Download Unit Packet 3: Periodic Properties

Chemistry
Unit 3
Periodic Properties
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Name: _________________
Name_________________________
Anatomy of the Periodic Table
1. Notice that there are 18 columns (vertical) numbered on the periodic table. These columns
are referred to as groups.
2. Notice that there are 7 rows (horizontal). These rows are referred to as periods. The number
of energy levels in the atom denotes the number of the period.
3. Label the following families on your periodic table.
Noble Gases - group 18
Transition Elements - group 3 - 12
Alkali Metals - group 1
Alkaline Earth Metals - group 2
Lanthanide Elements - Atomic numbers 58 - 71 and are part of the transition elements.
Actinide Elements - Atomic numbers 90-103 and are part of the transition elements.
Halogens - group 17
Hydrogen - forms its own chemical family
4. Label the following regions on your periodic table.




Draw a line under the following elements to create a stair step line: Boron, Silicon,
Arsenic, and Tellurium. All the elements above this line, excluding the Noble gases, are
considered nonmetals.
Metals - Groups 1-12 and all elements below the stair step line excluding Hydrogen
Metalloids - Those elements above and below the stair-step line (6 elements) B, Si, Ge,
As, Sb, Te
Noble Gases - Group 18
5. Create a legend and somehow denote whether each element is a gas or liquid at room
temperature. Elements at room temperature exhibit the following states of matter:
Gas - Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, and group 18
Liquid -Bromine, Mercury
Solid -All the rest of the elements
6. Include in your legend the electronic configuration and somehow denote the electronic
configuration of the following blocks of elements.
Groups 1-2 and Helium
Groups 13-17
Groups 3-12
Lanthanide and Actinide
s-block
p-block
d-block
f block
7. Notice that your table does not include the atomic mass for the elements. Using a periodic
table that includes atomic mass, write them into your table. Round the values to the 10th
decimal place. Do you notice a pattern? Do you notice any inconsistencies?
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Name_________________________
Anatomy of the Periodic Table
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Name_______________________
Effective Nuclear Charge
When moving across a period (left to right), each subsequent atom gains one more proton
and one more electron to the same energy level. As the protons increase, the positive charge of
the nucleus increases causing a greater pull on the electrons. Therefore, as electrons are added
within a period, they are pulled closer to the nucleus and the atomic size decreases. This trend is
less pronounced in periods where there are many electrons between the nucleus and the outer
most energy level.
Note that each electron is simultaneously attracted by the nucleus and repelled by other
electrons. Because there is so much going on within this atom, we will estimate the energy of
each electron by examining how it interacts with its average environment created by the nucleus
and other electrons. Any electrons found between the outer electron and the nucleus will reduce
the nuclear charge that the outer electron “feels.”
This overall positive charge that the outer electron “feels” is called the effective nuclear
charge (Z*). We can estimate this using the following equation:
Z* = Z – S
Z* is the effective nuclear charge
Z is the number of protons in the nucleus
S is the number of electrons that are between
the outer electron and the nucleus. These are
called the shielding electrons.
For example, consider a lithium atom
Li
1s2 2s1
The outer electron is the 2s electron.
Both of the 1s electrons are shielding electrons
3+
Z* = 3 – 2 = 1
To the outer electron in a lithium atom it seems that there is only a +1 charge attracting it to the
nucleus.
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1. Give the number of shielding electrons that exist for the circled electrons in the following
electron configurations.
a. Al
1s2 2s2 2p6 3s2 3p1
b. Ca
1s2 2s2 2p6 3s2 3p6 4s2
c. Ga
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1
d. Na
1s2 2s2 2p6 3s1
e. I
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5
2. Now calculate the effective nuclear charge (Z*) for the circled electrons in the atoms from
question 1.
a.
b.
c.
d.
e.
3. Explain why the effective nuclear charge experienced by a 1s electron in potassium is greater
than that for a 1s electron in lithium.
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Name _____________________________
Atomic Size
We will be examining how the properties of atoms change depending on their location within the
periodic table. You should be able to explain why these properties change based primarily on
the what the electrons are doing (their electronic configuration).
Give the electron configurations and effective nuclear charges (on an outer electron) for the
following atoms. This will illustrate how these change as you move down a column or across a
row in the periodic table.
moving down a column
element
electron configuration
Z*
H
Li
Na
K
moving across a row
element
electron configuration
Z*
Na
Mg
Al
Si
Atomic size
The size of an atom depends on two things: the number of energy levels that contain electrons
and Z*. Describe how you think each of these factors will affect atomic size.
# of energy levels - More energy levels will make the size of the atom.....
Z* - A higher effective nuclear charge will make the size of the atom.....
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Describe how each of these factors changes as you move down a column or across a row in the
periodic table. (Use the terms increases, decreases, or stays the same.)
# of energy levels
Z*
moving down
a column
Using the information above, fill in the blanks in the following rule which describe how atomic
size changes within the periodic table.
As you move down a column in the periodic table, atomic size _______________
# of energy levels
Z*
moving across
a row
As you move across a row in the periodic table, atomic size _______________
Circle the larger atom in each of the following pairs.
a. Mg or Ba
b. Li or B
c. S or Sn
Put the following in order of increasing size.
a. S, Cl, and F
b. Br, Ca, Cl, and Ba
c. Si, O, Pb, S, and Cs
What is the largest atom in the periodic table?
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Name _____________________________
Ionization Energy
We have discussed the fact that it takes energy to move an electron to a higher energy orbital. If
we add enough energy we can get the electron to completely leave the atom. This energy is
called the ionization energy.
The amount of energy required to remove an electron is related to the force that holds that
electron in place. The greater that force is the more energy it will take to remove the electron.
