Download Topic 3 Periodicity File

Topic 3
Periodicity
‘Two atoms walking down the road.
One says to the other "Oh no, I think I've
lost an electron".
"Are you sure?" says the other.
"Yes, I'm positive" he replies.’
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Definitions
Alkali metals: Group 1 elements.
Catalyst: A substance that increases the rate of reaction while being
recoverable unchanged at the final stage of the reaction. Examples of catalytic
transition metals:
Fe is used in Haber process;
V2O5 in Contact process;
Ni in hydrogenation reactions;
MnO2 in the decomposition of hydrogen peroxide
Catalyst, heterogeneous: In different state than reactants
Catalyst, homogeneous: In the same state as reactants
Colored complex: A complex is a compound in which molecules or ions form
dative bonds to a metal atom or ion. Colors are due to e- transitions between
different d orbitals.
Co-ordination number: Number of lone pairs bonded to the metal ion.
Cl- often gives 4 coordinate bonds, CN- gives 6, H2O gives 6 and NH3 gives 4
or 6.
Electronegativity: Relative measure of the ability an atom has to attract a
shared pair of electrons.
Group: Elements with the same number of valence e-.
Halide ions: Ions of the halogens.
Their presence can be detected by the addition of silver nitrate. AgCl is
white, AgBr is cream-colored, and AgI is yellow. Silver halides react with
light to form silver metal.
Ligand: A molecule or ion that can donate an electron pair.
Metalloid: An element that possesses some of the properties of a metal and
some of a non-metal. While metal oxides tend to be basic and non-metal oxides
tend to be acidic, metalloid oxides such as aluminium oxide can be amphoteric.
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Transition element: An element that possesses an incomplete d sub-level in one
or more of its oxidation states.
Often very efficient catalysts as they can exist in a variety of oxidation
states (all except Ti have oxidation state of +2). Form coloured complexes.
Web Links
Interactive Periodic Table
Link to Physics 2000 : Go to science trek: Elements as atoms: Beyond hydrogen
http://www.colorado.edu/physics/2000/index.pl
Weblinks for finding information:
http://www.dayah.com/periodic/
http://www.rsc.org/chemsoc/visualelements/PAGES/pertable_fla.htm
http://www.webelements.com/
http://www.chemicalelements.com/index.html
Now have a go at some online quizzes
http://www.funbrain.com/cgibin/pt.cgi?A1=s&A2=1&ACOMMON=1&submit=Play+Proton+Don
http://quizhub.com/quiz/f-elements.cfm
http://www.sporcle.com/games/elements.php
http://www.sheppardsoftware.com/Elementsgames.htm
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Electron Shells and the Periodic Table
Refer to the periodic table to answer the following questions.
1.
In the periodic table, what does the group number refer to?
2.
In the periodic table, what does the period number refer to?
3.
List, by name and symbol, all the elements in group 5 of the periodic table.
4.
How many electrons are in the outermost shell of all the elements in group 5?
5.
What element in group 2 of the periodic table has three electron shells?
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6.
Complete the following table:
Atom
Number of
Number of
electron
electrons in
shells
outermost
shell
In a neutral atom
Number of
Number
protons
of
electrons
Lithium
Silver
Cadmium
Nitrogen
Chlorine
Neon
Argon
Gold
Lead
Magnesium
Sc
Zr
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The Periodic Table Questions
Q1 Use the periodic table to write down the atomic symbols for:
(a) gold
(b) silver
(e) iron
(c) lead
(f) sodium
(d) copper
(g) aluminium
(h) zinc
Q2 Write down the element whose symbol is:
(a) N
(b) I
(c) C
(d) Mg
(e) O
(f) Cl
(g) Hg
(h) He
(i) Ne
Q3 State the atomic number of:
(a) Rubidium
(b) Hafnium
(c) Indium
(d) Tungsten
(e) Cd
(f) Pu
(g) K
(h) Es
Q4 List the elements of group IV from lightest to heaviest.
i)……………………….ii) …….. …………………iii) …………….……………
iv) …..…………………v)………………………….
Q5 List two elements named after planets:
i)
………………………… ii)…………………………..
