Download Page 8: Review 1

Chapter 6 - Periodic Table
OBJECTIVES
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
List the early attempts of classification of the elements.
Match scientists and their contributions to the development of the P.T.
State the modern periodic law.
Distinguish groups and periods in the P.T.
Define chemical stability using the octet rule.
Use the periodic table to predict electron configurations.
List properties of metals and nonmetals.
Distinguish metals from nonmetals in terms of electronic configuration.
Use the P.T to locate the : alkali metals, alkaline earth metals, chalcogens, halogens, noble gases, transition
elements, lanthanides, and actinides.
State the relationship between the properties of elements and their electronic configuration.
Define periodic, and list reasons why a group has properties related to position in the periodic table.
State the general relationship between size of the atoms with their positions in the P.T.
Given a pair of atoms, ions, or an atom and an ion, select the particle with the larger or smaller radius.
Predict and explain the oxidation number of an element given the element's location in the P.T.
Define first ionization energy.
Predict the relative first ionization energies of two elements given their positions in the P.T.
Distinguish metals from nonmetals in terms of ionization energy.
Define the effective nuclear charge, and list the factors that affect ionization energy, radius, and electron
affinity.
Explain the variations in ionization energies of elements as their atomic numbers increase in a period or family
of the P.T.
Define electron affinity and electronegativity.
Predict trends in electron affinities and electronegativities for elements, based on their location in the P.T.
Define family, or group, and explain what members of a chemical family have in common.
Describe some properties of the elements of the representative groups in the P.T.
State the relationship between the activities of metallic and nonmetallic elements in their relation to their
locations in the P.T.
Given two elements, determine which is more chemically active within the metals and nonmetals.
READING:
Chapter 2, Section 2.3, Introduction to Periodic Table pages 35-38,
Chapter 6, Section 6.7: Periodic Trends and the properties of Aroms, pages 166-172
Notes.
ASSIGNMENTs:
Dittos in the packet.
Textbook, page 176,
page 177
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questions # 54-60 evens;
questions # 64, 74
1
Chapter 6 - Periodic Table
Introduction
1.
How many groups are in the periodic table? ______________ How many periods? ____________
2.
Define atomic number _____________________ What information does it convey? _____________
3.
What is mass number ? _____________________________________________________________
4.
What number, other than atomic number is represented for every element in the periodic table?____
5.
Why is this number a decimal number, rather than a whole number? _________________________
6.
What is an isotope? ______________________________________________________________
7.
How can you use the periodic table to determine the number of electrons in an atom?
_____________
8.
How many elements are in period 1_____________? 2? __________ 3? ____________ 4? _____
9.
In which group are the noble gases found? __________________________________
10.
Why are the noble gases special? _____________________________________________________
11.
In how many groups are the s-sublevels filled? ________ p-sublevels filled? __________________
d-sublevels filled? ________________ f-sublevels filled? ____________________________
12.
How does the number of electrons in each sublevel compare with the number of groups in each block?
_________________________________________________________
13.
What block in the periodic table contains the transition elements? _______
The rare earth elements (lanthanides and actinides)? _____________
14.
What do the elements in a group have in common? _______________________________________
15.
What is an ion? ___________________________________________________________________
16.
How does an atom form an ion?________________________________________________________
17.
How does an atom form a cation?______________________________________________________
18.
How does and atom form an anion?____________________________________________________
19.
Which groups form cations? _________________________________
20.
Which groups form anions? _________________________________
21.
Draw the electron diagram for elements in groups 1, 2, 13, 17, and 18.
22.
What is the charge on the ions formed by elements in groups 1, 2, 17, and 18. ( Indicate + or - as well as
magnitude of the charge)
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2
Chapter 6 - Periodic Table
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3
Chapter 6 - Periodic Table
Periodic Properties:
1.
For each of the following pairs, use the periodic table to select the atom that is larger in radius.
a. Rb, Sr
b. Cl, I
c. Na, Rb
d. Mg, Be
e. S, P
f. Ac, U
g. B, Al
h. Au, Ba
2.
For each of the following pairs of particles, select the particle that is larger in radius.
a. Ca, Ca+2
b. Cl, Clc. As-3, P-3
d. Pb4+ , Pb
e. Mg2+ , Be2+
f. Te2- , Te
g. C, C4h. Ag, Ag1+
3.
