Chapter 5: The Periodic Table

Discovered triads.

Determined the combining weights of 43 elements with oxygen.

Determined the atomic weights of the elements.

Arranged elements by increasing atomic weight. Found similar properties occurred every 7 elements (“Law of Octaves”).

Published a table of elements based on the elemental properties and order of atomic weights.

How is Mendeleev’s table different from the modern one?

The properties of the elements are periodic functions (repeating patterns) of their atomic weight

Discovered the noble gases.

Saw that x-rays showed increasing size of nucleus (increasing number of protons from element to element).

The physical and chemical properties of the elements are periodic functions of their atomic numbers.

Approximately what % of elements do metals make up?

What is unique about the actinide series?

How many electrons do halogens have in their valence energy levels? What do halogens like to react with?

Why are noble gases happy?

The 14 elements with atomic numbers from 90-103

The 14 elements with atomic numbers from 58-71

An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group

The elements of Group 1 of the periodic table

The elements of Group 2 of the periodic table

The elements of Group 17

The p-block elements together with the s-block elements

The d-block elements that are metals with typical metallic properties

A negative ion

May be defined as one-half the distance between the nuclei of identical atoms that are bonded together

A positive ion

The energy change that occurs when an electron is acquired by a neutral atom

Measure of the ability of an atom in a chemical compound to attract electrons

An atom or group of atoms that has a positive or negative charge

Any process that results in the formation of an ion

The energy required to remove one electron from a neutral atom of an element

Electrons available to be lost, gained, or shared in the formation of chemical compounds

Which atom is larger: Calcium or Copper?

Which atom is larger: Nitrogen or Arsenic?

If energy is gained, the reaction is ________________. (+EA)

If energy is lost, the reaction is _________________.

Which element has the greatest electronegativity and is the most reactive nonmetal?

Which group of elements has no tendancy to gain electrons?

What are the trends of electronegativity down a group and across a period?

Why does the first energy level in every diagram only contain two electrons?

How many electrons can fit in the second energy level of any atom?

How many electrons can fit in the third energy level?

How many energy levels does aluminum have? How many electrons should be in each energy level?

Explain why the second energy level of aluminum only feels a +11 attraction instead of a +13 attraction from aluminum’s electrons.

How large is the charge that the third energy level of an aluminum atom “feels” from the nucleus?

Which atom is larger: nitrogen or carbon? Why?

In a silicon atom, the force of attraction from the nucleus to the outer energy level is +4. Explain in detail why this is true.

Which atom do you predict to be larger: sulfur or chlorine? Why?

1/2 the distance between two adjacent like atoms

Cations are always ________ than the neutral atom.

What are the trends for ionization energies across a period and down a group?

Why are electrons harder to remove across a period?

Why are electrons easier to remove down a group?

ATOMIC RADIUS: For each of the following sets of atoms, rank the atoms from smallest to largest atomic radius.
a. Li, C, F
b. Li, Na, K
c. Ge, P, O
d. C, N, Al
e. Al, Cl, Ga

IONIZATION ENERGY: For each of the following sets of atoms, rank them from lowest to highest ionization energy.
a. Mg, Si, S
b. Mg, Ca, Ba
c. F, Cl, Br
d. Ba, Cu, Ne
e. Si, P, He

IONIC RADIUS: For each of the following sets of ions, rank them from smallest to largest ionic radius.
a. Mg²⁺, Si⁴⁻, S²⁻
b. Mg²⁺, Ca²⁺, Ba²⁺
c. F⁻, Cl⁻, Br⁻
d. Ba²⁺, Cu²⁺, Zn²⁺
e. Si⁴⁻, P³⁻, O²⁻

ELECTRONEGATIVITY: For each of the following sets of atoms, rank them from lowest to highest electronegativity.
a. Li, C, N
b. C, O, Ne
c. Si, P, O
d. K, Mg, P
e. S, F, He

In the modern periodic table, elements are ordered according to increasing _______________.

Mendeleev noticed that properties of elements appeared at regular intervals when the elements were arranged in order of increasing ____________.

The modern periodic table law states that the _________ and __________ properties of the elements are functions of their ________________.

The discovery of noble gases changed Mendeleev’s periodic table by adding a new _____________.

The most distinctive property of the noble gases is that they are _____________.

Lithium, the first element in Group 1, has an atomic number of 3. The second element in theis group has an atomic number of _________.