Remember that F 
Z*
d2
distance
3+
Z*
In other words the force holding the
electron in place depends on the
effective nuclear charge and the
distance from the nucleus.
force
The symbol  means "is proportional to".
Describe how each of these factors will affect the ionization energy of an atom.
Z* - A higher Z* will do what to the force holding the electron in place?
atomic size (d) - A larger atom (outer electron is farther away) will do what to the force
holding the electron in place?
Describe how each of these factors changes as you move down a column or across a row in the
periodic table. (Use the terms increases, decreases, or stays the same.)
Z*
atomic size (d)
moving down
a column
As you move down a column in the table, ionization energy will ______________ because
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Z*
atomic size (d)
moving across
a row
As you move across a row in the table, ionization energy will ______________ because
Circle the atom with the larger ionization energy.
a. Na or Si
b. Ca or Mg
c. F or P
Put the following in order of increasing ionization energy.
a. I, Sn, and F
b. Si, C, Ne, and K
c. Sn, Br, Ba, I, and Pb
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Name ___________________________
Periodic Properties
When answering these questions, be as precise and thorough as you possibly can. Answer them
as though you are explaining the answer to someone in the class that does not understand.
1. Why is an oxygen atom smaller than a lithium atom even though it has more protons,
neutrons, and electrons?
2. Why is it harder to remove an electron from a fluorine atom than a bromine atom?
3. Why is it easier to remove an electron from a barium atom than a carbon atom?
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In this section we will look at how the size of an ion compares to the size of the neutral atom
from which it was formed.
The size of an ion depends on the same two factors as the atomic size: Z* and the number of
energy levels which contain electrons.
Give the electron configurations and find Z* on an outer electron for the following atoms and
ions.
element
or ion
electron configuration
Z*
Na
Na+
Ca
Ca2+
F
FS
S2For each of the atom-ion pairs above, put an X by the one that would be larger.
When an atom loses an electron to become a positive ion, is the ion smaller or larger?
Explain why.
When an atom gains an electron to become a negative ion, the ion is larger. Try to explain why.
Compare the electron configuration of each neutral atom above to that of the ion it forms. Come
up with a rule that describes how to predict the type of ion an atom will form.
Oxygen has the electron configuration 1s2 2s2 2p4. What will the charge on an oxygen ion be?
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Name ______________________________
More Periodic Properties
First group (column) elements:
sodium (Na)
1s2 2s2 2p6 3s1
potassium (K)
1s2 2s2 2p6 3s2 3p6 4s1
Second group (column) elements
magnesium (Mg) 1s2 2s2 2p6 3s2
calcium (Ca)
1s2 2s2 2p6 3s2 3p6 4s2
1. Why do elements in the first column of the periodic table form +1 ions?
2. Why do elements in the second column of the periodic table form +2 ions?
3. Why are metals in the first column more reactive than metals in the second column?
4. How do the ionization energies of the metals in the first column change as you move down the
family?
5. Why are metals that are lower on the periodic table more reactive?
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When two atoms are covalently bonded together, the bond is made of valence electrons which
can be shared by both of the atoms. These electrons no longer belong to just one of the atoms.
Electronegativity is defined as the ability of an atom to attract these shared electrons to itself and
away from the other atom. For example, when hydrogen and fluorine form a compound they
share two electrons.
These two electrons are being shared by
hydrogen and fluorine. Whichever atom
has the higher electronegativity will draw
the electrons in closer to itself and away
from the other atom.
F
H
Fill out the summary below for electronegativity.
Property: Electronegativity
Factor(s) which affect electronegativity and how they affect it:
How the electronegativity changes:
Going down a column why it changes this way -
Going across a row why it changes this way -
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Name _____________________________
Moles
1. Using the average atomic masses from the periodic table (in g/mol), calculate how many
moles are present in each of the following samples.
a. 2.0 x 10-3 g of strontium
b. 62.8 g of barium
c. 1.20 tons of iron
2. Using the average atomic masses from the periodic table (in g/mol), calculate the masses of
each of the following samples.
a. 5.0 mol of potassium
b. 0.000305 mol of mercury
c. 4.67 x 10-5 mol of radon
3. Find the number of atoms present in each of the following samples.
a. 0.0908 g of calcium
b. 4.980 x 10-5 mol of carbon
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4. How many grams would exactly one billion atoms of copper weigh?
5. A single sheet of paper has a thickness of 0.10 mm. How high would stack of paper that
contains a mole of paper be in miles?
Distance to the moon = 348,400 km
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Name _____________________________
Moles 2
1. Calculate the number of moles present in each of the following samples.
a. 37.5 g of aluminum
b. 0.0045 g of sulfur
c. 8.057 x 10-5 g of gold
d. 3.98x105 kg of nickel
e. 2.5 x 1016 atoms of phosphorus
f. 429 g of uranium
2. Calculate the mass of each of the following samples.
a. 0.00945 mol of neon
b. 4.05 mol of magnesium
c. 6.095 x 10-5 mol of carbon
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d. 0.0000386 mol of lead
e. 7.056 x 1028 atoms of zinc
f. 2.54 mol of bromine
3. Normally occurring lithium consists of a mixture of two isotopes. Use the following data
taken from a mass spectrometer to find the atomic mass of lithium.
Isotope
6
Li
7
Li
Mass (g/mol)
6.015
7.016
% found in nature
7.42
92.58
4. Normally occurring zinc consists of a mixture of five isotopes. Use the following data taken
from a mass spectrometer to find the atomic mass of zinc.
Isotope
64
Zn
66
Zn
67
Zn
68
Zn
70
Zn
Mass (g/mol)
63.929
65.926
66.927
67.925
69.925
% found in nature
48.89
27.81
4.11
18.57
0.62
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