Periodic Table Review

1. The vertical columns on the periodic table are called groups/families.
2. The horizontal rows on the periodic table are called periods.
3. Most of the elements in the periodic table are classified as metals.
4. The elements that touch the zigzag line are classified as metalloids.
5. The elements in the far upper right corner are classified as nonmetals.
6. Elements in the first group have one outer shell electron and are extremely reactive. They are called alkali metals.
7. Elements in the second group have 2 outer shell electrons and are also very reactive. They are called earth alkali metals.
8. Elements in groups 3 through 12 have many useful properties and are called transition metals.
9. Elements in group 17 are known as “salt formers”. They are called halogens.
10. Elements in group 18 are very unreactive. They are said to be “inert”. We call these the noble gasses.
11. The elements at the bottom of the table were pulled out to keep the table from becoming too long. The first period at the bottom is called the lanthanides.
12. The second period at the bottom of the table is called the actinides.

The table is also arranged in vertical columns called “groups” or “families” and horizontal rows called “periods.” Each arrangement is significant. The elements in each vertical column or group have similar properties. Group 1 elements all have one electron in their outer shells (valence electron). This gives them similar properties. Group 2 elements all have 2 electrons (valence electrons)  in their outer shells. This also gives them similar properties.

There are a number of major groups with similar properties. They are as follows:

Hydrogen: This element does not match the properties of any other group so it stands alone. It is placed above group 1 but it is not part of that group. It is a very reactive, colorless, odorless gas at room temperature. (1 valence electron)

Group 1: Alkali Metals – These metals are extremely reactive and are never found in nature in their pure form. They are silver colored and shiny. Their density is extremely low so that they are soft enough to be cut with a knife. (1 valence electron)

Group 2: Alkaline-earth Metals – Slightly less reactive than alkali metals. They are silver colored and more dense than alkali metals. (2 valence electrons)

Groups 3 – 12: Transition Metals – These metals have a moderate range of reactivity and a wide range of properties. In general, they are shiny and good conductors of heat and electricity. They also have higher densities and melting points than groups 1 & 2. (1 or 2 valence electrons)

Lanthanides and Actinides: These are also transition metals that were taken out and placed at the bottom of the table so the table wouldn’t be so wide. The elements in each of these two periods share many properties. The lanthanides (elements 58-71) are shiny and reactive. The actinides (elements 90-103) are all radioactive and are therefore unstable. Elements 95 through 103 do not exist in nature but have been manufactured in the lab.

Group 13: Boron Group – Contains one metalloid and 4 metals. Reactive. Aluminum is in this group. It is also the most abundant metal in the earth’s crust. (3 valence electrons)

Group 14: Carbon Group – Contains on nonmetal, two metalloids, and two metals. Varied reactivity. (4 valence electrons)

Group 15: Nitrogen Group – Contains two nonmetals, two metalloids, and one metal. Varied reactivity. (5 valence electrons)

Group 16: Oxygen Group – Contains three nonmetals, one metalloid, and one metal. Reactive group. (6 valence electrons)

Groups 17: Halogens – All nonmetals. Very reactive. Poor conductors of heat and electricity. Tend to form salts with metals. Ex. NaCl: sodium chloride also known as “table salt”. (7 valence electrons)

Groups 18: Noble Gases – Unreactive nonmetals. All are colorless, odorless gases at room temperature. All found in earth’s atmosphere in small amounts. (8 valence electrons)

CLASSICATION OF ELEMENTS

1.  For these elements, write the noble gas configuration and Lewis Dot diagrams except use the complete electron configuration for H.

H    1s¹

Li   [He] 2s¹

Na  [Ne] 3s¹

K    [Ar] 4s¹

2.  What is the same for each?

                        They have the same number of valence electrons.

Valence electrons are the electrons found in highest energy level  (s and p electrons) The number of valence electrons is the same as the family number in the A group.  (see figure 6-9)

***Atoms in the same group have similar chemical properties because they have the same number of valence electrons.

T/F.  Remember make the false into true.

T    F     3.  Lithium is an alkali metal.

T    F    4.  Iodine is not chemically similar to tellurium.

T    F     5.  The electron configuration for the halogens is s² p⁵.

T    F     6.  The alkali earth metals are more less reactive than the alkali metals.

T    F     7.   Lavoisier Mendelev would be considered the father of the modern periodic table.

T    F     8.   Moseley’s periodic table was arranged by increasing atomic number.

T    F     9.   Sodium is a nonmetal metal.

T    F     10.  Aluminum is a metalloid metal.

T    F     11.  The “stairstep” line divides metals from nonmetals on the periodic table.

Name the element and write the noble gas configuration for the element fitting each of the following descriptions.