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Q6 The following elements were named after scientists. Can you name the scientists?
a) Es………………………………
b) Curium……………………………….
c) Fm …………………………….
d) No …………………………………..
e) Md…………………………….
f) Lr……………………………………
Q7 a) How many atomic symbols begin with the letter C? ……………………..
b) Which letters of the alphabet have not been used as an atomic symbol? …….
Q8
How many elements are there between Lanthanides and Hafnium? ………….
Q9
Name the heaviest atom of the periodic table? ……………………………
Q10
Name the element that can be described as
(a) period 4 group II …………………………………
(b) period 6 group VIII ……………………………….
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Periodic Table Questions
1.
The diagram shows an outline of part of the Periodic Table of Elements.
H
region 3
region
1
region 2
(a)
What is the name of the element with the symbol H?
………………………………….
1 mark
(b)
In which regions of the Periodic Table are the following types of element
found?
(i)
non-metals (such as oxygen and chlorine);
region …………
1 mark
(ii)
very reactive metals (such as sodium and potassium);
region …………
1 mark
(iii) less reactive metals (such as copper and zinc).
Region …………
1 mark
(c)
Why is copper sulphate not found in the Periodic Table?
……………………………………………………………………………………….
……………………………………………………………………………………….
1 mark
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(d)
An iron nail is placed into some blue copper sulphate solution.
A reaction takes place between the iron and the copper sulphate.
(i)
Complete the word equation for the reaction.
iron + copper sulphate
……………………… +…………………………
1 mark
(ii)
Describe one change you would see on the surface of the nail.
………………………………………………….…………………………….
……….……………………………………………………………………….
1 mark
Maximum 7 marks
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Periodic Table
Locate, label and colour the following groups in your periodic table
1. Alkali Metals - lithium, sodium, potassium, caesium, francium
2. Halogens
- fluorine, chlorine, bromine, iodine, astatine
3. Transition metals - a large groups of common metals in the centre of the Periodic table
4. Noble Gases
- helium, neon, argon, krypton, xenon etc
5. The Boundary between metals and Non-metals
6. Label the Group numbers
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PHYSICAL PROPERTIES OF THE
PERIODIC TABLE
Term
Definition
Trend: Period 3
Explanation
Trend: Alkali
Metals
Explanation
Trend: Halogens
Explanation
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PHYSICAL PROPERTIES OF THE
PERIODIC TABLE
Term
Definition
Trend: Period 3
Explanation
Trend: Alkali
Metals
Explanation
Trend: Halogens
Explanation
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Periodicity Database Activity
Using the database skills you learned in Unit 2 (isotopes activity), you will create your own database
ofelements and their properties to illustrate the nature of periodicity in the periodic table. Some of
theproperties & trends you will learn as part of your class activities, others you will have to research on
yourown. Chapter 3 of your text will be useful towards understanding the trends. You will submit yourdatabase
for review and assessment at the end of the unit.
Your database must include the following:
Use your database skills to sort and answer the following questions. Capture a screenshot of
your sorting and paste below the answer to each question:
1. How does an element’s position on the periodic table relate to its electron arrangement?
Use examples between Z = 1 and Z = 20.
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2. How does an element’s position on the periodic table relate to the number of electrons
occupying the highest energy level? Use examples between Z = 1 and Z = 20.
3. Describe the trend in atomic radii for the alkali metals.
4. Describe the trend in melting point for the halogens.
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5. Describe the trend in first ionization energies for the elements across period 3
6. Which is the most electronegative atom? Which is the least electronegative atom?
7. Does reactivity of alkali metals with water increase or decrease as you go down the
group?
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8. Do metal oxides tend to be acidic or basic?
9. Do non-metal oxides tend to be acidic or basic?
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Periodic Trends Worksheet
1)
Rank the following elements by increasing atomic radius: carbon, aluminum, oxygen,
potassium.
2)
Rank the following elements by increasing electronegativity: sulfur, oxygen, neon, aluminum.
3)
What is the difference between electron affinity and ionization energy?
4)
Why does fluorine have a higher ionization energy than iodine?
5)
Why do elements in the same family generally have similar properties?
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Ionization Energies – A peek inside the atom
Line emission spectra strongly support the idea that electrons exist in discrete energy
levels – but theyaren’t the only forms of evidence. Other support comes from examining the
amount of energy tocompletely remove an electron from an atom. This energy is referred to
as the ionization energy,since when one or more electrons are completely removed from an
atom, the atom becomes apositively charged substance known as an ion.
