THE PERIODIC TABLE
DIRECTIONS: Answer correctly the following questions.
1. Who is considered to be the ‘father’ of the periodic table? Dimitri Mendeleev
2. How were the elements arranged? By atomic mass and chemical and physical properties
3. What does periodic mean? Repeats in a regular pattern
4. How did Mendeleev predict in his Periodic Table the future discovery of then unknown elements?
He left spaces in his table where unknown elements should exist based on properties and mass
5. Mendeleev noticed that properties of elements appeared at regular intervals when the elements were
arranged in order of increasing:
A. density
B. reactivity
C. atomic number
D. atomic mass
6. What was Henry Moseley’s contribution to the periodic table? Arranged by Atomic Number, not
atomic mass
7. The modern periodic law states that:
A. no two electrons with the same spin can be found in the same place in an atom
B. when elements are arranged according to increasing atomic number, there is a periodic
repetition of their chemical and physical properties
C. electrons exhibit properties of both particles and waves
D. the chemical properties of elements can be grouped according to periodicity, but physical
properties cannot
8. In a modern periodic table, every element is a member of both a vertical column and a horizontal row.
Which one is the group, which one is the period and which one is also called the family?
Group: vertical columns, also called “family”; Periods: rows
9. On the periodic table, where are the:
A. main-group elements? Groups 1, 2, 13-18
B. transition elements? Groups 3-12
10. True or false:
A. __T__All elements are either metals, non-metals or metalloids.
B. __F__Most elements are non-metals
C. __T__Group 1 elements (except for hydrogen) are known as the alkali metals
D. __T__The transition elements include iron, copper, gold, nickel and zinc.
E. __F__Group 13 elements are known as the alkaline earth metals
F. __T__Inner transition metals are also called the f block elements
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The Periodic Table
DIRECTIONS: Use only the below periodic table to answer correctly the following questions.
11. What is the purpose of the dark ‘stairs’ on the right side of the periodic table?
Separates metals from non-metals; is where the metalloids are located
12. Name the group identified by each of the following letters:
A. s-block
B.
d-block or transition metals
C. p-block
D.
f-block or inner transition metals
E. Alkali metals
F.
Alkaline earth metals
G. Halogens
H.
Noble gases
I.
J.
Actinide series
Lanthanide series
13. Provide several properties for each of the following groups (from notes, book or online):
E.
1 valence electron = s1. Highly reactive, soft, shiny, lustrous, form hydroxides in water,
very low ionization energy
F.
2 valence electrons = s2. Fairly reactive, harder than alkalis, form hydroxides in water,
fairly low ionization energy
7 valence electrons = s2p5. Highly reactive, mostly gases and liquids, brittle when solid,
G. very high electronegativity
8 valence electrons (full octet) = s2p6. Non-reactive, gases at room temperature, full octet
H. of valence electrons
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The Periodic Table
14. Write the name and the noble-gas configuration for each of the following elements:
A. Group 14 (IVA) element in Period 4:
= Ge, Germanium, [Ar] 4s2 3d10 4p2
B. The only metal in Group 15 (VA): Bismuth (Bi) = [Xe] 6s2 4f14 5d10 6p3
C. The transition element with the smallest atomic mass: Scandium (Sc) = [Ar] 4s2 3d1
D. The alkaline earth metal with the largest atomic number: Radium (Ra) = [Rn] 7s2
15. What are five major properties of metals?
1) Highly conductive of energy _____
4) shiny (and lustrous) __________________
2) Malleable (flatten into thin sheet)
5) solid at room temp (high melting points) _
3) Ductile (draw into wires) _______
16. Which tends to be more reactive: alkali metals or alkaline earth metals? Alkali Metals
17. An ionic compound that is brightly colored generally contains which type of metal? Transition Metals
18. True or false:
A. T____Group 18 elements are very unreactive
B. F____Fluorine has a higher electrical conductivity than Chromium
C. T____Sulfur is an element that is brittle and does not conduct hear or electricity well
D. T____Out of Se, Si, P and Mg, only Mg is soft and malleable in its solid state
E. F____Most Halogens are solid at room temperature
19. Transition Metals - d-Block Elements
Transition metals generally have 2 valence electrons.
