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Chapter 6 - The Periodic Table
6.1 Organizing the Elements
The same way products in the grocery store are logically grouped according to similar characteristics, elements are in
the periodic table
Chemist use the properties of the elements to sort them into groups. In the early 1800s some elements with similar
properties were grouped into triads. Ex. Cl, Br, I
Russian School Teacher, Dmitri Mendeleev, sorted over 60 elements with their properties written on cards according to
their properties. He arranged the elements into groups based on a set of repeating properties. He also arranged the
elements in the order of increasing atomic mass. This was the first periodic table (published in 1869)
Mendeleev was able to predict the existence and properties of atomic masses 68 and 70 before they were discovered
because he knew that F, Cl, and I must be grouped together. These elements, gallium and germanium, were isolated
and discovered in 1875 and 1886 respectively. Mendeleev's predicted properties were very close to the actual
properties showing the periodic table was a very powerful tool.
The periodic Law
In the modern periodic table, elements are arranged in order of increasing atomic number (Z).
Why? The atomic mass of iodine is 126.90 and the atomic mass of tellurium is 127.6. Based on chemical properties,
iodine belongs in a group with bromine and chlorine. This prompted Mendeleev to break his rule regarding
arrangement by atomic mass, and place Te before I (He assumed the masses were measured incorrectly).
Later, in 1913, all atomic numbers (proton number) were determined for each known element and this verified that
tellurium (Z=52) should come before iodine (Z=53). The masses that Mendeleev assumed to be measured wrong were
actually correct.
There are seven horizontal rows or periods in the periodic table. Each period corresponds to a principle energy level.
The elements in a vertical column or group have similar properties and are given "family" names.
Properties in a period vary from left to right, but this pattern of varying properties repeats from one period to the next.
Periodic Law - When elements are arranged in order of increasing atomic number, there is a periodic repetition of
their physical and chemical properties.
Metals, Nonmetals, and Metalloids - The 3 main classes of elements
Across a period (from left to right), the properties of elements become less metallic and more nonmetallic
1. Metals - most of the elements are metals (about 80%). Metals are good conductors of heat and electricity. They have
luster or sheen due to their ability to reflect light. All metals are solid at room temperature except for mercury. Most
metals are ductile (can be drawn into wire) and malleable (able to be hammered into thin sheets). Due to their
properties, metals have many uses.
2. Nonmetals comprise the upper right hand corner of the periodic table. Most nonmetals are gases at room
temperature (Z = 1, 2, 7, 8, 9, 10, 17, 18, 36, 54, 86); but others are solids (Z= 6, 15, 16, 34, 53) and bromine (a dark red
liquid Z = 35) is a liquid. Nonmetals have variable properties, but typically have properties opposite that of metals. They
are generally poor conductors of heat and electricity except for carbon and nonmetal solids tend to be brittle.
3. Metalloids - border the stair-step line that separates metals from non metals and generally have properties that are
similar to metals and nonmetals. They can behave like metals under certain conditions and nonmetals under certain
conditions. Semi conductors are created from these elements. For instance, mixing B with Si created semiconducting
silicon computer chips and wafers.
Section Review Questions
1. How are elements arranged in the modern periodic table?
2. Name the three broad classes of elements and give 3 examples of each. Note each the state of matter of each
example at STP
3. Which of the following sets of elements have similar physical and chemical properties.
a. oxygen, nitrogen, carbon, boron
b. strontium, magnesium, calcium, beryllium
c. nitrogen, neon, nickel, argon
4. Identify each element below as a metal, metalloid, or nonmetal
a. gold
b. sulfur
c. silicon
d. barium
e. uranium
f. hydrogen
e. arsenic
5. Name two elements that have properties similar to those of the element sodium
6.2 Classifying Elements
The periodic table displays the symbols and names of the elements, along with information about the structure of
their atoms.
Our periodic table displays the same information, but in different areas
Groups of elements - p. 162- 163 - The background colors of the squares distinguish groups of elements
•
The Group 1A (1) elements are called alkali metals.
•
The Group 2A (2) elements are called alkaline earth metals.
•
The nonmetals of Group 7A (17)are called halogens.
•
The nonmetals of Group 8A (18) are called noble gases.
Elements can be sorted into noble gases, representative elements, transition metals, or inner transition metals based
on their electron configurations.
Electron configurations determine the properties of elements, and thus elements in the same groups have similar
electron configurations.
1. The noble gases are the elements in Group 8A (18) of the periodic table. Consider the electron configurations for the
first four noble gases. They all have their highest occupied energy level full with electrons and are generally unreactive
as a result. We call them inert gases because they are not chemically active.
2. The Representative Elements - Elements in groups 1A (1), 2A (2), 3A (13), 4A (14), 5A (15), 6A (16), and 7A (17) are
called representative elements because they display a wide range of physical and chemical properties. In these
elements, the s and p sublevels of the highest occupied energy level are not filled.
Ex. Consider the electron configurations of the first three alkali ( group 1) metals. There is only one electron in the
outermost energy level and the electron is in the s sublevel
Ex. Consider the first three elements in group (4A) 14. They have four electrons in the outermost energy level
*** For any representative element, the group number indicates the number of electrons in the highest energy level***
3. Transition elements - are the B group elements in the middle that connect the A group representative elements.
They are group numbers 3-12. These elements are classified into 2 groups based on their electron configurations.
a. transition metals - the highest occupied s sublevel and a nearby d sublevel contain electrons. These elements
are characterized by the presence of electrons in d orbitals.
b. inner transition metals - the highest occupied s sublevel and a nearby f sublevel contain electrons. These
elements are characterized by the presence of electrons in f orbitals.