Predict the oxidation numbers for the following elements:
a. Al
b. N
c. Cl
d. Zn
g. Na
h. Mn
e. Mg
f. S
4.
Which electrons were gained or lost to complete the outer octet and produce the following ions?
a. K+1
b. O-2
c. Ga+3
d. P3e. Sn4+
f. Br1g. Ca2+
h. Sc3+
5.
Which atom in each of the following pairs would have the lower first ionization energy?
a. N, O
b. Te, Sn
c. Ne, F
d. C, Ge
e. Br, I
f. Mg, Ca
g. I, Sb
h. Al, N
I. F, S
6.
Within a group, does the radii of atoms increase or decrease as the atomic number increases?
7.
Does the radii of atoms within a period increase or decrease as the atomic number increases?
8.
In each of the following pairs of atoms, pick the one that is larger.
a. Mg, Na
b. K, Ca
c. Al, B
e. F, N
f. Ne, Ar
d. Br, Cl
In each of the following pairs of particles, pick the one that is smaller.
a. Fe, Fe3+
b. S2- , S
c. Ac3+ , U3+
e. Mo6+ , Mo
f. As3- , As
d. Br1- , Se2-
9.
10.
11.
Predict the oxidation number of the following elements:
a. Li
b. Be
c. B
d. C
g. F
h. Ar
i. K
Predict the oxidation number of the following elements:
a. Rb
b. Co
c. Pu
d. Bi
e. P
f. O
e. V
f. Ba
12.
In a group, will the ionization energy tend to increase or decrease with increasing atomic number?
13.
In a period, will the ionization energy tend to increase or decrease with increasing atomic number?
14.
Do metals generally have lower ionization energies? Explain?
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Chapter 6 - Periodic Table
15.
Carbon has a first ionization energy of 1086.5 kJ/mol. Predict whether the first ionization energies of
the following elements will be more or less than that of carbon.
a. helium
b. lithium
c. fluorine
d. silicon
16.
As the distance between the nucleus and the outer electrons of an atom increases, will the ionization
energy increase or decrease?
17.
As the shielding effect increases, will the ionization energy increase or decrease?
18.
As the positive charge on an ion increases, will the ionization energy increase or decrease?
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Chapter 6 - Periodic Table
Oxidation Numbers:
1.
2.
Predict the oxidation numbers for the following elements:
a. Al
__________
b. N
__________
c. Cl
__________
d. Zn __________
e. Mg __________
f. S
__________
g. Na __________
h. Mn __________
Which electrons were gained or lost to complete the outer octet and produce the following ions?
a. K+1 ________________
b. O-2 ________________
c. Ga+3 ________________
d. P3- ________________
e. Sn4+ ________________
f. Br1- ________________
g. Ca2+ ________________
h. Sc3+ ________________
First ionization Energy:
1.
Which atom in each of the following pairs would have the lower first ionization energy?
a. N, O
e. Br, I
I. F, S
b. Te, Sn
f. Mg, Ca
c. Ne, F
g. I, Sb
d. C, Ge
h. Al, N
2.
Within a group, does the radii of atoms increase or decrease as the atomic number increases?
3.
Does the radii of atoms within a period increase or decrease as the atomic number increases?
4.
In each of the following pairs of atoms, pick the one that is larger.
a. Mg, Na
e. F, N
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b. K, Ca
f. Ne, Ar
c. Al, B
d. Br, Cl
6
Chapter 6 - Periodic Table
General Trends in the Periodic Table:
1.
Arrange the elements Rb, Te, and I in order of
a. increasing atomic radius _______________________
b. increasing ionization energy ____________________
c. increasing electronaffinity (electonegativity) ____________________
2.
List the following species in order of decreasing radius:
a. K, Ca, Ca+2, Rb ____________________________
b. S, Te-2, Se, Te _____________________________
3.
Name and give the symbol for the element with the characteristics given below.
a. Electron configuration 1s22s22p63s23p3 . _______________________________
b. Lowest ionization energy in Group 17. ____________________________
c. Alkali metal with the largest atomic radius. __________________________
d. Largest ionization energy in the third period. _________________________
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7
Chapter 6 - Periodic Table
Review(1)
1.
In a column , the ionization energy tends to ____ with increasing atomic number.
a. increase
b. decrease
c. remain the same
2.
