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Electrons and Atomic Properties Practice Test

1.  The electron configuration for an atom is 1s22s22p63s23p64s23d104p6. How many valence electrons does this atom have?

(A) 18 (B) 8 (C) 6 (D) 36

Count the electrons in the highest principal energy level, which in this case is level 4. There are 8 in this level, so there are 8 valence electrons.

2.  The electron configuration for helium is 1s2. The “1” stands for….

(A) the orbital

(B) how many electrons helium has

(C) the energy level the electrons are in

(D) the shape of the area that the electron’s in 90% of the time

1 is the energy level, s is the subshell, 2 tells you how many electrons are in that subshell.

3.  Whenever an electron in hydrogen electron promoted from a low energy level to a high energy level, it:

(A) Absorbs energy

(B) Jumps back to its excited state

(C) Gives off energy

(D) Collapses into the nucleus.

In order to be promoted, it absorbs energy. When it relaxes, it releases it.

4. An atom composed of 8 protons, 10 electrons, and 8 neutrons is

  1. 16O2+

b.  16O2-

  1. 18O
  2. 18O2-

5. What does a group refer to on the periodic table?

a.  a column

  1. a row
  2. a period
  3. one element individually

6. Among the groups of elements listed below, which have the same number of energy levels?

a.  Li, Na, K, Rb

b.  Na, Mg, B, C

c.  K, Ca, Rb, Sr

d.  N, P, As, Sb

Note: on the practice test I handed out, there is no correct answer. The must be in the same row (also called period to have the same # of energy levels). If we change answer B to Na, Mg, Al, Si it is correct.

7. How many valence electrons does strontium (Sr) have?

a. 1

b.  2

c.  88

d.  38

Look for the group number.

8. Which of the following has the greatest electronegativity?

  1. Calcium

b.  Gallium

  1. Barium
  2. Thallium

Gallium and Thallium have the same effective nuclear charge, but Gallium has less electron shielding, so it will have a greater pull on bonded electrons. You can also think about it as Ga being closest to the northeast of the periodic table.

9.  Which of the following elements has the largest atomic radius?

a.  Potassium

  1. Sodium
  2. Gallium
  3. Krypton

Potassium has more energy levels than sodium and a lower effective nuclear charge than the other elements in its period.

10. Which of the following elements has an electron configuration that ends with p3?

  1. Boron
  2. Tin

c.  Antimony

  1. Neon

11. Which has the longest wavelength of the options below?

a.  Infrared

  1. Red
  2. Green
  3. Ultraviolet

12. Which of the following statements are true?

I. The longer the wavelength, the more energy

II. The longer the wavelength, the higher the frequency

III. The higher the frequency, the higher the energy

a) I only b) I + II c) III only d) I + III e) None of the above

Longer wavelength equals lower frequency and less energy.

For these 3 questions, match each element with the correct electron configuration. Some answers will not be used.

A = 1s1 B = [Ar]4s23d6 C = [Kr]5s24d105p3 D = [Kr]5s2

E = [He]2s22p1 AB = [Ne]3s23p2

13. Sr __D___

14. H __A___

15. Si __AB___

16. Which of the following best explains why fluorine is more electronegative than boron?

A)  Electronegativity is the measure of how negative an atom is.

B)  Fluorine has more electrons than boron

C)  Fluorine has a higher effective nuclear charge

D)  Fluorine is smaller than boron.

Be careful, while B + D is true, the important factor is the effective nuclear charge.

17. Which of the following best explains why sodium is smaller than potassium?

A) Sodium has a higher nuclear charge

B) Sodium has fewer valence electrons

C) Sodium has a higher electronegativity

D) Sodium has fewer energy levels

18) Which of the following is NOT a difference between the Bohr model and quantum mechanical model.

a. Electrons are in orbits for the Bohr model and orbitals for the quantum mechanical model.

b. Electrons are in energy levels for the Bohr models and the quantum model doesn’t address energy.

c. The Bohr model and quantum models are the same

19) Which of the following groups will have similar properties?

a.  Li, B, C, F

b.  Na, Mg, Al, S

c.  K, Ca, Rb, Sr

d.  N, P, As, Sb

Elements in the same group have similar properties because they have the same number of valence electrons.

20) What is the valence energy level for S?

a. 3 b. 6 c. 16 d.4

The valence energy level is also equal to the period #.

Open Response:

1)  Compare and contrast the Bohr and quantum mechanical atomic models.

Both the Bohr model and quantum mechanical talk about the electrons around the nucleus of an atom. They both talk about the electrons in terms of energy levels (no electron can be between energy levels in either model). However, the Bohr model shows electrons orbiting the nucleus in a circular pattern (like planets around the sun), while the quantum mechanical model talks about electrons in orbitals. Orbitals are 3-D shapes that are defined by the probability of finding an electron in that space (since we can never exactly know where an electron is). Also, the energy levels in the quantum model are broken into sublevels (or subshells).

2)  Fill in the chart below

Element / Electron Config. / Valence level / Valence electrons
Be / 1s22s2 / 2 / 2
Co / [Ar]4s23d7 / 4 / 2
Pd / [Kr]5s23d8 / 5 / 2
Sr / [Kr]5s2 / 5 / 2
F / 1s22s22p5 / 2 / 7
Mn / [Ar]4s23d5 / 4 / 2
Al / [Ne]3s23p1 / 3 / 3
Br / [Ar]4s23d104p5 / 4 / 7
Rn / [Xe]6s24f145d104p6 / 6 / 8

3)  Draw the Bohr model and box and arrow diagram for Calcium

Bohr Model
/ Box and Arrow Diagram

4)  On the periodic table below draw arrows that describe the trends in electronegativity and atomic size across a period and down a group. Below explain those trends in terms of effective nuclear charge and electron shielding.

Atomic Size

Electronnegativity

E. A.

Neg size

Atomic size decreases across a period because the effective nuclear charge increases. The effective nuclear charge is equivalent to the force felt by the valence electrons. Since , within a period, every element has the same number of energy levels, the elements on the right will be pulling their outer electrons closer, making atoms of those element smaller. Going down a group, the size increase because you are adding energy levels which are further away from the nucleus.

Electronegativity increases across a period because the effective nuclear charge increases. This means that electrons that are in a bond (always valence electrons) will be held more tightly by atoms on the right of the periodic table (except for the noble gases, which have 0 electronegativity). The electronegativity decreases down a group because the increased number of energy levels shield the valence electrons from the nucleus, meaning they will be held less tightly.