Topic 6 Topic 6 Topic 6: Periodic Table Table of Contents Topic 6 Topic 6 Basic Concepts Additional Concepts.

Topic 6Topic 6

Topic 6: Periodic Table

Table of ContentsTable of ContentsTopic 6Topic 6

Basic Concepts

Additional Concepts

• By 1860, scientists had already discovered 60 elements and determined their atomic masses.

The Search for a Periodic Table

Periodic Table: Basic ConceptsPeriodic Table: Basic Concepts Periodic Table: Basic ConceptsPeriodic Table: Basic Concepts

• They noticed that some elements had similar properties.

• They gave each group of similar elements a name. Copper, silver, and gold were called the coinage metals; lithium, sodium, and potassium were known as the alkali metals; chlorine, bromine, and iodine were called the halogens.

Topic 6Topic 6

• Chemists also saw differences among the groups of elements and between individual elements.

The Search for a Periodic Table

• They wanted to organize the elements into a system that would show similarities while acknowledging differences.

• It was logical to use atomic mass as the basis for these early attempts.

Periodic Table: Basic ConceptsPeriodic Table: Basic Concepts Periodic Table: Basic ConceptsPeriodic Table: Basic Concepts Topic 6Topic 6

Döbereiner’s Triads

• The elements in a triad had similar chemical properties, and their physical properties varied in an orderly way according to their atomic masses.


• In 1829, the German chemist J.W. Döbereiner classified some elements

into groups of three, which he called triads.

• Döbereiner’s triads were useful because they grouped elements with similar properties and revealed an orderly pattern in some of their physical and chemical properties.

Döbereiner’s Triads

• The concept of triads suggested that the properties of an element are related to its atomic mass.


• Triads show a relationship among the densities that is true for many triads. Density increases with increasing atomic mass.

• The Russian chemist, Dmitri Mendeleev, was a professor of chemistry at the University of St. Petersburg when he developed a periodic table of elements.

Mendeleev’s Periodic Table

• Mendeleev was studying the properties of the elements and realized that the chemical and physical properties of the elements repeated in an orderly way when he organized the elements according to increasing atomic mass.


• Mendeleev later developed an improved version of his table with the elements arranged in horizontal rows.


• This arrangement was the forerunner of today’s periodic table.

• Patterns of changing properties repeated for the elements across the horizontal rows.

• Elements in vertical columns showed similar properties.


• Mendeleev’s insight was a significant contribution to the development of chemistry.


• He showed that the properties of the elements repeat in an orderly way from row to row of the table.

• This repeated pattern is an example of periodicity in the properties of elements.

• Periodicity is the tendency to recur at regular intervals.


• One of the tests of a scientific theory is the ability to use it to make successful predictions.


• Mendeleev correctly predicted the properties of several undiscovered elements.

• In order to group elements with similar properties in the same columns, Mendeleev had to leave some blank spaces in his table.

• He suggested that these spaces represented undiscovered elements.


• Mendeleev was so confident of the periodicity of the elements that he placed some elements in groups with others of similar properties even though arranging them strictly by atomic mass would have resulted in a different arrangement.





Periodic Table of the Elements


The Modern Periodic Table

• There are several places in the modern table where an element of higher atomic mass comes before one of lower atomic mass.

• This is because the basis for ordering the elements in the table is the atomic number, not atomic mass.


The Modern Periodic Table• The atomic number of an element is equal to

the number of protons in the nucleus. • Atomic number

increases by one as you move from element to element across a row.

• Each row (except the first) begins with a metal and ends with a noble gas.



• In between, the properties of the elements change in an orderly progression from left to right.

• The pattern in properties repeats after column 18.

• This regular cycle illustrates periodicity in the properties of the elements.



• The statement that the physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of increasing atomic number is known as the periodic law.


Relationship of the Periodic Table to Atomic Structure

• In the modern periodic table, elements are arranged according to atomic number.

• The atomic number tells the number of electrons it has.



• The lineup starts with hydrogen, which has one electron.

• Helium comes next in the first horizontal row because helium has two electrons. Lithium has three.