By definition,
The first ionization energy of an element is the energy required to remove one
mole of electrons from one mole of atoms of the element in the gaseous state to
form gaseous ions.
Examples are shown in the following equations:
Na(g) →Na+(g) + e- First ionization energy = 496 kJ mol-1
Mg(g) →Mg+(g) + e- First ionization energy = 738 kJ mol-1
A knowledge of ionization energies provides valuable information about the arrangement of
electrons within atoms. The discussion of the ionization of an atom has so far considered
the removal of electron only; but if an atom containing several electrons is treated with
sufficient vigour, then more than one electron may be removed from it. A succession of
ionization energies is therefore possible. These may be determined, principally from
spectroscopic measurements; a table of successive ionization energies for a number of
elements is given in Table 1. just below.
Study Task
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Use Table 1 to answer the following questions. In all questions use a set of data that
includes 3 elements: either the three alkali metals (Li, Na, and K) or the three alkaline earth
metals (Be, Mg, Ca).
1. Examine the first ionization energies for the set of elements you are investigating.
Do you notice any trend? Explain why this trend exists.
2. Make a plot of ionization energy, on the vertical axis, against the number of electrons
removed, on the horizontal axis for your set of three elements (use the graph templates on
the next page).
• What do you notice about the general trend in values?
• Why do you think this trend in values occurs?
3. Now plot a graph of the logarithm (to base 10) of the ionization energy against the
number of electrons removed. You may want to use the Table 2. below to help you organize
your data and then use the graph templates on the next page.
• Does this type of plot give any information about groups of electrons that can be removed
more readily than others?
• How many electrons are there in each group?
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Investigating Trends in the Periodic Table
Purpose
In this exercise you will be investigating the relationship between the elements’ atomic numbers
and properties, such as ionization energy (kJ/mol), electronegativity, atomic radius (pm), ionic
radius (pm) and melting point (K) of elements in the Periodic Table using graphs.
Pre-Lab
Define the terms first ionization energy, electron affinity, electronegativity, atomic radius,
ionic radius and melting point.
Procedure
Predict the relationships between an element’s atomic number and each of the five properties
listed below. (Write one hypothesis per property.)
A.
B.
C.
D.
E.
First ionization energy
Electronegativity
Atomic radius
Ionic radius
Melting point
Using the periodic tables in your data booklet, construct a series of graphs for Atomic Number
vs. Atomic Radius and then each of the other properties (A through E) in the list above for
elements 3 – 38.
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You may place up to two sets of data on any one graph (ie. You may place atomic radius and ionic
radius on the same graph).
Make sure each graph has a name and is properly labeled with x- and y-axis variables and units.
Finally, create a division line that indicates the transition from one period to the next and place
a square around each element in Group 1.
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Conclusion
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This section should be thorough, but concise, and should respond to the Testable Question.
Address the answered in paragraph form, NOT as individual questions.
 Summarize the results/data that were obtained
1. Look at the shape of your graphs. What general patterns do you observe for each.
2. What is the trend shown for the alkali metals by each graph? Give evidence for your
answer.
3. What is the trend shown for the halogens by each graph? Give evidence for your
answer.
4. What is the trend shown across Period 3 by each graph? Give evidence for your
answer.
 Explain whether or not the results matched your hypothesis
 Explain why your results were similar to or different from your original hypothesis
 What is the significance of the results? What did you learn? What was surprising?
 State the relationship of the manipulated and responding variables
 Identify and explain a real world example/application of the concepts
 Present a new question to be answered in a future experiment / investigation
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Periodic Trends
Here we summarize trends for the main group elements (Columns 1A - 8A). Trends for
the transition metals, the lanthanides, and the actinides may differ.
Sizes of Atoms and Ions
Neutral Atoms (or Ions with the Same Charge).
•
Size increases as you go down a column. Why? As you go down a column, electrons are
•
Size decreases as you go across a row. In this case electrons are being added to the
same shell. Thus they experience little additional shielding. On the other hand, the
filling orbitals farther and farther out from the nucleus. Each row adds a new shell.
Outer electrons are shielded from the nucleus by electrons in inner shells; thus they are
less tightly held (in spite of the much increased nuclear charge).
nuclear charge of the atom increases with the atomic number. Thus as you go across a
row, the electrons are held more tightly and the size decreases.
