name

  1. In the space below, write the following notations for Fe.
  2. Write the noble gas configuration.[2]

[Ar]4s23d6

  1. Write the orbital notation for the electrons past the noble gas. Include the correct subshell in the boxes and use arrows to represent electrons. [2]

[Ar ]______

Electron Configuration and the Periodic Table______

Name

Part IV: Free Response Questions:All questions must be answered in complete sentences to receive credit.

  1. Explain, in terms of electrons, why the radius of a nitride ion is larger than the radius of a nitrogen atom. Show the electron configuration of each species. [2]

Electron Configuration and the Periodic Table______

Name

N 1s22s22p3, 7 protons, 7 electrons

N-31s22s22p6, 7 protons, 10 electrons

Nitride is larger than nitrogen because it has 3 extra electrons that repel each other in the outer shell.

Electron Configuration and the Periodic Table______

Name

  1. A sample of a potassium-containing compound is heated in the flame of a Bunsen burner until the atoms are in the excited state. The color of the flame appears pink.
  2. Explain the pink color of the flame in terms of both subatomic particles and energy states. [1]

Atoms of potassium become excited when they are heated in a flame. The electrons of potassium absorb energy and move to a higher energy level (the excited form) and when they release the energy the electron falls back to the n=2 energy level, where visible light is produced. A photon is released as the electron relaxes with a wavelength that corresponds to the visible spectrum, specifically pink light.

  1. In the space below, answer the following questions for arsenic
  2. Write an electron configuration for an atom of arsenic in an excited state. [1]

Ground state: 1s22s22p63s23p64s23d104p3

Excited state: 1s22s22p63s23p64s23d104p26s1

Note: there is more than one possible configuration for the excited state, the configuration can be different, but the quantity of electrons must be the same

  1. Explain, in terms of atomic structure and electron configuration, how arsenic reacts. [1]

Arsenic reacts by gaining three valence electrons to completely its 4p orbitals and achieves the noble gas configuration of krypton.

  1. The latest element to have a claim of discovery, with six atoms having been detected by a joint Russia-US collaboration at Dubna, Moscow Olbast, Russia, in 2009-2010, is Unuseptium. The temporary name for this element is assigned until the discovery is acknowledged by the IUPAC. Ununseptiums symbol is Uus. Scientist created this element by fusing Calcium-48 and Berkelium-249. The new element has the following electron configuration 2-8-18-32-32-18-7.
  2. State the group and period this element is in. [1]

Element #117 – 117 electrons & 117 protons

Group 17, Period 7

  1. Draw the Lewis electron-dot diagram for the atom. [1]

Uus – 7 valence electrons = 7 dots

Ask for this in class

  1. List two other elements with similar properties as this element. [2]

F, Cl, Br, I, At

Electron Configuration and the Periodic Table______

Name

  1. State the trend in first ionization energy for the elements in the table as the atomicnumber increases. [1]

The first ionization energies in the table decrease as the atomic number increases.

  1. Explain, in terms of atomic structure, why this trend occurs. [1]

This trend occurs because orbital size increases as you go down a group.

  1. Explain, in terms of atomic structure, why cesium is more reactive than lithium. [1]

Electron Configuration and the Periodic Table______

Name

Cesium is more reactive than lithium because it has lower ionization energy. Ionization Energy is a measure of metal reactivity. Less energy is required to remove an electron from cesium than lithium.

Electron Configuration and the Periodic Table______

Name

Honors Only

Electron Configuration and the Periodic Table______

Name

  1. Bohr’s model of the atom was able to accurately explain:the energy for spectral lines for each atom
  2. Why spectral lines appear when atoms are heated.
  3. The energies of the spectral lines for each element.Best answer of the two
  4. Why electrons travel in circular orbits around the nucleus.False, this is not true
  5. all of the above answers is correct.If c is false, d is false
  1. Which of the following is true of the distance of an electron from the nucleus of a 1H atom?Heisenberg’s Uncertainty Principle states that you can either know the location or the energy of an electron, we know the energy.
  2. It is 1 amu.
  3. It remains constant over time.
  4. Its distance at any given time can only be predicted by looking at a“wave function”.
  5. It is impossible to say where an electron will be at any given time.
  1. The magnetic quantum number of an orbital defines:

m, defines the orientation of the orbital

  1. The energy level of the orbitaln
  2. The shape of the orbitall
  3. The spatial orientation of the orbital
  4. The spin of the electrons in the orbitals
  1. For which of the following atoms does the electron in the outermost occupied orbital have thequantum numbers n = 3 and l = 1?3p
  2. N 2p
  3. S3p
  4. Mg 3s
  5. Cs 6s
  6. Ge4p
  1. Which of the following atoms has a half-filled p-orbital in its ground state?
  2. B p1
  3. C p2
  4. Pp3
  5. Sp4
  6. Tep4
  1. Which of the following orbitals cannot exist?