• If elements are ordered in the periodic table by atomic number, then they are also ordered according to the number of electrons they have.


• Notice on the periodic table that lithium starts a new period, or horizontal row, in the table.

• Why does this happen? Why does the first period have only two elements?

• Only two electrons can occupy the first energy level in an atom. The third electron in lithium must be at a higher energy level.



• Lithium starts a new period at the far left in the table and becomes the first element in a group.

• A group, sometimes also called a family, consists of the elements in a vertical column.



• Groups are numbered from left to right.

• Lithium is the first element in Group 1 and in Period 2. Check this location on the periodic table.





• Elements with atomic numbers 4 through 10 follow lithium and fill the second period.

• Each has one more electron than the element that preceded it.

• Neon, with atomic number 10, is at the end of the period.



• Eight electrons are added to Period 2 from lithium to neon, so eight electrons must be the number that can occupy the second energy level.





• The next element, sodium, atomic number 11, begins Period 3.

• Sodium’s 11th electron is in the third energy level.

• The third period repeats the pattern of the second period. Each element has one more electron than its neighbor to the left, and those electrons are in the third energy level.




Atomic Structure of Elements Within a Period

• The first period is complete with two elements, hydrogen and helium.

• Hydrogen has one electron in its outermost energy level, so it has one valence electron.



• Therefore, Group 1 elements have one valence electron.

• These elements have one electron at a higher energy level than the noble gas of the preceding period.

• Every period after the first starts with a Group 1 element.



• As you move from one element to the next across Periods 2 and 3, the number of valence electrons increases by one.

• Group 18 elements have the maximum number of eight valence electrons in their outermost energy level.





• Group 18 elements are called noble gases.

• The noble gases, with a full complement of valence electrons, are generally unreactive.



• The period number of an element is the same as the number of its outermost energy level, so the valence electrons of an element in the second period, for example, are in the second energy level.

• A Period 3 element such as aluminum has its valence electrons in the third energy level.




Atomic Structure of Elements Within a Group

• The number of valence electrons changes from one to eight as you move from left to right across a period; when you reach Group 18, the pattern repeats.

• For the main group elements, the group number is related to the number of valence electrons.

• The main group elements are those in Groups 1, 2, 13, 14, 15, 16, 17, and 18.



• For elements in Groups 1 and 2, the group number equals the number of valence electrons.

• For elements in Groups 13, 14, 15, 16, 17, and 18, the second digit in the group number is equal to the number of valence electrons.





• Because elements in the same group have the same number of valence electrons, they have similar properties.

• Sodium is in Group 1 because it has one valence electron.

• Because other elements in Group 1 also have one valence electron, they have similar chemical properties.



• Chlorine is in Group 17 and has seven valence electrons.

• All the other elements in Group 17 also have seven valence electrons and, as a result, they have similar chemical properties.

• Throughout the periodic table, elements in the same group have similar chemical properties because the have the same number of valence electrons.



• Four groups have commonly used names: the alkali metals in Group 1, the alkaline earth metals in Group 2, the halogens in Group 17, and the noble gases in Group 18.



• The word halogen is from the Greek words for “salt former” so named because the compounds that halogens form with metals are saltlike.

• The elements in Group 18 are called noble gases because they are much less reactive than most of the other elements.



• Because the periodic table relates group and period numbers to valence electrons, it’s useful in predicting atomic structure and, therefore, chemical properties.



• For example, oxygen, in Group 16 and Period 2, has six valence electrons (the same as the second digit in the group number), and these electrons are in the second energy level (because oxygen is in the second period).

• Oxygen has the same number of valence electrons as all the other elements in Group 16 and, therefore, similar chemical properties.


Electrons in Energy Levels—Group 16


Physical States and Classes of the Elements

• The color coding in the periodic table on pages 92 and 93 identifies which elements are metals (blue), nonmetals (yellow), and metalloids (green).



• Nonmetals occupy the upper-right-hand corner.

• Metalloids are located along the boundary between metals and nonmetals.