Isoelectronic Series. These are series of atoms and ions in which the number of electrons
stays constant, but the number of protons increases with the atomic number. In this type
of series, the size of the atom decreases as the number of protons increases. The reason
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for the size decrease is that more protons are pulling in the same number of electrons.
Examples include the series below in which the largest member of the series is listed first:
10 electron series: Ne > Na+ > Mg2+ > Al3+
18 electron series: P3- > S2- > Cl – > Ar
Cation Size as Compared to Parent Atom. The size decreases when cations form. The
effect is particularly pronounced when all the valence electrons are lost and only the noble
gas core of electrons remains. For example, the Mg2+ ion (65 pm radius) is considerably
smaller than the Mg atom (160 pm radius).
Anion Size as Compared to Parent Atom. The size increases when anions form. The added
electrons are going into the same shell. They repel each other and so the size increases.
Thus the Cl– ion (181 pm radius) is considerably larger than the Cl atom (99 pm radius).
Ionization Energies
The ionization energy I is the minimum energy needed to remove an electron from the
ground state of a gaseous atom, A(g).
A(g) A+(g) + e–(g)
∆E = I = I1
More precisely, this is the first ionization energy I1. Additional electrons may be removed
with ionization energies I2, I3, etc., for the removal of the second, third, etc., electrons.
Ionization is always an endothermic process: it requires energy to remove an electron from
an atom or ion.
The overall trends in ionization energy are opposite to those for atomic and ionic radii.
The more tightly electrons are held, the higher the ionization energy, and the smaller the
atom or ion size. Some generalities are as follows:
•
Noble gases have the highest ionization energies of the atoms in each row.
•
Alkali metals have the lowest ionization energies of the atoms in each row.
•
In general, ionization energies increase as you go across a row, but there are a few local
ups and downs. Dips occur with the loss of the first and the fourth p electron: Thus in
the second row, there are dips for boron and for oxygen.
•
The ionization energy decreases for atoms as you go down a column.
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• Higher ionization energies are always larger than lower ionization energies: I1< I2< I3 ,
etc. A huge jump in ionization energy occurs when you first pull an electron out of the noble
gas core.
Electron Affinity
The electron affinity EA is the energy released when an electron is added to a gasphase atom (or ion) of the element. The sign convention is opposite to that for
. If the
process is exothermic, ∆H is (-) and EA is (+); if it is endothermic, ∆H is (+) and EA is (-).
A(g) + e–(g) A–(g)
∆H = - EA
While ionization energies are always positive numbers, electron affinities can be either
positive or negative. A high positive EA (and thus a (-) value of
) indicates that gaining
an electron is a very favorable process. The halogens have the most positive electron
affinities of all the elements.
Electronegativity
The electronegativity Χ (Greek letter chi) is a measure of the ability of an atom to
attract and hold electrons. Elements that readily form negative ions have high
electronegativities, while a low electronegativity correlates with the tendency to lose
electrons and form positive ions. Values of Χ range from a high of Χ = 4.0 for F to a low of
Χ = 0.7 for Cs. In general electronegativities increase diagonally from the lower left (Cs) to
the upper right (F) of the periodic chart.
In practice, chemists use electronegativities far more than ionization energies or
electron affinities.
_____________________________________________________________
Exercises:
1. In each of the following pairs, circle the species with the higher first ionization energy:
(a) Li or Cs
(b) Cl- or Ar
(c) Ca or Br
(d) Na+ or Ne
(e) B or Be
2. In each of the following pairs, circle the species with the larger atomic radius:
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(a) Mg or Ba
(b) S or S2-
(c) Cu+2 or Cu (d) He or H-
(e) Na or Cl
_____________________________________________________________
3. Circle the best choice in each list:
(a) highest first ionization energy:
C, N, Si
(b) largest radius: S2–, Cl–, Cl
(c) highest electronegativity: As, Sn, S
(d) smallest atom: Na, Li,
Be
(e) most paramagnetic: Fe,
Co,
Ni
(f) lowest first ionization energy: K, Na, Ca
(g) highest second ionization energy: Na, Mg, Al
(h) lowest second ionization energy: Ar, K, Ca
Answers (be sure you can explain the reason for each answer!):
1. (a) Li; (b) Ar (isoelectronic pair); (c) Br; (d) Na+ (isoelectronic pair); (e) Be (common
exception: what is the rule here?).