Transition metals generally have a +2 ionic charge or oxidation state
Define Ferromagnetismproperty of some metals like iron which allow them to hold a magnetic force.
20. Inner transition Metals - f-Block Elements
Lanthanide Series = Elements # 58 to 71
Lanthanides are sometimes called “trace” elements. What does this mean?
“Trace” means to be found in very small amounts.
Actinide Series = Elements # 90 to 103
Actinides are all radioactive and those with a higher atomic number than Uranium are also
referred to as “transuranic” and are man-made (synthetic) elements which are not normally
found in nature .
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The Periodic Table
CHEM WKST
PERIODIC TRENDS
21. DIRECTIONS:
List the valence electrons for the group, then note how realatively reactive the
group is (i.e. reactivity is: high, moderate, fair or low)
# Valence Electrons
Relative Reactivity
Group 1 Alkali Metals
___1_____
Very High
Group 2 Alkaline Earth Metals
___2_____
Moderately High
Group 13 Boron Group
___3_____
Moderate
Group 14 Carbon Group
___4_____
Low to Fair
Group 15 Nitrogen Group
___5_____
Moderate
Group 16 Oxygen Group
___6_____
Moderately High
Group 17 Halogens
___7_____
Very High
Group 18 Noble Gases
___8_____
None
22. The energy required to remove an electron from an atom is the atom’s ________________________
A. electron affinity
B. electron energy
C. electronegativity
D. ionization energy
23. Moving from left to right across a period on the periodic table,
A. ionization energy values tend to become larger or smaller (circle correct answer)
B. atomic radii tend to become larger or smaller (circle correct answer)
24. Name:
A. the alkali metal with the highest ionization energy:
A. the element in Period 3 with the smallest atomic radius:
Lithium (Li)
Argon (Ar)
25. Define the following terms:
A. atomic radius: distance between nucleus and outer edge of neutral atom’s electron cloud
B. ionization energy: energy need to remove an electron from an atom
C. ionic radius: distance between nucleus and outer edge of ION’s electron cloud
D. valence electrons: number of electrons in atom’s outermost shell (s and p sublevels)
E. anion: a negatively charged ion (an atom with more electrons than protons)
F. cation: a positively charged ion (an atom with more protons than electrons)
G. nuclear charge: overall attraction for electrons by the positively charged nucleus
26. Rank in order of increasing ionization energy: Xe, Te, Mo, Sn. .
1: Mo 2: Sn, 3: Te, 4: Xe
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The Periodic Table
27. Explain the shielding effect and its importance in ionization periodic trends down groups.
The shielding effect is caused by electrons located in inner shells blocking the outer electrons
from the attraction of the positive nuclear charge at the center of the atom. This causes the
ionization energy to decrease as you go down the groups (and thus increasing the number of
shells).
28. The size of the atom gets bigger as one goes down the column of a group on the Periodic Table is due
to two factors: Shells and Shielding increase
29. As one goes from left to right across a period, the atomic radius decreases because of the increased
attraction of the electrons by an increased number of: protons
30. Explain the trend for ionization energy going across the periods.
Increasing net nuclear charge as protons are added increases the attraction on electrons
making it harder to remove them.
31. As one moves down a group, the reason that the ionization energy decreases, making it easier to
remove an electron is that the electrons are farther from the: positively charge nucleus.
32. Define electronegativity: the relative ability of an element’s atoms to attract electrons
33. Moving from left to right across a period on the periodic table, electronegativity values tend to
become: larger or smaller (circle correct answer)
34. Name:
A. the element with the highest electronegativity value:
B. the alkali metal with the highest electronegativity value:
C. the element in Period 2 with the smallest atomic radius:
D. the Group 14 element with the lowest electronegativity:
35. Rank in order of decreasing electronegativity: F, B, Be, C. .
Fluorine
Lithium
Neon
Lead
1: F, 2: C, 3: B, 4: Be
36. Explain the trend for electronegativity going across the periods.
Again, the increasing net nuclear charge as protons are added increases the attraction of that
atom for electrons in nearby atoms (i.e. atoms that it is chemically reacting with).