• The majority of the elements are metals. They occupy the entire left side and center of the periodic table.



• Elements are classified as metals, metalloids, or nonmetals on the basis of their physical and chemical properties.

• Each of these classes has characteristic chemical and physical properties, so by knowing whether an element is a metal, nonmetal, or metalloid, you can make predictions about its behavior.


Metals

• With the exception of tin, lead, and bismuth, metals have one, two, or three valence electrons.

• Metals are elements that have luster, conduct heat and electricity, and usually bend without breaking.


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Metals

• All metals except mercury are solids at room temperature; in fact, most have extremely high melting points.

• The periodic table shows that most of the metals (coded blue) are not main group elements.


Metals

• The elements in Groups 3 through 12 of the periodic table are called the transition elements.

• All transition elements are metals.


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Metals

• Some are less common but still important, such as titanium (Ti), manganese (Mn), and platinum (Pt).

• Some period 7 transition elements are synthetic and radioactive.


• Many are commonplace, including chromium (Cr), iron (Fe), nickel (Ni), copper (Cu), zinc (Zn), silver (Ag), and gold (Au).

Metals

• In the periodic table, two series of elements, atomic numbers 58-71 and 90-103, are placed below the main body of the table.

• These elements are separated from the main table because putting them in their proper position would make the table very wide.

• The elements in these two series are known as the inner transition elements.


Metals• The first series of inner transition elements is

called the lanthanides because they follow element number 57, lanthanum.

• The lanthanides consist of the 14 elements from number 58 (cerium, Ce) to number 71 (lutetium, Lu).

• Because of their natural abundance on Earth is less than 0.01 percent, the lanthanides are sometimes called the rare earth elements.

• All of the lanthanides have similar properties.


Metals

• The second series of inner transition elements, the actinides, have atomic numbers ranging from 90 (thorium, Th) to 103 (lawrencium, Lr).

• All of the actinides are radioactive, and none beyond uranium (92) occur in nature.

• Like the transition elements, the chemistry of the lanthanides and actinides is unpredictable because of their complex atomic structures.


Nonmetals

• Although the majority of the elements in the periodic table are metals, many nonmetals are abundant in nature

• The nonmetals oxygen and nitrogen make up 99 percent of Earth’s atmosphere.


Nonmetals

• Carbon, another nonmetal, is found in more compounds than all the other elements combined.

• The many compounds of carbon, nitrogen, and oxygen are important in a wide variety of applications.


Nonmetals• Most nonmetals don’t conduct electricity,

are much poorer conductors of heat than metals, and are brittle when solid.

• Many are gases at room temperature; those that are solids lack the luster of metals.

• Their melting points tend to be lower than those of metals.

• With the exception of carbon, nonmetals have five, six, seven, or eight valence electrons.


Properties of Metals and Nonmetals


Metalloids• Metalloids have some chemical and physical

properties of metals and other properties of nonmetals.

• In the periodic table, the metalloids lie along the border between metals and nonmetals.


MetalloidsTopic 6Topic 6

• Silicon (Si) is probably the most well-known metalloid.

• Some metalloids such as silicon, germanium (Ge), and arsenic (As) are semiconductors.

Metalloids

• The ability of a semiconductor to conduct an electrical current can be increased by adding a small amount of certain other elements.

• Silicon’s semiconducting properties made the computer revolution possible.


• A semiconductor is an element that does not conduct electricity as well as a metal, but does conduct slightly better than a nonmetal.

Semiconductors and Their Uses

• Your television, computer, handheld electronic games, and calculator are electrical devices that depend on silicon semiconductors.

• All have miniature electrical circuits that use silicon’s properties as a semiconductor.


Semiconductors and Their Uses

• You learned that metals generally are good conductors of electricity, nonmetals are poor conductors, and semiconductors fall in between the two extremes.


Basic Assessment QuestionsBasic Assessment Questions

Question 1

Match each element in Column A with the best matching description in Column B. Each Column A element may match more than one description from Column B.

