Chemistry Unit 3 Lesson 1 and 2

What does the d block start with?

3d

What does the f block start with and what are they called?

4f and lanthanides

What is the second row of f and what are they called?

5f and actinides

Transition Elements

A d-block element located in Groups 3 through 12 on the Periodic Table of the Elements; it can have multiple valences, and form colored compounds and stable complex ions

Octet

A full s and p sublevel (eight electrons) that completes the outer shell of an atom and makes the element very stable and chemically nonreactive.

Anion

A negatively charged ion formed when a neutral atom gains electrons

Halogen

A p-block element located in Group 17 on the Periodic Table of the Elements; a halogen is a very reactive nonmetal that combines with metals to form salts

Cation

A positively charged ion formed when a neutral atom loses electrons

Law of Triads

A relationship between atomic mass and chemical properties of elements proposed by Johann Döbereiner, which states that if three elements are arranged in ascending order of their atomic masses, such that the atomic mass of the middle element is the arithmetic mean of the first and third elements, then these elements will show similar properties

By looking at the ionization energies of various elements on the Periodic Table, a chemist can tell at a glance if certain elements will give up electrons to react with other elements.

Alkali metals (group 1) have relatively low first ionization energies (because they are so reactive because they easily give up electrons). But each noble gas has full s and p sublevels completed by valence electrons, so their first ionization energies are very high.

Descriptions of Group 2 Elements

Alkaline-earth metals have many of the same properties as alkali metals, but they are less reactive. Even so, they are still too reactive to be found in nature in their elemental form. Alkaline-earth metals are also more dense and have higher melting temperatures than alkali metals.

Ion

An atom or a group of atoms that has a positive or negative charge due to the loss or gain of electrons

Noble Gases

An element in Group 18 of the Periodic Table of the Elements; any noble gas except helium is a d-block element; a noble gas is chemically stable and has a low reaction rate

Metalloid

An element that is not as malleable as a metal, but not as brittle as a nonmetal, and which will conduct electricity under certain conditions

Alkaline-earth metal

An s-block element located in Group 2 on the Periodic Table of the Elements; a metal that reacts with halogens to form salts

Lanthanide

Any one of a series of metallic f-block elements with atomic numbers 58 through 71

Actinide

Any one of a series of radioactive f-block elements with atomic numbers 90 through 103

Actinides

Are the 14 elements with atomic numbers 90 through 103. Since they follow the element actinium, they are called actinides. Actinides have electrons in the 5f sublevel, but their valence electrons are in the 7s sublevel. Actinides are all radioactive and only four of them—thorium, proactinium, uranium, and neptunium—are naturally occurring. The rest have been synthesized in laboratories.

Lathanides

Atomic numbers 58-71 and are all metals with a silver-gray luster. They are about as reactive as alkaline-earth metals and will tarnish if exposed to air. Most will burn at a relatively low temperature and a few will ignite if they are scratched with a knife.

Ionization Energies

Atoms rarely exist alone. They are usually bonded to other atoms.Atoms form chemical bonds by either sharing electrons or by becoming ions. When an neutral atom loses an electron, it has less negative charge, so it becomes a positively charged ion (Cation). When a neutral atom gains an electron, it becomes a negatively charged ion (anion). The process of gaining or losing electrons to create an ion is called ionization.

Electron Affinities Example

Consider sodium chloride (salt). It is made when sodium and chlorine atoms react.You see, sodium and chlorine have the perfect relationship. If sodium gives up an electron, it achieves a more stable energy state. If chlorine gains an electron, it achieves a more stable energy state. It's a win-win relationship. When a sodium atom meets a chlorine atom, several things happen at the same time. Sodium gives up an electron to become a sodium cation, Na+. Chlorine receives the electron to become a chlorine anion, Cl-. The energy released by the sodium is absorbed by the chlorine, the two ions form a lovely couple (NaCl) and everyone lives happily ever after.

Electronegativies ___ down each group and ___ across each period.

Decrease down each group and increase across each period

Electron affinities ___ down each group and ___ across each period.

Decrease down each group and increase across each period.

Ionization energies ___ down each group and ___ across each period

Decrease down each group and increase across each period. The decrease is because each successive period adds another level of electrons, which places them farther from the nucleus. As a rule, the farther valence electrons are from the nucleus, the less hold the nucleus has on them and the easier it is for them to be given up.The increase is because as the number of protons increases, the force of attraction of the nucleus increases, which makes removing an electron more difficult. However, ionization energies take a bit of a roller coaster ride to get from Group 1 to Group 18.You would think that because boron has more protons, it should have a higher IE1. But its IE1 (801 kJ/mol) is less than that of beryllium. Boron's IE1 is lower because when a boron atom loses one electron, it returns to a more stable configuration with a lower energy state. The IE1 of the next two elements, carbon and nitrogen, increase as expected, but then we take another dive when we get to oxygen. We can see why if we look at the orbital notation for nitrogen and oxygen. Nitrogen has one electron in each of the 2p orbitals (Hund's Rule), which is a stable configuration (lower energy state). Oxygen has two electrons in the 2px orbital and one each in the 2py and 2pz. If oxygen gives up its extra electron in the 2px orbital, it returns to more stable configuration with a lower energy state. By the time we get to Periods 5, 6, and 7, the number of protons has increased so much that this roller coaster effect is not as noticeable.