In a period, the ionization energy tends to ___ with increasing atomic number.
a. increase
b. decrease
c. remain the same
3.
The _____ the electron affinity, the greater the ionization energy.
a. greater
b. lesser
c. more constant
4.
In an experiment designed to measure the first four ionization energies of aluminum, the I.E. would be
the greatest.
a. first
b. second
c. third
d. fourth
5.
___________ generally have the lowest ionization energies.
a. Noble gases
b. Metalloids
c. Nonmetals
d. Metals
6.
The most active ____ have the highest electronegativities.
7.
Carbon has a first ionization energy of 1121.6 kilojoules per mole. Predict whether the first ionization
energies of the following elements will be more or less that that of carbon.
a. helium _______________
c. fluorine _______________
b. lithium _______________
8.
d. silicon
_______________
Fill in the following blanks with the word "increases" or "decreases."
a. As the distance between the nucleus and the outer electrons of an atom increases, the I.E.
__________.
b. As the shielding effect increases, the I.E.
_______________.
c. As the positive charge on an ion increases, the I.E. _______________.
Explain the following statements:
9.
Ionization energy tends to increase with increasing atomic number along any horizontal row.
10.
Ionization energy decreases with increasing atomic number down any vertical column.
11.
Explain why there is a tremendous increase between the fourth and fifth ionization energies of the
element carbon.
12.
Underline the atom in each of the following pairs that has the lower first ionization energy.
a. Li, Na
b. Kr, Rb
c. Cs, Ba
d. Cl, Br
e. F, Ne
f. S, Cl
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8
Chapter 6 - Periodic Table
Properties of Metals and Nonmetals
1.
Compare the following in terms of :
Metals
Nonmetals
Ionization Energy
Electronegativity
Luster
Deformability
Conductivity of Heat
Conductivity of Electricity
Phase at Room Temperature
Ion Formation
Number of Electrons in outermost Energy Level
2.
How do metals and nonmetals differ in terms of electronic configurations?
3.
The most active metals are found in what part of the Periodic Table?
___________________
What are the properties that determine the activity of a metal? __________________________
____________________________________________________________________________
4.
The alkali metals react with water when cold producing ________________________________
5.
The alkaline metals react only with hot water or steam producing ________________________
6.
The most active nonmetals are found in what part of the Periodic Table? What are the factors that
determine the activity of a nonmetal? __________________________________________
___________________________________________________________________________
7.
The most active nonmetal is ________________________
8.
Draw an arrow showing the increased metallic properties in a period and in a group.
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9
Chapter 6 - Periodic Table
Review(2)
1.
a.
b.
c.
d.
2.
How many sublevels are possible for a. n = 6 ________;
3.
Match the quantum number with their proper descriptions:
a.
b.
c.
d.
4.
If n= 4, there may be __________ sublevels.
The p sublevel mat contain ____ pairs of electrons.
Two electrons occupying the same orbital must have opposite _____.
The possible quantum numbers for the 4p1 electron are : n = ________; l = _________;
m = ___________; s = ___________
number of energy level and describes electron cloud size ________
shape of electron cloud __________
direction in space of each orbital ___________
spin of the electron
______________
(1) l
(2) m
(3) n
(4) s
Since the activity of a nonmetal depends upon the ease with which the atom gains electrons, using the periodic
table arrange the following elements in order of increasing activity.
iodine, fluorine, bromine, and chlorine
5.
b. n = 1 __________
___________________________________
Metallic ions are ______ their corresponding atoms.
a. smaller than
b. the same size as
c. bigger than
6.
The tendency to lose electrons ____ as we move down a column.
a. increases
b. remains the same
c. bigger than
7.
______ first proposed the law of octaves.
a. Mendeleev
b. Dobereiner
c. Meyer
d. Newlands
The most stable atoms are those of
a. metals
b. metalloids
c. noble gases
d. nonmetals
8.
9.
The elements in a(n) ________________ are all in the same horizontal line.
10.
A(n) __________usually has five or more electrons in its outer energy level.
11.
In the ___________________ series electrons are added to the 5f sublevel.
12.
The _________ is formed when elements with similar electron configurations are placed in columns in order of
their increasing principal quantum numbers.
13.
Which group in the Periodic Table has the outer configuration:
a. ns2
b. ns2np2
c ns2np5
14.
Give the total number of outermost level electrons (valence electrons) of an element in Group
a. 1
b. 13
c. 14
d. 16
e. 17
15.