An isotope of fluorine has a mass number of 19 and an atomic number of 9.
a. How many protons are in this atom?
b. How many neutrons are in this atom?
c. What is the symbol of this fluorine atom including its mass number and atomic number?

Samarium, Sm, is a member of the lanthanide series.
a. Identify the element just below samarium in the periodic table.
b. The atomic numbers of these two elements differ by how many units?

A certain isotope contains 53 protons, 78 neutrons, and 54 electrons.
a. What is its atomic number?
b. What is the mass of this atom in amus?
c. Is this element Pt, Xe, I, or Bh?
d. Identify two other elements that are in its group.

In a modern periodic table, every element is a member of both a horizontal row and a vertical column. Which one is the group, and which one is the period?

Explain the distinction between atomic mass and atomic number.

In the periodic table, the atomic masses of Te and I decrease rather than increase, while their atomic numbers increase. This phenomenon happens to other neighboring elements in the periodic table. Name two of these sets of elements.

Identify the element, and write the noble-gas notation for each of the following:
a. the Group 14 element in Period 4
b. the only metal in Group 15
c. the transition metal with the smallest atomic mass
d. the alkaline earth metal with the largest atomic number

Give the symbol, period, group, and block for the following:
a. sulfur
b. nickel
c. [Kr] 5s¹
d. [Ne] 3d⁵ 4s¹

There are 18 columns in the periodic table; each has a group number. Give the group numbers that make up each of the following blocks:
a. s-block
b. p-block
c. d-block

When an electron is added to a neutral atom, energy is ________________.

The energy required to remove an electron from an atom is the atom’s _______________.

Moving from left to right across a period on the periodic table:
a. ionization energy values tend to become ______ (larger or smaller)
b. atomic radii tend to become ________ (larger or smaller)

a. Name the Halogen with the least-negative electron affinity.
b. Name the alkali metal with the highest ionization energy.
c. Name the element in Period 3 with the smallest atomic radius.
d. Name the Group 14 element with the largest electronegativity.

Write the electron configuration of the following:
a. Na
b. Na⁺
c. O
d. O²⁻
e. Co²⁺

a. Compare the size of the radius of a positive ion to its neutral atom.
b. Compare the size of the radius of a negative ion to its neutral atom.

a. Give the approximate positions and blocks where metals and nonmetals reside in the periodic table.
b. Of metals and nonmetals, which tend to form positive ions? Which tend to form negative ions?

a. Identify the electron that is removed in the first ionization energy of Mg.
b. Identify the electron that is removed in the second ionization energy.
c. Identify the electron that is removed in its third ionization energy.

Explain why the second ionization energy is higher than the first, the third is higher than the second, and so on.

Explain the role of valence electrons in the formation of chemical compounds.

Consider a neutral atom with 53 protons and 74 neutrons to answer the following questions:
a. What is its atomic number?
b. What is its mass in amus?
c. Is the element’s position in a modern periodic table determined by its atomic number or its atomic mass?

Consider an element whose outermost electron configuration is 3d¹⁰ 4s² 4pⁿ.
a. To which period does the element belong?
b. If it is a halogen, what is the value of n?
c. The group will equal (10+2+n). True or False?

a. Metalloids are found in which block: s, p, d, or f?
b. The hardest, densest metals are found in which block: s, p, d, or f?

a. Name the most chemically active halogen.
b. Write its electron configuration.
c. Write the configuration of the most-stable ion this element makes.

a. Which has the larger radius, Al or In?
b. Which has the larger radius, Se or Ca?
c. Which has a larger radius, Ca or Ca²⁺?
d. Which has greater ionization energies as a class, metals or nonmetals?
e. Which has the greater ionization energy, As or Cl?
f. An element with a large negative electron affinity is most likely to form a positive ion, a negative ion, or a neutral atom?
g. In general, which has a stronger electron attraction, large atoms or small atoms?
h. Which has greater electronegativity, O or Se?
i. In the covalent bond between Se and O, to which atom is the electron pair more closely drawn?
j. How many valence electrons are there in a neutral atom of Se?

Identify all of the ions below that do not have noble gas stability:
K⁺ S²⁻ Ca⁺ I⁻ Al³⁺ Zn²⁺

Give the noble-gas notation for the following:
a. Br
b. Br⁻
c. the element in Group 13, Period 5

Use the position in the periodic table and electron configuration to describe the chemical properties of Calcium and Oxygen.