Glenn Seaborg

Earlier we mentioned that Elements 58 through 71 are called lanthanides and are placed below the main body of the table. This is done for two reasons. All of the lanthanides have very similar chemical properties. As such, they cannot be placed in any one group. Also, the lanthanides have a particular electron configuration which makes them different from elements in any other group. In 1945, American chemist Glenn Seaborg suggested that a new group (with the recently discovered new elements) be added to the Periodic Table. Elements thorium (90) through lawrencium (103) fall between actinium and rutherfordium, so they were named the actinide series and form the actinides.

Valence Electrons

Electrons that occupy the highest energy level of an atom of a particular element and determine its chemical properties

Protons attract -

Electrons. The strength of electrostatic attraction between protons and electrons depends on how many protons are in the nucleus. A nucleus with many protons will hold electrons tighter and the ones closer to the nucleus are more strongly held to the atom.

Descriptions of Groups 13-18 Elements

Elements in Groups 13 through 18 are called p-block elements. They are metals, nonmetals, or metalloids. Nonmetals are poor conductors of heat and electricity and are usually brittle in their solid state. Metalloids are elements with properties of both metals and nonmetals. They are not as malleable as metals but not as brittle as nonmetals and are also called semiconductors because they will conduct electricity under certain conditions.

Descriptions of Group 1 Elements

Except hydrogen are all alkali metals. Metals are good conductors of heat and electricity and have metallic luster. Alkali Metals are soft enough to be cut with a knife, have a relatively low melting point, and are so chemically reactive that they are not found in nature in their elemental form—they are always found combined with nonmetals to form compounds. When they are in contact with water, they react vigorously to form compounds called oxides and releases Hydrogen gas.

How to find out the Group with the Electron Configuration

For the S-Block elements the group number is the same as the number of valence electrons (ns). For the d-blocke elements the grou number is the sum of the electrons in the valence level (ns) and the (n-1) d sublevel. For example Iron's is [Ar]4s²3d⁶ and 6 plus 2 is 8 so group 8. For the p block elements its the same as d plus 10.

Descriptions of Group 17 Elements

Halogens are all very reactive nonmetals and contain elements that exist in all three states of matter at room temperature.

Periods

Horizontal rows

Ionic Radii

If an atom gains electrons, its radius will increase. If an atom loses electrons, its radius will decrease. As a result, cations<original atoms<anions. Metals in Groups 1-15 tend to lose electrons to become cations. As their positive charge increases from losing electrons, their ionic radius decreases. Nonmetals in Groups 14-17 tend to gain electrons to become anions. As they gain electrons to become more negative, their ionic radii increase in size.

Dimitri Mendeléev

In 1860, he attended the Karlsruhe Congress. The next year, he published a book on spectroscopy. In 1865, he began researching the periodicity of the elements. In 1869, he submitted a paper to the Russian Chemical Society in which he outlined his Periodic Table of the Elements.When Mendeléev began working on his table, he used different chemical properties just as Meyer was doing. He arranged them according to different information. When he arranged the cards in order of increasing atomic mass, they did not always fit into groups according to their properties. When that happened, he always placed the card in the group with similar properties even if their atomic masses were out of order. Eventually, he worked out a table that was surprisingly similar to Meyer's table and the one we use today. When Mendeléev completed his table, he noticed there were gaps, just as Meyer had discovered before but many of these were filled in later after they were discovered. If he had arranged the table strictly according to atomic mass, elements with the wrong properties would have gone in those gaps. Mendeléev correctly placed them according to their properties instead of according to their atomic masses. Because Mendeléev published his table before Meyers, it was more complex and because he was able to predict the existence of new elements, he is generally given credit for the discovery of periodicity and the compilation of the Periodic Table of the Elements.

The Cylinder of de Chancourtois

In 1862, a French geologist named Alexandre-Emile Béguyer de Chancourtois [shan-koo-TWAH] noticed a relationship between elements with similar properties and their atomic masses. When he arranged the elements in order of increasing atomic mass, he noticed that elements with similar properties recurred about every 16 amu (atomic mass unit). To demonstrate his theory, de Chancourtois placed the names of the elements in a spiral around a cylinder such that one revolution of the cylinder was equivalent to 16 amu. Many elements with similar properties lined up vertically on the cylinder, but many did not. He presented his findings in a paper on the subject; however, because he tried to apply the concept to geology instead of chemistry and because the paper was so confusing, very few scientists took him seriously. De Chancourtois was the first chemist to publish a paper on the periodicity of elements, but due to bad timing and a poor presentation, the Cylinder of de Chancourtois was not given much recognition.