In what group of the Periodic Table do all the elements have:
a. 2 valence electrons
b. 5 valence electrons
c. 6 valence electrons?
NOTES
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10
Chapter 6 - Periodic Table
Objective: Describe the origin of the periodic table. State the periodic law.
Origin:
1.
Dobereiner, Johann (1780-1849)
Arranged elements in triads:
Ca
Sr
Ba
2.
40
87.6
137
Av: 88.5
Cl
Br
I
35.5
79.9
127
Av 81.3
John Newland ( 1837-1898)
1863: arranged elements in order of increasing atomic masses.
Noticed: properties repeat every 8th element.
Law of Octaves: Properties of elements repeat every 8th element.
He arranged his elements in 7 groups of 7 elements each.
> Li Be B C N O F Na
> Na Mg Al Si P S Cl K etc.
Concluded: atomic masses are related to chemical properties.
3.
Dimitri Mendeleev ( 1834 -1907)
1869 Periodic Law: Properties of elements are a function of their atomic masses.
Strengths of Mendeleev’s Table:





Carefully planned.
Contained details.
Suggested periods of different lengths.
Elements with similar properties were arranged in horizontal rows.
Left empty boxes, if no element fitted in the spot.
Predicted properties of the undiscovered elements - for the elements that
were unknown in his time.
He predicted the existence of the following elements: ekaboron (Sc)
eka-aluminum (Ga); ekasilicon (Ge); ekamanganeese (Tc);
dvi-manganese (Rh); ekatantalum (Po).
Noticed irregularities in masses, but used properties to arrange the elements:
I 126.9
Te 127.6;
switched their position in periodic table.
Stated the periodic law.
4.
Moseley (1887-1915)
1914: discovered the relationship between the atomic number (# of protons in nucleus) and the place of the
element in the Period Table using x-ray techniques.
MODERN PERIODIC LAW: Properties of elements are a periodic function of their atomic numbers.
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11
Chapter 6 - Periodic Table
Objective: Describe the nature of periods and groups in the Periodic table.
Period:




Horizontal row.
Begins on the left with an active metal, and ends with a noble gas (except 7th period)
Numbered 1 through 7
Period 1 is a short period, contains only two elements.
Group:


Family, vertical columns numbered 1 - 18.
Elements in a group have similar physical and chemical properties.
Every element is a member of both a group and a period.
Periodic Table and Electron Structure:
1.
The electron configuration of the outer energy levels (valence level) determines the chemical properties of the
elements.
2.
Every member of a group has the same electron arrangement in its valence energy level. Helium is an exception: 2
electrons.
3.
The number of the period in which an element is found is the same as the number of the energy level of is valence
electrons.
4.
It is also the same as the number of the occupied energy levels in atoms of the element.
Short Periods: period 1 ( 2 elements), period 2, period 3 (8 elements each)
Long Periods: periods 4 and 5 ( 18 elements)
periods 6, 7 (32 elements each)
Transition elements: any element with an atom that has an incomplete d-sublevel, or that gives rise to a cation or cations
with incomplete d-sublevels.
Rare Earth Elements: elements with incomplete f-sublevels, or cations with incomplete f sublevels. Lanthanides and
Actinoids.
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12
Chapter 6 - Periodic Table
Chapter 6 – Periodic Table Packet
Page 2:
3.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
4.
15.
a. 18 b. 7
a. # of protons
#p +#n
atomic mass
mixture of isotopes
same # of p; different # of n
atomic number
a. 2 b. 8 c. 8 d. 18
18
inert
a. 2 b. 6 c. 10 d. 1
s1-1 s2-2 s2p1 – 13
a. d b. f
same electron configuration in
outermost energy level
charged atom or group of
atoms
gains or loses an electron
loses electrons
gains electrons
groups 1-14
15-17
16.
17.
18.
19.
20.
21.
22. group 1 - +1 group2=+2
group17=-1 group 18=0
Page 3:
1. Periodic Table
2. Atomic mass
3. Periodic law
4. Atomic mass
5. Period
6. Group
7. Family
8. Transition element
9. Lanthanoid series
10. Actinoid series
11. Metal
12. Nonmetal
13. Semimetal
14. Alkali metal
15. alkali earth metal
16. halogen
17. noble gas
Page 4/5:
1. a. Rb b. I c. Rb d. Mg e. P
f. Ac g. Al h. Ba
2. a. Ca b. Cl- c. As-3 d. Pb e.