12.   the metal in group VA – Bismuth          [Xe] 6s²4f¹⁴5d¹⁰6p³

13.   the halogen in period 3 – Chlorine        [Ne] 3s²3p⁵

14.   the alkali metal in period 2 – Lithium    [He] 2s¹

15.   the transition metal that is a liquid at room temperature – Mercury     [Xe] 6s²4f¹⁴5d¹⁰

16.  FILL IN THE BLANK:  Write “increase(s)” or “decrease(s)” in the blank provided when describing each periodic trend.

PERIODIC TRENDS

Atomic radius – the radius of an atom  (See figure 6-11 page 163).

• ·   The atomic radii decreases as you move left-to-right, because the positive charge increases (adding protons) and the number of energy levels stays the same. The nuclear charge pulls the outermost electrons close to the nucleus.  The positive pull of the protons is called the effective nuclear charge, also called the “Z effect”.
• ·   The atomic radii increases as you move down a group, because the added number of energy levels overcomes the increase in nuclear charge. The electrons on the inner levels are shielding the positively charge nucleus and the electrons spread out even further.

Ionic radius – the radius of an ion (remember that ions have gained or lost electrons).  Figure 6-16, pg 167.

• ·   When atoms lose electrons to form cations, their radii decreases for 2 reasons:
• o The valence electrons are lost leaving an empty orbital
• o The electrostatic repulsion between the now remaining electrons decreases, allowing them to be pulled closer to the nucleus.
• ·   When atoms gain electrons to form anions, their radii increases.
• o The addition of an electron increases the electrostatic repulsion (like charges repel) forcing them farther apart.
• ·   As you move left-to-right across a period the size of the ions gradually decreases. At 5A (15) and 6A (16) the size increases and then decreases again.
• o Elements that form cations are generally in groups 1A-4A  (1,2,13, 14).
• o Elements that form anions are generally in 5A-7A  (15-17).
• ·   As you move down a group, the radii increases due to added energy levels.

Ionization energy – the energy required to remove an electron from an atom. Think of ionization energy as how strong an atom’s nucleus holds onto its electrons.

• ·                           First ionization energy is the energy needed to remove the first electron to form a cation.
• ·                           Second ionization energy is the energy needed to remove the second electron from a +1 ion, if it is possible.
• ·                           Third, fourth, fifth, sixth, seventh, eighth and ninth ionizations are also possibilities.
• o   As you move left-to-right across a period the first ionization energies generally increases because anions are found on the right and anions want to keep their electrons.
• o   As you move down a group, first ionization energies generally decreases because the farther the electron is from the positive force (protons in the nucleus), the easier it is to remove.

Octet rule – Atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons.

• o   The exception is the first period elements, H and He, are complete with only 2 electrons.
• o   The octet rule determine the type of ion likely to form.
• o   Elements on the right side tend to gain electrons to form a noble gas configuration.
• o   Elements on the left side tend to lose electrons to form a noble gas configuration.

Electronegativity – an element’s relative ability of its atoms to attract electrons in a chemical bond.         

• o   The values are based on fluorine, which is 4.0.
• o   Francium has the lowest electronegativity with 0.70.
• o   As you move left-to-right the electronegativity increases because anions are found on the right and those elements tend to gain electrons.
• o   As you move down a group the electronegativity decreases because the farther the electron is from the positive force (protons in the nucleus), the more difficult it is to attract electrons.

REACTIVITY OF METALS AND NON-METALS

Reactivity – the ability of an element to undergo a chemical change – how easy it is for the atom to lose, gain or share electrons.

• ·         Metals are located on the left side of the periodic table. ALL metals lose electrons in chemical reactions
• o    The reactivity of metals decreases as you move left to right on the periodic table, because the atomic radii are decreases making it harder to remove electrons.
• o   The reactivity of metals increases as you move down a family, because the electrons are easier to remove when the atomic radius is larger.
• ·         Non-metals are located on the right side of the periodic table. Non-metals are more likely to gain electrons during a chemical reaction.
• o   The reactivity of non-metals increases as you move left to right on the periodic table, because fewer electrons are needed to complete the outer shell.
• o   The reactivity of non-metals decreases as you move down a family because the atomic radii is increasing shielding the nuclear charge which attracts the electrons.

23.  Which metal is the most reactive?  Why?

      Francium; because it has the lowest ionization energy.

24.   Which non-metal is the most reactive?  Why?

Fluorine; because it has the highest electronegativity.

True/False.  Make the false into true.  Use the trends to answer these questions.  Think about why these trends occur.  Try to answer these questions without looking at the summary below.

T    F    25.  Atoms that attract electrons become positively negatively charged.

T    F    26.  Cl is larger smaller than Cl^-1.

T    F    27.  Li is larger than B.

T    F    28.  P is more less electronegative than N.

T    F    29.  I is more less reactive than Cl.

T    F    30.  Ba is larger than Ba⁺².

T    F    31.  Ar is more less reactive than S.