Newlands' Octaves

In 1863, 62 elements had been discovered. Newlands arranged the elements according to their atomic masses in a table with 8 rows that he called octaves and 7 columns that he called groups. Like de Chancourtois he discovered that certain elements with similar properties lined up periodically. But in Newlands' table, the (similar) elements lined up on every 8th element instead of every 16 amu. In other words, elements 1 and 8 would have properties similar to each other (at least in theory). Since the pattern repeated on every 8th element, he called each row an octave (like the musical scale).It didn't work out for Group 1 but it did work out for part of Groups 2, 3, 4, 5, 6, and 7, but only for octaves A and B. After that, his table did not work.When you compare Newlands' table with the modern Periodic Table, you will notice several things: •Some of his atomic mass values were incorrect and this led to elements being out of order. •Some of the elements were not in columns of similar properties. For instance, he placed iron in the same column with oxygen and sulfur, but their properties are not even close to being similar. Sometimes, he had to place two elements together in the same box to try to make the table work out correctly. •The noble gases were left out because they had not yet been discovered. His table didn't quite work like he theorized it would so scientists of his day didn't take it too seriously.

History of the Noble Gases

In 1894, an English physicist with the title of Lord Rayleigh and a Scottish chemist named William Ramsay discovered a new element that didn't fit anywhere on the Periodic Table. Using fractional distillation, the two researchers were able to isolate the element, which they named argon. Rayleigh and Ramsay continued to search for other elements with similar properties. In 1895, they isolated helium. In 1898, they separated neon, krypton, xenon, and radon from air. They then created group 18 to hold the other elements that had all the elements that didn't go anywhere else.

Pauling

In 1932, he developed a table of electronegativity values. He used ionization energies and electron affinities to calculate electronegativities. An element with a high electronegativity means an atom of that element will attract electrons more strongly than an element with a low electronegativity.

Electron Affinities

In the last section, we saw how neutral atoms can lose electrons to become cations. Now we will look at how neutral atoms can gain electrons to become anions . Energy is required (ionization energy) to release an electron from an atom. The opposite happens—energy is released—when an atom gains an electron. If an atom gains an extra electron, energy is added, but, at the same time, energy must be released to restore the equilibrium of the energy in the atom. If an electron is released from one atom, it must be gained by another atom. Several electron affinity values, such as those for most of the noble gases, are zero. This is because the atoms are so stable that they will not readily accept an electron. Periodic trends are not so evident for electron affinities, but an element's position on the Periodic Table can help a chemist determine the reactivity of an element. The noble gases are nonreactive, but the halogens are very reactive.

Atomic Radii ____ down each group and ___ across each period.

Increase down each group and decrease across each period.

Henry Moseley

It is important to note at this point that chemists such as Mendeléev and Meyer were using Dalton's model of the atom. They didn't even know about electrons and protons. However, they could not explain periodicity and why the elements should be arranged as they are on the Periodic Table. In 1913, Henry Moseley discovered that every element had a unique amount of positive charge, which was later explained by the number of protons in the nucleus. The amount of positive charge of the atoms of an element is called its atomic number. He also noticed that the arrangement of the elements on the Periodic Table was in order of their atomic numbers. He discovered that periodicity is based on atomic numbers instead of atomic masses.

How to find out the Period with the Electron Configuration

Look at the primary energy level of valence electrons

Atoms will tend toward the ____ energy state possible for that particular situation

Lowest. A completed energy level has a lower energy state and is more stable. So an atom will try to achieve a complete energy level by gaining, losing, or sharing an electron.

Generally the ___ electrons an atom has, the larger is its diameter

More

Atomic Radius

One-half the distance between the nuclei of two atoms of the same element when the atoms are bonded together

Electrons repel-

Other electrons. Electrons in the lower levels repel electrons in the outer or valence levels. The force of repulsion does not cancel out the force of attraction by the protons but it does weaken the hold of the nucleus on the valence electrons. This is called shielding.

Electronegativity Explained

Recall that elements form compounds by either sharing electrons or by becoming ions. The type of bonding depends on how strongly the electrons of one atom are attracted to the nucleus of another atom. The measure of the ability of an atom to attract electrons from an atom of another element to form a compound is called electronegativity. The symbol for electronegativity is c. By comparing the electronegativity of two atoms, a chemist can determine what kind of chemical bond will be formed if atoms of the two elements react.