Calculate the l values for each n value below

  1. 3dfor n = 3, l = 2,1,0 (s, p, d)
  2. 4f for n = 4, l = 3,2,1,0 (s, p, d, f)
  3. 2sfor n = 2, l = 0 (s)
  4. 5d for n = 5, l = 4,3,2,1,0 (s, p, d, f, g)
  5. 2dfor n = 2, l = 1,0 (s, p)
  1. One of the outermost electrons in aArgon atom in the ground state can be described by which of the following sets of four quantum numbers?

3p6, n = 3, l = 1

  1. 3, 0, 0, -½
  2. 3, 1, -1, -½
  3. 2, 1, 0, ½
  4. 3, 0, 1, -½
  5. 2, 1, -1, ½
  1. Which of the following sets of quantum numbers (n, l, ml, ms) are disallowed for an electron inan atom?

n values must be whole numbers, l must be n-1, m values are –l through +l, and spin values are either – ½ or + ½

  1. 4, 2, 0, ½
  2. 3, 3, -3, -½Three answers – oops!
  3. 2, 0, +1, ½
  4. 4, 3, 0, ½
  5. 3, 2, -2, -1
  1. In which of the following groups are all the species isoelectronic?Same number of electrons
  2. Mg+, Na, Ne11, 11, 10
  3. Br-, Rb+, Sr2+36, 36, 36
  4. K, Ca+, Cl-19, 19, 18
  5. Be2+, He, Li+2+2, 2, 1
  6. Ba2+, Rn, I-54, 86, 54
  1. One of the outermost electrons in a strontium atomin the ground state can be described by which of the following sets of four quantum numbers? (2 points)

Sr has 2 valence electrons in 5s2

  1. 5, 2, 0, ½5d3
  2. 5, 1, 1, ½5p3
  3. 5, 1, 0, ½5p2
  4. 5, 0, 1, ½Doesn’t exist
  5. 5, 0, 0, ½5s1

Electron Configuration and the Periodic Table______

Name

For each of the following questions show all work where necessary and use complete sentences in your explanations.

  1. Give the four quantum numbers which describe the circled electrons (n,l,ml,ms) [2 each]
  2. K ↑↑n = 4, l = 0, m = 0, s = + ½

4s 4p

  1. U↑↓↑ ↑ ↑ ↑ n = 5, l = 3, m = -2, s = + ½

7s 5f

  1. A certain violet light has a wavelength of 413 nm. What is the frequency of the light? What is the energy of the light? [2]

c=ν ν=c/c = 3.0 x 108 m/s = 7.26 x 1014 s-1

c= 3.0x108 m/s 413 x 10-9m

= 413 nm = 413x10-9m

E=hνE = (6.626 x 10-34Js)(7.2 x 1014 s1) = 4.81 x 10-19 J

E=?

h= 6.626x10-34 Js

ν= 7.26 x 1014 s-1

  1. If an electron transitions from n=5 to n=2, how much energy does the photon released carry? [2]

DUPLICATE FROM PROBLEMS WORKED IN CLASS – OOPS!!

According to the Bohr Model of Hydrogren a transition from n=5 to n=2 will release a photon with a wavelength of 434nm. This can be found using your reference pages and can be used to calculate the frequency or energy of a photon.

E=hc/E = (6.626 x 10-34Js)(3.0 x 108 m/s) = 4.58 x 10-19 J

E=? 434 x 10-9 m

h= 6.626x10-34 Js

c= 3.0x108 m/s

= 434 nm = 434x10-9m

  1. Using orbital notation and stability arguments explain why tin and lead form both +2 and +4 ions. [2]

Both tin and lead have the valence configuration s2p2 and have a relatively unstable configuration with a full s subshell and a partially filled p subshell. When the two atoms form the +2 ion they lose the 2 p electrons leaving the ion with a more stable configuration (full s subshell) than the atom. The atoms then lose the remaining 2 s valence electrons when they form the +4 ion and have a full energy level. This configuration is the most stable.

  1. Explain the why the anomalies between group 15 and group 16 become less pronounced as you go down the periodic table. [2]

The difference between the orbitals is less pronounced as you increase the atomic radius of the atoms in group 15 and 16 so the energy difference between the two configurations decreases.

  1. Of the four trends which trend is best to predict metallic and non-metallic reactivity? Why? [2]

Ionization energy predicts metallic reactivity because it describes the loss of electrons.

Electron affinity predicts non-metallic reactivity because it describes the gaining of electrons.

  1. Silver (Ag) is an exception in nomenclature because as a transition metal, it only forms one cation, Ag+1. Using orbital notation and stability arguments, explain why silver doesn’t want to form Ag+2 as well. [2]

Silver does not want to form Ag+2 because its configuration is [Kr] 5s14d10, when it loses the first electron it has a full energy level.