Mg+2 f. Te-2 g. C-4 h. Ag
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5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
a. +3 b. –3 c. –1 d. +2 e. +2
f. –2, +4, +6 g. +1 h. +2, +7
a.4s b.2p c. 4s24p1 d.3p e.
5s25p2 f.4p g. 4s
h.
4s23d1
a. O b. Sn c. F d.Ge e. I
f.Ca g. Sb h. Al i. S
increase
decrease
a. Na b. K c. Al d. Br e.N f.
Ar
a. Fe+3 b. S c. U+3 d. Br- e.
Mo+6 f. As
a. +1 b. +2 c. +3 d. +4, -4 e.
–3, +3, +5 f. –2 g. –1 h. 0
i.+1
1. +1 b.+2,+3 c. +3, +4, +6
d. +3, +5 e. +2,+4,+5 f. +2
decrease
increase
yes
a. more b. less c. more d. less
decrease
decrease
increase
Page 6: Oxidation numbers
1.a +3
b. –3
c. –1
d. +2
e. +2
f. –2
g. +1 h. +2
2. see Page 4 #4
First Ionization Energy
1. a. N b. Sn c. F
d. Ge e. I
f. Ca
g. Sb h. Al
Page 7: General Trends in the
Periodic Table
1 a. Rb>Te>I b. Rb<Te<I c.
Rb<Te<I
2. a. Rb>K>Ca>Ca+2
b.Te2>Te>Se>S
3. a. P b. I c. Cs d. Ar
Page 8: Review 1
1. b
2. a
3. a
4. d
5. d
6. nonmetal
7.
a. more b. less c. more d.
less
8. a. decreases b. decreases c.
increases
9. answer unreadable
10. no answer here
11. no answer here
12. a. Na b. Rb c. Cs D. Br e. F
f. S
Page 9: Properties of Metals and
Nonmetals
1.
2.
3.
4.
5.
6.
7.
a. high, low b.low, high c.
high, low d. high, low e.
high, low f. high, low g.
solid, gas h. high, high i.
Low, high
Metals never have an valence
electron in the p orbital.
Group 1; the # of valence
electrons
Hydroxide bases
Hydroxide bases
Group 18; the octet # of
electrons
Depends which PT you look at
– some it’s H, and others it’s
Li.
8.
Page 10: Review 2
1. a. 4 b. 6 c. spins d. 4; 1; 1,+1,0; + or - ½
2. 6;1
3. a. n b. l c. ml d. ms
4. I, Br, Cl, F
5. A
6. A
7. D
8. C
9. Period
10. Nonmetal
11. Actinoid
12. Group
13. A. 2 b. 14 c. 17
14. A. 1 b. 3 c. 4 d. 6 e.7
15. A. 2 b. 15 c. 16
13
Chapter 6 - Periodic Table
Atomic mass = 1
Atomic mass = 7
Atomic mass = 9
Atomic mass = 11
clear gas
barely soluble in air
burns in air when ignited
d = 0.0001 g/cm3
melting point = -259 oC
boiling point = -253 oC
does not conduct electricity
forms XCl and X2O;
XCl is acidic in water
soft silvery-white solid
reacts moderately with water
to release H2 gas and form a
basic solution,
tarnishes in air
density = 0.53 g/cm3
melting point = 186 oC
boiling point = 1336 oC
conducts electricity
forms XCl and X2O;
X2O is basic in water
Atomic mass = 14
silvery-white solid,
reacts with hydrochloric acid
to release H2 gas
density = 1.8 g/cm3
melting point = 1278 oC
boiling point = 2970 oC
conducts electricity
forms XCl2 and XO;
XO is basic in water
yellow brown solid
relatively unreactive with
air, or acid
density = 2.3 g/cm3
melting point = 2300 oC
boiling point = 2550 oC
poor conductor of electricity
forms XCl3 and X2O3
Atomic mass = 16
Atomic mass = 19
black solid, relatively
unreactive with water, or
acid
density = 2.2 g/cm3
melting point = 3550 oC
boiling point = 4200 oC
poor conductor of electricity
forms XCl4 and XO2;
XO2 is acidic in water
clear gas, barely soluble in
water
density = 0.0013 g/cm3
melting point = -210 oC
boiling point = -196 oC
does not conduct electricity
forms XCl3 and X2O3;
X2O3 is acidic in water
clear gas, barely soluble in
water
density = 0.0014 g/cm3
melting point = -218 oC
boiling point = -183 oC
does not conduct electricity
forms XCl2 and XO;
pale yellow gas,
reacts violently with water to
release oxygen gas and form
an acidic solution
density = 0.0017 g/cm3
melting point = -223 oC
boiling point = -188 oC
does not conduct electricity
forms XCl and X2O
Atomic mass = 23
Atomic mass = 24
Atomic mass = 27
Atomic mass = 28
soft, silvery solid
