Chapter 5: The Periodic Table

Johann Wolfgang Dobereiner
Discovered triads.
Determined the combining weights of 43 elements with oxygen.
We will write a custom essay sample on
Chapter 5: The Periodic Table
or any similar topic only for you
Order now
Determined the atomic weights of the elements.
Arranged elements by increasing atomic weight. Found similar properties occurred every 7 elements (“Law of Octaves”).
Dmitri Mendeleev
Published a table of elements based on the elemental properties and order of atomic weights.
-Ours is ordered by atomic number, his by atomic weight.
-He had no noble gases. (he had a blank space babyyy)
-His was vertical and ours is horizontal.
How is Mendeleev’s table different from the modern one?
Mendeleev’s Periodic Law
The properties of the elements are periodic functions (repeating patterns) of their atomic weight
Sir William Ramsay
Discovered the noble gases.
Henry Gwyn-Jefferys Mosely
Saw that x-rays showed increasing size of nucleus (increasing number of protons from element to element).
Modern Periodic Law
The physical and chemical properties of the elements are periodic functions of their atomic numbers.
Approximately what % of elements do metals make up?
They are all radioactive, and some have been made in labs
What is unique about the actinide series?
7 electrons; metals and elements from Group 1 (Alkali metals)
How many electrons do halogens have in their valence energy levels? What do halogens like to react with?
Using the Bohr description of electron shells, happy atoms have full shells. All noble gases have full outer shells (8 electrons).
Why are noble gases happy?
The 14 elements with atomic numbers from 90-103
The 14 elements with atomic numbers from 58-71
Periodic Table
An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group
Alkali Metals
The elements of Group 1 of the periodic table
Alkaline Earth Metals
The elements of Group 2 of the periodic table
The elements of Group 17
Main-Group Elements
The p-block elements together with the s-block elements
Transition Elements
The d-block elements that are metals with typical metallic properties
A negative ion
Atomic Radius
May be defined as one-half the distance between the nuclei of identical atoms that are bonded together
A positive ion
Electron Affinity
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
Ionization Energy
The energy required to remove one electron from a neutral atom of an element
Valence Electrons
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?
Je ne sais pas 69
Which group of elements has no tendancy to gain electrons?
Down a group→ Decreases
Across a period→ Increases
What are the trends of electronegativity down a group and across a period?
The first energy level has one orbital, and each orbital can hold at most two electrons.
Why does the first energy level in every diagram only contain two electrons?
8 electrons
How many electrons can fit in the second energy level of any atom?
18 electrons
How many electrons can fit in the third energy level?
3 levels; First→2 Second→8 Third→3
How many energy levels does aluminum have? How many electrons should be in each energy level?
The two electrons in the first energy level shield them.
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?
Carbon because the attraction is weaker than the one from Nitrogen.
Which atom is larger: nitrogen or carbon? Why?
The ten electrons in the first two energy levels shield the electrons in the outer level. Since there are 14 electrons in a neutral silicon atom, and ten electrons are shielding, then the outer level has a +4 attraction.
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.
Sulfur because the attraction from the nucleus to the outer shell is weaker.
Which atom do you predict to be larger: sulfur or chlorine? Why?
Covalent Atomic Radius
1/2 the distance between two adjacent like atoms
Cations are always ________ than the neutral atom.
Across a period→Increases
Down a group→Decreases
What are the trends for ionization energies across a period and down a group?
-Elements/Atoms get smaller, and electrons are closer to the nucleus.
-Elements gain electrons and become closer to 8.
Why are electrons harder to remove across a period?
Electrons are farther from the nucleus.
Why are electrons easier to remove down a group?
a. F, C, Li
b. Li, Na, K
c. O, P, Ge
d. N, C, Al
e. Cl, Al, Ga
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
a. Mg, Si, S
b. Ba, Ca, Mg
c. Br, Cl, F
d. Ba, Cu, Ne
e. Si, P, He
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
a. Mg²⁺, S²⁻, Si⁴⁻
b. Mg²⁺, Ca²⁺, Ba²⁺
c. F⁻, Cl⁻, Br⁻
d. Zn²⁺, Cu²⁺, Ba²⁺
e. O²⁻, P³⁻, Si⁴⁻
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⁻
a. Li, C, N
b. Ne, C, O
c. Si, P, O
d. K, Mg, P
e. He, S, F
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
Atomic Number
In the modern periodic table, elements are ordered according to increasing _______________.