Ductility

The ability of a substance to be drawn, pulled, stretched, or extruded into a wire

Malleability

The ability of a substance to be hammered, shaped, or beaten into thin sheets

Electron Affinity

The amount of energy released when a neutral atom gains an electron to become an anion

Why do elements on the right side of the Periodic Table have smaller atomic radii than elements on the left side of the table when elements on the right have more electrons than those on the left?For instance, why does potassium, which has only 19 electrons, have an atomic radius of 227 pm when krypton, which has almost twice as many electrons (36), has an atomic radius of only 112 pm?

The answer to that question is that krypton also has 36 protons as opposed to 19 protons in a potassium atom. The increased positive charge of the protons in krypton holds its valence electrons tighter and reduces its atomic radius.

Julius Loather Meyer

The credit of who discovered the periodic table usually goes to Dimitri Mendeléev but it was also discovered by Meyer. He noticed that when he arranged the elements according to their atomic masses the elements didn't quite line up. So he began to consider other factors, such as atomic volumes (sizes of atoms) and valences (combining power). When he took all these things into consideration, a pattern of periodicity was revealed. He compiled a chart of about half of the elements known at that time and published his findings in a textbook in 1864. His table differed from Newlands' table in that instead of trying to force the elements to fit into octaves, he arranged them according to their actual chemical properties. In order to arrange them in groups with similar properties he had to leave several positions empty. He assumed that the empty positions were due to errors in calculation. He probably would have gotten credit for discovering the Periodic Table, except that Dimitri Mendeléev beat him to it.

Octave

The distance between the first and last note of a musical scale (8 tones = 1 octave); according to John Newlands, an octave is a row of 8 elements that show periodicity

Ionization Energy

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

Electronegativity

The measure of the ability of an atom of one element to attract electrons from an atom of another element to form a compound

Ionic Radius

The measure of the size of an ion.

Periodicity

The pattern of periodic repetition of properties of the elements when arranged according to their atomic numbers

Ionization

The process of gaining or losing electrons to create an ion

Periodic Trend

The tendency for a particular chemical or physical property to change in a particular direction in relation to the location of the elements on the Periodic Table

Karlsruhe Congress

There was no systematic method for calculating atomic mass so it was difficult to determine a relationship between atomic masses when the values were not accurate. In 1860, a group of European scientists decided to meet to discuss the problem. There wasn't many choices made until on the last day of the conference, an Italian chemist named Stanislao Cannizzaro submitted a paper he had written two years previously on the subject of atomic mas and they used it as an international standard. Two other delegates, both from Heidelberg, Germany, were Emi Erlenmeyer and Robert Wilhelm Bunsen. Erlenmeyer is best known for his flask. Bunsen developed a laboratory burner—the Bunsen burner—that has become a standard in most labs.

Descriptions of Groups 3-12 Elements

Transition elements are called d-block elements. They are all metals with typical characteristics of metals—they are good conductors of heat and electricity and they have metallic luster. Transition metals also possess malleability and ductility. Malleability is the ability of a material to be hammered, shaped, or beaten into thin sheets. Ductility is the ability of a material to be drawn, pulled, stretched, or extruded into a wire. Transition metals are less reactive than alkali metals or alkaline-earth metals. In fact, several transition metals, such as gold, palladium, and platinum, do not easily form compounds.

Groups

Vertical Columns

Different levels of Ionization Energies

When one electron is removed from a neutral atom, the atom becomes a cation with a positive charge . Instead of an atom with a neutral charge, it now has a charge of +1. ex- Li+. A lithium cation that has lost a second electron becomes Li2+ with a +2 charge. The second ionization energy is always greater than the first ionization energy, because when an atom loses an electron, the positive attraction of the nucleus is spread out over fewer electrons, so it holds them tighter. Some elements have dramatic differences between their first and second ionization energies. We can see why if we look again at its electron configuration . Lithium will easily give up its one 2s electron to become an Li+ cation with an electron configuration of 1s2 (full). An electron configuration of 1s2 is the same as that for helium, a very stable noble gas with a low energy state. If you try to take away a second electron, lithium puts up a fight because stable atoms (or ions) don't like to give up electrons. So it takes considerably more energy to remove a second electron (Fig. 3.7C). Also, if you tried to remove a second electron from lithium, the Li+ cation would become an Li2+ cation with an electron configuration of 1s1. An Li2+ cation has only 1 electron, but its nucleus still has 3 protons, so it would take a great deal more energy to overcome the attraction of the nucleus and take away the second electron.

Alkali Metals

an s-block element, excluding hydrogen, located in Group 1 on the Periodic Table of the Elements; a soft metal that reacts violently with water to form hydrogen gas

Luster

the ability of a substance to shine by reflecting light