reacts vigorously with water
to release H2 gas and form a
basic solution, tarnishes in
air
density = 0.97 g/cm3
melting point = 98 oC
boiling point = 880 oC
conducts electricity
forms XCl and X2O;
X2O is basic in water
silvery-white solid,
reacts with hydrochloric
acid to releasw H2 gas
density = 1.7 g/cm3
melting point = 651 oC
boiling point = 1107 oC
conducts electricity
forms XCl2 and XO;
XO is basic in water
silvery-white solid,
reacts with hydrochloric acid
to release H2 gas
density = 2.7 g/cm3
melting point = 660 oC
boiling point = 2057 oC
conducts electricity
forms XCl3 and X2O3;
gray solid, relatively
unreactive with air, water ,
or acid
density = 2.4 g/cm3
melting point = 1420 oC
boiling point = 2355 oC
poor conductor of electricity
forms XCl4 and XO2
XO2 is acidic in water
Atomic mass = 31
Atomic mass = 32
Atomic mass = 35.5
Atomic mass = 39
red or yellow solid,
yellow ignites spontaneously
in air
density = 2.2 g/cm3(red)
density 1.8 g/cm3 -(yellow)
melting point = 44 oC
boiling point = 280 oC
poor conductor of electricity
forms XCl3 and X2O3;
X2O3 is acidic in water
yellow solid, relatively
unreactive with air, water, or
acid
density = 2.0 g/cm3
melting point = 116 oC
boiling point = 445 oC
does not conduct electricity
forms XCl2 and XO;
XO2 and XO3
are acidic in water
green-yellow gas,
reacts with water to form an
acidic solution
density = 0.0032 g/cm3
melting point = -103 oC
boiling point = -35 oC
does not conduct electricity
forms XCl and X2O;
X2O is acidic
soft-silvery solid,
reacts violently with water to
release hydrogen gas and
form a basic solution,
tarnishes in air
density = 0.87 g/cm3
melting point = 62 oC
boiling point = 760 oC
conducts electricity
forms XCl and X2O,
X2O is basic in water
Atomic mass = 12
A tom Properties Exercise
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Chapter 6 - Periodic Table
Different properties of the elements are related in systematic way to their atomic number. This is known as the
Periodic Law.
In this exercise you will investigate the relationships between the atomic numbers of the first 20 elements (H
through Ca) and the following properties:
1. First Ionization Energy
2. Atomic Radius (covalent radius)
You will do this by making full-page graphs of the two properties (Y-axis) versus atomic number (X-axis). The
property values can be found in your Periodic Table or in the provided Table.
Procedures:
1.
Prepare 2 graphs:
a) Ionization energy versus atomic number (ionization energy on the y-axis, atomic number on the xaxis).
b) Atomic radius versus atomic number
2.
Connect each point with solid lines.
3.
Now, using a different color for each set of elements, connect all of the elements found in
Group 1 of the Periodic Table. Repeat this procedure for the elements in Group 16 and
Group 18.
Questions:
1.
As you go from element 1 to element 20, what is the general overall pattern for?
a. ionization energy?
b. for atomic radius?
2.
A horizontal row on the Periodic Table is called a “period” .
a. List the elements in period 1; in period 2; in period 3; in period 4.
b. Within a period, how do the values change for
(1) Ionization energy
(2) atomic radius?
3.
Within the same column ( 1, 16, 18 ), how do the values for ionization energy and atomic radius change
with increasing atomic number?
4.
Explain why the radius and the first ionization energies of the elements are considered periodic
properties.
5.
State the modern version of the Periodic Law. How does it differ from Mendeleyev’s version?
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