CLASS COPY – ANSWER QUESTIONS ON YOUR OWN PAPER
Review Guide For Test 1
Atomic Structure and the Periodic Table
1. What experiment did Thomson perform and what did he discover?
Cathode ray tube – beam of negative charge running from cathode to anode
Discovered electrons – tiny particles with no mass
2. What experiment did Rutherford perform and what did he discover?
Positive alpha particles shot a gold foil and deflected off the nucleus of the atom
Discovered protons – positive charge in the atom that deflect positive alpha particles
3. What did Chadwick look at specifically and what did he discover?
He looked at the nucleus and determined there was more mass than just protons
Discovered neutrons – particles in the nucleus that have the same mass as protons
4. Describe Bohr’s model.
Model of the atom where protons and neutrons are in the nucleus and electrons orbit the
nucleus in specific patterns or energy shells
5. What makes up most of the mass of an atom and where is this located in the atom?
Protons and neutrons, located in the nucleus
6. Given the model below, draw a Bohr model of Chlorine including the correct number and
location of protons, neutrons and electrons (remember Bohr’s electron rules!).
7. Write an equation for how you would find the number of
neutrons. (Hint: you will need 3 things).
Atomic mass – protons = neutrons
For isotopes: mass # – protons = neutrons
8. How do you find the number of TOTAL electrons an
element has?
Same as atomic number or # of protons
9. Why are groups “special” on the periodic table?
Each group contains special properties that are unique to
that group. Ex. Alkali metals are very reactive with water,
halogens have 7 valence electrons etc.
10. What is the difference between periods and groups?
Periods are horizontal – tell you # of shells
Groups are vertical – tell you # of valence electrons
11. How do you find the number of VALENCE electrons an element has?
Above each group (remember groups 13-18 have 3, 4, 5, 6, 7, 8 valence electrons)
12. How do you find the number of energy shells (orbitals) an element has?
Look at the period number
13. Use Calcium (Ca) to answer the following questions:
a. What group is Calcium in on the periodic table (name!)
Alkali earth metals
b. How many valence electrons does Calcium have?
2 valence electrons
c. How many energy shells does Calcium have?
4 shells
14. What is the magic number for a stable, outer shell of electrons? Which group (name) has
accomplished this?
8 electrons
CLASS COPY – ANSWER QUESTIONS ON YOUR OWN PAPER
15. What is the difference between atomic mass and mass number?
Atomic number = average mass of all isotopes of an element, found on periodic table
Mass # = the mass of one isotope (one version of the element), always a whole number
16. What is an isotope?
A version of an element with the same number of protons (atomic #) and different number of
neutrons and therefore mass #
17. Draw an isotope symbol for Fluorine-20.
a. How many neutrons does Fluorine 20 have?
𝟐𝟎
𝐅 20 – 9 = 11 neutrons
𝟗
18. Is Chlorine-18 an isotope of Chlorine-17?
Yes because they both have the atomic number 17 because they are both Chlorine atoms but
they have a different mass number and therefore a different number of neutrons
14
13
19. Is C an isotope of C? Why?
6
6
Yes because they both have the same atomic number, 6 and a different mass number and
therefore a different number of neutrons
For questions 20-25, describe each trend and be able to explain that you know what each trend
means!
**Consider the following factors: # of protons, # of valence electrons, # of shells
20. Why do the elements on the LEFT side of the periodic table have a larger atomic radius?
Elements on the left side of the periodic table have MORE protons, which pull on the electrons
in the outer shell LESS, allowing the atom to be bigger
21. Why do the elements on the BOTTOM of the periodic table have a larger radius?
Elements at the bottom of the periodic table have more energy shells and therefore a larger
radius
22. Why do the elements on the RIGHT side of the periodic table have a higher ionization energy?
Elements on the right side of the periodic table have more valence electrons and are closer to a
full outer shell and are therefore are harder to pull away
23. Why do elements on the TOP of the periodic table have a higher ionization energy?
Elements on the top of the periodic table have less energy shells so electrons are closer to the
nucleus (tight pants) and are harder to pull away
24. Why do elements on the RIGHT side of the periodic table have a higher electronegativity?
Elements on the right side of the periodic table have more valence electrons and the desire for
more or attraction to electrons is high
25. Why do elements on the TOP of the periodic table have a higher electronegativity?
Elements at the top of the periodic table have less energy shells so electrons are closer to the
nucleus (tight pants) and are more attracted to the positive nucleus
CLASS COPY – ANSWER QUESTIONS ON YOUR OWN PAPER
26. Use the Bohr models below to predict AND EXPLAIN each trend.
a. Which element has a larger
atomic radius? WHY?
Be – less protons, less pull
b. Which element has a higher
ionization energy? WHY?
N – more valence electrons,
harder to pull away
c. Which element has a higher
electronegativity? WHY?
N – more valence electrons,
desire more for a full shell
27. Use the Bohr models below to predict AND EXPLAIN each trend.
Selenium (Se)
a. Which element has a larger
atomic radius? WHY?
Se – more shells
b. Which element has a higher
ionization energy? WHY?
O – less shells, electrons
closer to nucleus, harder
to pull away
c. Which element has a higher
electronegativity? WHY?
O – less shells, electrons
closer to nucleus, more
attraction to positive
nucleus
28. What are the characteristics of metals? Give an example from the periodic table.
Lustrous (shiny), malleable, ductile, conduct heat and electricity well, high melting point
Anything left of metalloids (zig-zag)
29. What are the characteristics of nonmetals? Give an example from the periodic table.
Dull, brittle, not good conductors of heat and electricity, low melting point
Anything to the right of metalloids (zig-zag)
30. What are the characteristics of metalloids? Give an example from the periodic table.
Intermediate properties between metals and nonmetals
Anything in the zig-zag
31. What are the properties of a solid?
Molecules are close together and orderly, held together by strong forces, do not flow, volume
and shape are constant (definite)
32. What are the properties of a liquid?
Molecules are farther apart than a solid, they flow, difficult to compress, volume is constant
33. What are the properties of a gas?
Molecules are far apart, they flow, can be compressed, no definite shape or volume
CLASS COPY – ANSWER QUESTIONS ON YOUR OWN PAPER
34. What is one difference between the states of matter (solid, liquid, gas) and one similarity?
Similarity – all made of matter (atoms), water can be all of them, can become one another (solid
 liquid, liquid  gas etc.)
Difference – all different states, different distance between particles, different properties
35. How did you characterize your substance from our state of matter lab? What state of matter did
you decide it was and why?
Use properties of solid, liquid and gas
36. Synthesis question:
 Describe the element Sulfur (S) on the periodic table by predicting its properties given its
position and the fact that it is a solid at room temperature.
 Include:
Properties: metal vs. nonmetal, solid/liquid/gas properties, valence electrons, shells, reactivity
Trends: electronegativity, ionization energy, atomic radius  compared to the left side or the
bottom of the table.
Sulfur
 6 valence electrons, 3 shells
 Moderately reactive, wants to get to 8 electrons in outer shell but
 Solid – difficult to compress, does not flow, volume and shape constant
 Nonmetal – dull, brittle, does not conduct, low melting point
 Atomic radius – small because only 3 shells and low number of protons compared to
others
 Ionization energy – high because 6 valence electrons (hard to pull away) and only 3
shells (close to nucleus and harder to pull away)
 Electronegativity – high because 6 valence electrons (wants more) and only 3 shells
(close to nucleus and attracted to positive charge)