2. (a) Ba; (b) S2-; (c) Cu; (d) H- (isoelectronic pair); (e) Na.
3. (a) N; (b) S2- (S2- and Cl- are isoelectronic); (c) S; (d) Be; (e) Fe (hint: determine no. of
unpaired spins for each element); (f) K; (g) Na; (h) Ca.
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Periodicity Practice
1. How are elements organized on the periodic table?
2. What is the electron configuration of the following elements?
a. Helium
b. Carbon
c. Sodium
d. Argon
e. Calcium
3. How many energy levels are occupied for the elements from #2?
4. What is the relationship between the electron arrangement of elements and their position in the
periodic table? (Include in your explanation the relationship with groups/families and periods.)
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5. Define the following terms: group, period, alkali metals, halogens, ionic radius, electronegativity,
first ionization energy
6. What is the trend as you go down the periodic table from Li to Cs for:
a. Atomic radius
b. Ionic radius
c. First ionization energy
d. Electronegativity
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e. Melting point
7. What is the trend as you across the periodic table from Na to Cl for:
a. Atomic radius
b. Ionic radius
c. First ionization energy
d. Electronegativity
8. Between Mg, K and Ca which one has the highest electronegativity? Why?
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9. Compare the reactivity of the alkali metals. Explain.
10. Compare the reactivity of the halogens. Explain.
11. What type of bonds do metal oxides tend to form? Do they react with water to form acidic or
basic (alkaline) solutions?
12. What type of bonds do non-metal oxides tend to form? Do they react with water to form acidic
or basic (alkaline) solutions?
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Using Electronegativity
1) Use the Periodic Table to choose the element in each set that has the highest
electronegativity. Then explain your answer.
a) Si, As, Ge, P
b) P, Mg, Ba, Sb
c) B, F, Te, P
d) Ca, Ba, Zn, I
2) In each of the following bonds, choose the atom that carries the partial negative charge:
Then explainyour answer.
a) P – Br
b) Si - Cl
c) S - Cl
d) As – H
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e) F - C – Cl
f) Cs – F
3) In each of the following bonds, choose the atom that carries the partial positive charge:
The explain your answer.
a) N - S
b) Si - I
c) N - Br
d) Hg – I
e) P – I
f) Mg – N
4) For each of the bonds in question #3, indicate whether the bond is ionic, polar covalent
or covalent.
Then explain your answer.
a) _______________
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b) _______________
c) _______________
d) _______________
e) _______________
f) _______________
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Successive Ionization Energies
A graph of successive ionization energies vs number of electrons removed serves as
convincing evidence for the fact that the electron arrangements of lithium (Li), sodium (Na),
and potassium (K) are (2,1) ; (2, 8, 1) ; and ( 2, 8, 8, 1) respectively. Explain this statement
by discussing:
• What factors affect how much energy is required to remove a mol of electrons,
• What is the significance of a large difference in the amount of energy required to remove
a successive electron?
• What is the significance of small differences in the amount of energy required to remove
a set of successive electrons?
• How can the three points above be used to justify the electron arrangements of lithium,
sodium, and potassium?
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Periodic Trends
1. Identify each element as a metal, metalloid, or nonmetal.
a) fluorine
b) germanium
c) zinc
d) phosphorous
e) lithium
2. Give two examples of elements for each category.
a) noble gases
b) halogens
c) alkali metals
d) alkaline earth metals
3. What trend in atomic radius do you see as you go down a group/family on the periodic
table?
What causes this trend?
4. What trend in atomic radius do you see as you go across a period/row on the periodic
table?
What causes this trend?
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5. Circle the atom in each pair that has the largest atomic radius.
a) Al B
b) S O
c) Br Cl
d) Na Al
e) O F
f) Mg Ca
6. Define ionization energy.
7. Is it easier to form a positive ion with an element that has a high ionization energy or an
element that has a low ionization energy? Explain.
8. Use the concept of ionization energy to explain why sodium form a 1+ ion (Na+) but
magnesium forms a 2+ ion (Mg2+).
