Download Periodic Table and Trends Test Review KEY Describe the common

Periodic Table and Trends Test Review KEY
Describe the common characteristics and uses of the following families:
1. Alkali metals (# of valence electrons = 1)
• Most reactive metals that do not occur freely in nature
• Slivery-white and softer than most other metals (to the point that they can be cut easily with a knife)
• can explode if they are exposed to water
2. Alkaline earth metals (# of valence electrons = 2)
• They are harder, denser, and stronger than the alkali metals with higher melting points
• They are so reactive they are not found as free elements in nature
3. Transition metals (# of valence electrons = 2)
• Metals with typical metallic properties and uses
• Less reactive than group 1 and 2 metals – some are so unreactive that they exist in nature as freeelements
• Used in jewelry, coins, building materials
4. Inner Transition Metals (# of valence electrons = 2)
• The lanthanides (top) are shiny and similar in reactivity to the alkaline-earth metals.
• The actinides (bottom) are all radioactive and most are man-made
5. Halogens (# of valence electrons = 7)
• "halogen" means "salt-former" and compounds containing halogens are called "salts"
• Most reactive non-metals
• exist in all three states of matter at room temperature
6. Noble gases (# of valence electrons = 8)
• Nonreactive due to their complete valence shells
• Do not easily gain or lose valence electrons
• All gases at room temperature
• used in lighting and balloons
7. Metalloids
• Stair step between metals and non-metals
• Properties in between metals and non-metals
o Si: brittle but semi-conductor
• Many metalloids are used in computer chips to conduct electricity without conducting too much heat
8. Metals
• Left side of the periodic table
o except H
• Properties:
o lustrous (shiny)
o good conductors of heat & electricity
o malleable & ductile
o solids at RT (except Hg)
9. Nonmetals
• Right side of table
o plus H
• Properties:
o Dull appearance
o Brittle when solids
o Do not conduct heat or electricity well
o May be solid, liquid or gas at RT
Summarize the scientist’s contribution to the development of the periodic table:
10. Dmitri Mendeleev
• designed the first periodic table based on:
o Elements with similar properties grouped together
o Increasing atomic mass
o Left blank spots, predicting properties of elements that had not yet been discovered.
11. Henry Moseley
• Rearranged the periodic chart to its current form based on increasing atomic number elements with
similar properties grouped together.
o This solved problem of misplaced elements
• Periodic Law - when elements are arranged by increasing atomic number, there is a periodic repetition
of chemical and physical properties
Label the families on the periodic
table (including the metals,
nonmetals, and metalloids):
Practice Questions:
12. The modern periodic table is arranged according to increasing atomic number.
13. Electrons in the outermost energy level are known as valence electrons and are available to be lost, gained
or shared when molecules are formed.
14. Rows on the periodic table are called periods and columns on the periodic table are called groups or
families.
15. Elements in the same family have similar properties.
16. In general, ionization energies of the elements increase across each period.
17. Among the elements, ionization energies generally decrease down each group.
18. The measure of the ability of an atom in a chemical compound to attract electrons is called
electronegativity.
19. In general, this ability to attract electrons increases across each period.
20. In general, the atomic radii of the element increases down a group.
21. Arrange the following in order of increasing ion size: Cl ion, Al ion, P ion, Mg ion, Na ion
Al3+ ! Mg2+ ! Na+ ! Cl- ! P317) Arrange each group of elements in order of increasing ionization energy.
a. F, Br, I, Cl
I ! Br ! Cl ! F
b. Ga, Al, Tl, B
Tl ! Ga ! Al ! B
c. Al, Si, Cl, S Al ! Si ! S ! Cl
18) For each group of elements, choose the one with the largest atomic radius.
a. He, Rn, Xe, Ar
Rn
b. As, N, P, Bi
Bi
c. Ba, Hf, Os, Hg
Ba
22. Why does fluorine have a higher ionization energy than iodine?
Fluorine is in a lower energy level, meaning the electrons are closer to the nucleus, thus increasing the
coulombic attraction. Also, Iodine has more shielding present due to the multiple energy levels between the
nucleus and the valence electrons.
23. Why do elements in the same family generally have similar properties?
They have the same number of valence electrons
24. What happens to the size of an atom when it becomes an anion? Why?
Anions are atoms that have GAINED electrons to become NEGATIVELY charged. With the additional
electrons, the proton-electron ratio has decreased, decreasing the coulombic attraction. Also, with
additional electrons, there is an increase in electron-electron repulsion, forcing the electrons to move away
from each other, increasing the size of the ion.
25. What happens to the size of an atom when it becomes a cation? Why?
Cations are atoms that have LOST electrons to become POSITIVELY charged. This causes an increase in the
proton-electron ratio, increasing the coulombic attraction.
26. Define each of the following periodic trends, identify their trends along a period and down a group, and
why this happens:
• Atomic Radius
Size of an atom, distance from the center of an atom's nucleus to the edge of its electron cloud.
Period trend - decreases across a period because of the increase in positive charge of nucleus
(more p+ added) creates a greater Coulombic attraction on the e– towards nucleus while the new
e– are added within same energy level
o Group Trend - increases down a group because e– are added to higher energy levels, farther
from nucleus, making radius larger
Ionization Energy
o energy required to overcome Coulombic attraction and remove one valence electron
o ionization energy decreases down a group because The farther away from the nucleus an
electron is, the easier it is to remove because the Coulombic attraction of the protons has less
hold on the negative charge of the electron.
o Ionization energy increases across a period because the increased number of protons in the
nucleus, causes an increase in the force of attraction
Electronegativity
o tendency of an atom to attract e– towards its nucleus when chemically combined (bonded) with
another element
o electronegativity decreases down a group because because The farther away from the nucleus an
electron is, the harder it is to pull because the Coulombic attraction of the protons has less hold
on the negative charge of the electron.
o electronegativity increases across a period because the increased number of protons in the
nucleus, causes an increase in the force of attraction
o
o
•
•
27. Describe each of the states of matter in terms of shape, volume, and compressibility. Include a picture of the
state of matter at the molecular level.
28. What must happen in order for a change in state of matter to occur?
Add or remove energy