37. How does an atom turn into:
A. an anion? Gain electrons ________________________
B. a cation? Lose electrons _________________________
give an example: F-1
give an example: Ca+2
38. What is the octet rule? Atoms will gain or lose valence electrons in a reaction so that they end up
with eight electrons in valence shell (an “octet”). This means that both the s and p sublevels on
the outermost energy level are filled = very stable situation. Metals gain and non-metals lose
electrons to get to this stable electron configuration which is the same as their nearest noble gas.
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The Periodic Table
39. Use arrows on the following periodic tables to indicate increasing trends, by period and group, for
the following properties:
A. Atomic Radii
B.
Ionization Energy
\
C. Electronegativity
D.
Ionic Radius
40. A negative ion is ( smaller / larger ) than its parent atom. Why? (more electrons than protons)
41. A positive ion is ( smaller / larger ) than its parent atom. Why? (more protons than electrons)
42. Use the periodic table to write the symbol, the most likely oxidation number (ionic charge) and to
describe the key chemical properties of the following elements:
B. fluorine: F-1; 1s2 2s2 2p6 (like Ne); very reactive since it needs just one more electron for
an octet and it has a high electronegativity due to its small atomic radius and high net
nuclear charge (upper right position on the table).
C. xenon: Xe will not form an ion. Not reactive since it already has a full octet:
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6
D. magnesium: Mg2+ fairly reactive with somewhat low ionization energy
1s2 2s2 2p6
E. sodium: Na+1; 1s2 2s2 2p6 (like Ne); very reactive since it needs to lose only one electron for
an octet and it has a fairly low ionization energy.
F. oxygen: O2- 1s2 2s2 2p6 (like Ne)fairly high reactivity due to high electronegativity
G. gold: Au1+, low reactivity due to fairly low electronegativity and fairly high ionization
energy. Loses its one 6s electron 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p66s1 4f14 5d10
2. Where are the most active metals located and why are they so reactive? The alkali metals are most
reactive because they have a very low ionization energy and want to lose their one s 1 electron to
form an octet.
3. Where are the most active nonmetals located and why are they so reactive? The halogens are the
most reactive because of their very high electronegativities and their desire to add one more
electron to form an octet.
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The Periodic Table
4. As you go from left to right across a period, the atomic size ( decreases / increases). Why?
Increased nuclear charge (protons), but no additional shielding or shells.
5. As you go down a group/family, the atomic size ( decreases / increases ). Why? Although nuclear
charge does increase, this factor is overwhelmed by the increased size due to additional shells
and due to more shielding.
6. As you go down a group/family, the first ionization energy generally: ( decreases / increases ).
7. The tendency of an atom to attract other electrons to itself is referred to as its: electronegativity
8. Where is the atom with the following characteristic found on the periodic table (list name, group
name and period)?
A. highest electronegativity: fluorine, halogen group, period = 2
B. lowest electronegativity: francium, alkali metal, period = 7
9. Respond to the following questions:
A. Rank the following elements by increasing atomic radius: carbon, aluminum, oxygen, potassium.
O, C, Al, K
B. Rank the following elements by increasing electronegativity: sulfur, oxygen, neon, aluminum.
Ne, Al, S, O
C. Why does fluorine have a higher ionization energy than iodine?
Smaller size (fewer shells) and less shielding
D. Why do elements in the same group tend to behave similarly in chemical reactions?
Same number of valence electrons and similar relative ionization energies and
electronegativities compared to other elements in their period.
10. Use periodic trends we have learned to:
A. Rank in order of increasing atomic radius: K, Na, Rb.
Na, K, Rb
B. Rank in order of increasing atomic radius: N, O, F.
F, O, N
C. Rank in order of increasing ionization energy: Ca, Mg, Ra.
Ra, Ca, Mg
D. Rank in order of increasing ionization energy: Xe, Te, Sn.
Sn, Te, Xe
E. Rank in order of decreasing electronegativity: Li, Na, Cs.
Li, Na, Cs
F. Rank in order of decreasing electronegativity: Si, Cl, P.
Cl, P, Si
G. Rank in order of decreasing ionic radius: Be, Sr, Mg.
Sr, Mr, Be
11. Key point: Why is the electronegativity of an element such an important factor in determining its
chemical properties? It ranks the ability of an element to attract an electron from another atom.
This attraction, whether it results in the transfer of an electron (ionic bonding) or the sharing of
the electron (covalent bonding) is the underlying cause of all chemical reactions.