9. What trend in ionization energy do you see as you go down a group/family on the periodic
table? What causes this trend?
10. What trend in ionization energy do you see as you go across a period/row on the periodic
table? What causes this trend?
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11. Circle the atom in each pair that has the greater ionization energy.
a) Li Be
b) Na K
c) Cl Si
d) Ca Ba
e) P Ar
f) Li K
12. Define electronegativity
13. What trend in electronegativity do you see as you go down a group/family on the periodic
table? What causes this trend?
14. What trend in electronegativity do you see as you go across a period/row on the periodic
table? What causes this trend?
15. Circle the atom in each pair that has the greater electronegativity.
a) Ca Ga
b) Li O
c) Cl S
d) Br As
e) Ba Sr
f) O S
16. Define electron affinity.
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17. What trend in electron affinity do you see as you go down a group/family on the periodic
table? What causes this trend?
18. What trend in electron affinity do you see as you go across a period/row on the periodic
table? What causes this trend?
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Practice Questions
1. Which of the following statements are reasons that explain why transition metals are often good metal
catalysts?
I. they have multiple oxidation states so they can lose or gain electrons easily
II. They can form complex ions with ligands that can donate lone pairs of electrons
A. I only
B. II only
C. both I and II
D. neither I nor II
2 All of the following statements about the nitrogen family of elements are true EXCEPT:
A. the atomic radii increase with increasing atomic number.
B. the boiling points increase with increasing atomic number.
C. the electronic configuration of the valence shell of the atom is ns2 np3 where n is the quantum
number.
D. It contains both metals and nonmetals
E. All of the above are true.
3. Which of the following elements is not likely to form either negative or positive ions because its highest
energy level already has a stable octet?
A. K
B. Zn
C. He
D. Xe
E. Sc
4. Which of the following elements has atoms that show the greatest attraction or affinity for electrons?
A. Li
B. S
C. Cl
D. O
Use the following answers for questions 5-8 Match the correct element with the correct description.
A. Fluorine (F # 9)
B. Chlorine (Cl # 17)
C. Bromine (Br # 35)
D. Iodine (I # 53)
E. Astatine (At #85)
5. All isotopes of this element are radioactive.
6. A compound of this element is found in bleach.
7. Atoms of this element have the smallest radius
8. This element is the most reactive in the halogen group. It can displace the others from compounds
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9. The colors of the compounds of the d block elements are due to electron transitions
A. between different d orbitals
B. between d and s orbitals
C. among the attached ligands
D. from the metal to the attached ligands
E. None of the above
10. Which of the following reactions will NOT occur?
A. F2 + 2 NaBr à 2 NaF + Br2
B. Cl2 + 2 KI à 2 KCl + I2
C. Br2 + 2 NaF à 2 NaBr + F 2
D. F2 + 2 NaCl à 2 NaF + Cl2
11. Which of these ions is colorless?
A. [Cr(H2O)6]3+
B. [Fe(CN)6]4C. [Cu(NH3)4]2+
D. [Zn(H2O)4]2+
12. Which of these oxides would you expect to have the highest melting point?
A. Al2O3
B. SO2
C. Cl2O
D. SO3
13. Which of these chlorides would you expect to be the least acidic?
A. NaCl
B. MgCl2
C. AlCl3
D. PCl3
14. The greatest similarity in chemical properties is expected for elements with the atomic numbers
A. 3 and 4
B. 6 and 12
C. 17 and 25
D. 19 and 37
15. Arrange the following neutral gaseous atoms in order of decreasing atomic radius.
A. Cl > F > S > Mg
B. Cl > S > F > Mg
C. Mg > S > Cl > F
D. F > Cl > S > Mg
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16. Which of the following elements most readily gives up an electron?
A. F
B. Cl
C. K
D. Li
17. In the modern periodic table elements are arranged in order of increasing
A. atomic number.
B. atomic mass.
C. number of valence electrons.
D. electronegativity.
18. Barium is an element in group 2 of the periodic table with atomic number 56. Which of the following
statements about strontium is NOT correct?