Atomic Mass
Mendeleev noticed that properties of elements appeared at regular intervals when the elements were arranged in order of increasing ____________.
Physical; Chemical; Atomic Number
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 _____________.
Largely Unreactive
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 _________.
a. 9
b. 10
c. ¹⁹₉F
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?
a. Plutonium
b. 32
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. 53
b. 131 amus
c. I
d. Br, Cl
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.
The horizontal row is a period, and the vertical columns are groups.
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?
The atomic number is how many protons are in an atom, and the atomic mass is the number of protons plus the number of neutrons.
Explain the distinction between atomic mass and atomic number.
Ar & K
Co & Ni
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.
a. Ge: [Ar] 4s² 3d¹⁰ 4p²
b. Bi: [Xe] 4f¹⁴ 5d¹⁰ 6s² 6p³
c. Sc: [Ar] 4s² 3d¹
d. Ra: [Rn] 7s²
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
a. S; Period 3; Group 16; p-block
b. Ni; Period 4; Group 10; d-block
c. Rb; Period 5; Group 1; s-block
d. Cr; Period 4; Group 6; d-block
Give the symbol, period, group, and block for the following:
a. sulfur
b. nickel
c. [Kr] 5s¹
d. [Ne] 3d⁵ 4s¹
a. 1-2
b. 13-18
c. 3-12
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
Either absorbed or released
When an electron is added to a neutral atom, energy is ________________.
Ionization Energy
The energy required to remove an electron from an atom is the atom’s _______________.
a. larger
b. smaller
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. At
b. Li
c. Ar
d. C
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.
a. 1s² 2s² 2p⁶ 3s¹
b. 1s² 2s² 2p⁶
c. 1s² 2s² 2p⁴
d. 1s² 2s² 2p⁶
e. 1s² 2s² 2p⁶ 3s² 3p6 4s2 3d5
Write the electron configuration of the following:
a. Na
b. Na⁺
c. O
d. O²⁻
e. Co²⁺
a. The positive ion is SMALLER than its neutral atom.
b. The negative ion is LARGER than its neutral atom.
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. Metals are on the left side of the periodic table in the s and d blocks. Nonmetals are on the right side in the p blocks.
b. Metals form positive ions, and nonmetals form negative ions.
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. 3s²
b. 3s¹
c. 2p⁶
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.
The second ionization energy is higher than the first, because the same amount of protons in the nucleus are pulling on a lesser amount of electrons.
Explain why the second ionization energy is higher than the first, the third is higher than the second, and so on.
The valence electrons are the ones that can be lost or gained, and decides the reactivity of the compounds.
Explain the role of valence electrons in the formation of chemical compounds.
a. 53
b. 127 amus
c. atomic number
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?
a. Period 4
b. 5
c. True
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. p
b. d
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. Fluorine
b. 1s² 2s² 2p⁵
c. 1s² 2s² 2p⁶
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. In
b. Ca
c. Ca
d. nonmetals
e. Cl
f. negative ion
g. small
h. O
i. O
j. 6
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?
Ca⁺ & Zn²⁺
Identify all of the ions below that do not have noble gas stability:
K⁺ S²⁻ Ca⁺ I⁻ Al³⁺ Zn²⁺
a. [Ar] 4s² 3d¹⁰ 4p⁵
b. [Ar] 4s² 3d¹⁰ 4p⁶
c. [Kr] 5s² 4d¹⁰ 5p¹
Give the noble-gas notation for the following:
a. Br
b. Br⁻
c. the element in Group 13, Period 5
Calcium is located in Group 2, and has 2 valence electrons. Oxygen is located in Group 16, and has 6 valence electrons. Calcium is a metal, making it very reactive, and Oxygen is a nonmetal, which makes it less reactive.
Use the position in the periodic table and electron configuration to describe the chemical properties of Calcium and Oxygen.

Hi there, would you like to get such a paper? How about receiving a customized one? Check it out