Topic 6Topic 6


Question 1

1. strontium2. chromium3. iodine

Column A Column Ba. halogenb. alkaline earth metal

c. representative elementd. transition element

Topic 6Topic 6


Answers

1. strontium

2. chromium

3. iodine

b, c

d

a, c

Topic 6Topic 6


How many valence electrons are in an atom of each of the following elements?

Question 2

a. magnesium

b. selenium

c. tin

Topic 6Topic 6


Answers

a. magnesium

b. selenium

c. tin

2

6

4

Topic 6Topic 6

Periodic Table: Additional ConceptsPeriodic Table: Additional ConceptsPeriodic Table: Additional ConceptsPeriodic Table: Additional ConceptsTopic 6Topic 6

Additional Concepts

Periodic Table: Additional ConceptsPeriodic Table: Additional ConceptsPeriodic Table: Additional ConceptsPeriodic Table: Additional Concepts

Periodic Trends

• The electron structure of an atom determines many of its chemical and physical properties.

• Because the periodic table reflects the electron configurations of the elements, the table also reveals trends in the elements’ chemical and physical properties.

Topic 6Topic 6

Atomic radius

• The atomic radius is a measure of the size of an atom.

• The larger the radius, the larger is the atom.


Atomic radius

• Research shows that atoms tend to decrease in size across a period because the nuclei are increasing in positive charge while electrons are being added to sublevels that are very close in energy.

• As a result, the increased nuclear charge pulls the outermost electrons closer to the nucleus, making the atom smaller.


Atomic radius

• Moving down through a group, atomic radii increase.

• Even though the positive charge of the nucleus increases, each successive element has electrons in the next higher energy level.


Atomic radius

• Electrons in these higher energy levels are located farther from the nucleus than those in lower energy levels.

• The increased size of higher energy level outweighs the increased nuclear charge.

• Therefore, the atoms increase in size.


Ionic radius

• When an atom gains or loses one or more electrons, it becomes an ion.

• Because an electron has a negative charge, gaining electrons produces a negatively charged ion, whereas losing electrons produces a positively charged ion.


Ionic radius

• As you might expect, the loss of electrons produces a positive ion with a radius that is smaller than that of the parent atom.

• Conversely, when an atom gains electrons, the resulting negative ion is larger than the parent atom.


Ionic radius

• Practically all of the elements to the left of group 4A of the periodic table commonly form positive ions.

• As with neutral atoms, positive ions become smaller moving across a period and become larger moving down through a group.


Ionic radius

• Most elements to the right of group 4A (with the exception of the noble gases in group 8A) form negative ions.

• These ions, although considerably larger than the positive ions to the left, also decrease in size moving across a period.

• Like the positive ions, the negative ions increase in size moving down through a group.


Additional Assessment QuestionsAdditional Assessment Questions

For each of the following pairs, predict which atom is larger.

Question 1

a. Mg, Sr

b. Sr, Sn

c. Ge, Sn

d. Ge, Br

e. Cr, W

Topic 6Topic 6

Answers

a. Mg, Sr

b. Sr, Sn

c. Ge, Sn

d. Ge, Br

Sr

Sr

Sn

Ge

e. Cr, W W

Topic 6Topic 6


For each of the following pairs, predict which atom or ion is larger.

Question 2

a. Mg, Mg2+

b. S, S2–

c. Ca2+, Ba2+

d. Cl–, I–

e. Na+, Al3+

Topic 6Topic 6


Answers

a. Mg, Mg2+

b. S, S2–

c. Ca2+, Ba2+

d. Cl–, I–

Mg

S2–

Ba2+

I–

e. Na+, Al3+ Na+

Topic 6Topic 6


For each of the following pairs, predict which atom has the higher first ionization energy.

Question 3

a. Mg, Na

b. S, O

c. Ca, Ba

d. Cl, I

e. Na, Al

f. Se, Br

Topic 6Topic 6


Answers

a. Mg, Na

b. S, O

c. Ca, Ba

d. Cl, I

Mg

O

Ca

Cl

e. Na, Al Al

f. Se, Br Br

Topic 6Topic 6


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