A. Its first ionization energy is lower than that of calcium .
B. It has two electrons in the outermost energy level.
C. Its atomic radius is smaller than magnesium.
D. It forms a chloride with the formula BaCl2.
19. A molecule or ion that donates a lone pair of electrons while attaching to transition metal ion to form a
complex ion is called
A. a chelating agent
B. a ligand
C. a polyatomic ion
D. none of the above is correct
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Answers
1C
2C
3E
4D
5C
6E
7B
8A
9A
10 A
11 C
12 D
13 A
14 A
15 D
16 C
17 C
18 A
19 C
20 B
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Using a periodic table characterize the following elements according to the group to which they belong
A. Alkali Metal
B. Alkaline earth Metals
C. Halogen
D. Noble Gases
E. Transition Metals
1. Iodine (I # 53)
2. Rubidium (Rb # 37)
3. Astatine (At # 85)
4. Xenon (Xe # 54)
5. Vanadium (V# 23)
6. The first comprehensive periodic table was developed by
A. J.W. Dobereiner
B. Dmitri Mendeleev
C. Vladimir Putin
D. J.A.R. Newlands
E. Lothar Meyer
7. Which of the following elements is not likely to form either negative or positive ions because its
highest energy level already has a stable octet?
A. K
B. Zn
C. He
D. Xe
E. Ra
8. A solution of KSCN will give a dark red color when added to a solution containing
A.NaCl
B. NaI
C. FeSO4
D.FeCl3
E. NaBr
9. Which of the following elements in the third period has the highest ionization energy
A. Na
B. Mg
C. Al
D. P
E. Si
10. The greatest electrical conductivity at room temperature would be observed in
A. Al
B. C
C. P
D. Si
E. S
11. The most reactive group in the Periodic Table is
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A. the transition metals
B. the alkaline Earth Metals
C. the Noble Gases
D. the Alkali Metals
12. The property that decreases from top to bottom in the halogen family is
A. density
B. ionization energy
C. mass
D. boiling point
13. The greatest similarity in chemical properties is expected for elements with the atomic numbers
A. 9 and 10
B. 9 and 18
C. 9 and 17
D. 16 and 17
14. Barium is an element in group 2 of the periodic table with atomic number 56. Which of the following
statements about barium is NOT correct
A. Its first ionization energy is lower than that of calcium .
B. It has two electrons in the outermost energy level
C. Its atomic radius is smaller than calcium
D. It forms a chloride with the formula B BaCl2
15. Which of the following reactions will NOT occur?
A. Br2 + 2 NaCl à 2 NaBr + Cl2
B. Cl2 + 2 NaBr à 2 NaCl + Br2
C. Br2 + 2 NaI à 2 NaBr + I2
D. Cl2 + 2 KI à 2 KCl + I2
16. Which of the following series is arranged in order of increasing value?
A. The radii of: H- ion, H atom, H+ ion
B. The first ionization energies of oxygen, fluorine, neon.
C. The electronegativities of : chlorine, bromine, iodine
D. The boiling points of: iodine, bromine, and chlorine.
17. Which of the following has the largest ionic radius?
A. Al3+
B. Mg2+
C. P3D.
The following list contains oxides of elements in the third period. Use these answers for questions 18-20
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A. SO2
B. P4O10
C. Al2O3
D. MgO
E. Na2O
18. Which of these oxides has the highest melting point?
19. Which of the above is most likely to be a gas at room temperature?.
20. Which of the above would form water solutions that are highly acidic ?
21. Which of the following statements are reasons that explain why transition metals are often good metal
catalysts?
I. they have multiple oxidation states so they can lose or gain electrons easily
II. They can form complex ions with ligands that can donate lone pairs of electrons
A. I only
B. II only
C. both I and II
D. neither I nor II
22. Which of these ions is most likely colorless?
A. [Ni(NH3)4]2+
B. [Cd(H2O)4]2+
C. [Cr(H2O)6]3+
D. [Fe(CN)6]323. An ion or polar molecule that attaches to a simple ion to form a complex ion is called a(n)?
A. catalyst
B. ligand
C. coordinate
D. none of the above
24. The formula for the chloride of a certain element E has the form ECl 3. The element is most likely to be
A. lithium
B. calcium
C. silicon
D. boron
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Answer key
1. C
2. A
3. C
4. D
5. E
6. B
7. D
8. D
9. D
10. A
11. D
12. B
13. C
14. C
15. A
16. B
17. C
18. D
19. A
20. A
21. B
22. B
23. B
24. D
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