Blocks of Elements - The periodic table can be divided into blocks based on element position and electron configurations
The periodic table can be divided into 4 blocks ---> s, p, d, and f. From this version of the table, electron configurations
can be easily determined (d sublevel is filled in 1 lower principal energy level while f sublevels are 2 below)
Section Review
1. Use your periodic table to write the electron configuration and expanded electron configuration of each element.
a. carbon
b. strontium
c. vanadium
2. What information can be included in the periodic table?
3. Into what four classes can elements be sorted based on their electron configurations?(hint: look at the table above)
4. Why do the elements potassium and sodium have similar chemical properties?
5. Classify each element as a representative element, transition metal, or noble gas
a. 1s22s22p63s23p63d104s24p6
b. 1s22s22p63s23p63d64s2
c. 1s22s22p63s23p2
6. Which of the following elements are transition metals? Cu, Sr, Cd, Au, Al, Ge, Co
7. How many electrons are in the highest occupied energy level of the element in period 5, group 14?
6.3 Periodic Trends - This section includes periodic trends in atomic size, ionization energy, ion size, and
electronegativity.
1. Atomic size - is measured when atoms form molecules in terms of atomic radius - is one half the distance between
the nuclei of atoms of the same element when the atoms are joined. Atomic radius is measured in picometers because
atoms are so small. The measurements are found in table S.
In general, atomic size/radius increases from top to bottom within a group and decreases from left to right across a
period
Atomic Size Trends -Notice the data for the alkali metals and the noble gases. The atomic radius trend increases going
down the group, but decrease going across a period. Why? As atomic number increases going across a period, the
nuclear charge increases which pulls the electrons in, however, going down a group, the nuclear charge increases, but so
does the number of principle energy levels. The inner electrons shield the electrons in the highest occupied energy level
from the nuclear charge and atomic size "jumps up". This effect is called shielding.
Ions - an ion is an atom or group of atoms that has a positive or negative charge. Atoms are typically neutral because
they have equal numbers of electrons, but positive and negative ions form when electrons are transferred between
atoms.
Metals tend to form ions by losing one or more electrons. This results in a positively charged ion called a cation. The
charge for a cation is written as a number followed by a + sign as a superscript.
Ex. Na -----------> 1 e-
+
Na1+ (if the charge is 1 we usually drop the number and write Na+)
Mg ------------> 2 e-
+
Mg2+
**When metals lose electrons to form cations, the resulting cation has a smaller radius than the original neutral atom.
Nonmetals tend to form ions by gaining one or more electrons. This results in a negatively charged ion called an
anion. The charge for an anion is written as a number followed by a - sign as a superscript.
Ex. Cl +
1 e- --------> Cl-
Ex. O +
2e- --------> O2-
**When nonmetals gain electrons to form anions, the resulting anion has larger radius than the original neutral atom.
Ionization energy - The energy required to remove an electron from an atom (units are kilojoules per mole - kJ/mole)
First ionization Energy - the energy required to remove the first electron from an atom (+1 charge)- Table S
Second ionization Energy - the energy required to remove the second electron from an atom (usually much
higher, +2 charge)
Third ionization Energy - the energy required to remove the third electron from an atom (much, much higher,
+3)
* Ionization energies help us predict which ion is most likely to form
Ionization energy trend - First ionization energy tends to decrease from top to bottom within a group and increase
from left to right across a period.
Notice the noble gases and the alkali metals - as an atom gets bigger, the nucleus has a smaller affect on the outermost
electrons (ionization E goes down). Going across the period, nuclear charge increases while the shielding effect remains
constant, so there is more pull on the electrons toward the positive charge (ionization E goes up).
Trends in Ionic Size - During reactions between metals and nonmetals, metal atoms tend to lose electrons, and
nonmetal atoms tend to gain electrons. This transfer has a predictable effect on the size of the ions that form.
Cations are always smaller than the atoms from which they form. Anions are always larger than the atoms from which
they form.
Cations are smaller in size because when the electron(s) are lost the attraction between the remaining electrons and the
nucleus is larger.
Anions are larger in size because as electrons are added, the attraction of the nucleus for any one electron decreases.
Trends in Electronegativity
Electronegativity - is the ability of an atom of an element to attract electrons when the atom is in a compound
In general, electronegativity values decrease from top to bottom within a group. For representative elements, the
values tend to increase from left to right across a period.
The most electronegative elements are fluorine, then oxygen, then chlorine, and nitrogen.
Section Review Questions
1. How does atomic size change within groups and across periods?
2. When do ions form?
3. What happens to first ionization energy within groups and across periods?
4. Compare the size of ions to the size of the atoms from which they form? (hint: compare metals to nonmetals)
5. How does electronegativity vary within groups and across periods?
6. Describe the concept of shielding?
7. Arrange these elements in order of decreasing atomic size: sulfur, chlorine, aluminum, and sodium. Does this
arrangement represent a periodic trend or a group trend?
8. Which element in each pair has the largest first ionization energy
a. sodium, potassium
b. magnesium, phosphorus
c. fluorine, chlorine