Chapter 5: The Periodic Law

Chapter 5:The Periodic Law 5.1: History of the Periodic Table 5.2: Electron Configuration & the Periodic Table 5.3: Electron Configuration & Periodic Properties

5.1: History of the Periodic Table • September 1860- International Congress of Chemists in Karlsruhe, Germany. • Problem: how to determine the atomic mass… so that everyone is getting the same #. • Italian Chemist , Stanislao Cannizzaro’s method of determining atomic masses was accepted- created a standard value.

Döbereiner • In the early 1800’s J.W. Döbereiner began classifying elements into groups of three (TRIADS).For example, lithium, sodium and potassium; calcium, strontium and barium; AND chlorine, bromine, and iodine are three sets of triads.The elements within the triads share similar chemical properties.

Newlands • Then in 1865, J.A.RNewlands noticed that when the known elements were arranged in increasing atomic mass, regular patterns would occur, AND they would repeat every 8 elements(law of octaves).

Mendeleev & Chemical Periodicity • In 1869, Dmitri Mendeleev(and Lothar Meyer independently)published practically identical schemes for classifying elements. Mendeleev also noticed regular patterns in the elements when they were arranged by increasing atomic mass. He produced the first periodic table of elements and actually left spaced to be filled in by undiscovered elements. He was able to predict there would be more and about where they would go!

Mendeleev's Periodic Table of the Elements

Moseley & the Periodic Law • The periodic table Mendeleev came up with was a good start but it wasn’t perfect. It wasn’t until 1913 when H.G.J. Moseley determined that the correct way to organize the periodic table is by ATOMIC NUMBER. • The PERIODIC LAW- when elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern. Increases in atomic #

The Modern Periodic Table • Periodic Table- is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.

5.2 • Organizing the Squares: • When you look at the modern Periodic Table of Elements (PTE) you will notice it is arranged in to columns(vertically)and rows(horizontally). The columns are called GROUPS or FAMILIES(18). The rows are called PERIODS(7). Periods Groups

Labeling & Naming Groups • There are 3 methods for numbering the columns(the American, the European and the IUPAC) The only 1 your book uses is the IUPAC. • IUPAC: 1-18 • AMERICAN: 1A, 2A, 3B-8B, 11B, 12B, 3A-8A(on the PTE I gave you) • The families are given names: • 1- ALKALI METALS (except hydrogen) • 2-ALKALINE EARTH METALS • 17- HALOGENS • 18- NOBLE GASES

Electronic Configuration & the PTE: • The electrons in the outermost energy level are called theVALENCE electrons. • These are the electrons which will allow one atom to interact with another atom and are responsible for the elements chemical properties. • The members in each group/family have the same # of valence electron and that is why they share similar chemical properties.

The s-, p-, d- and f-block Elements: • S-BLOCK- groups 1&2 • P-BLOCK- groups 13-18 • D-BLOCK- groups 3-12 • F-BLOCK- the lanthanide & actinide series (the two rows at the very bottom of the PTE)

The s-Block: groups 1&2 • The metals that are the most reactive are found in the two-column s-block of the PTE. These are called the alkali metals and the alkaline earth metals.

ALKALI METALS • These are the metals in Group 1 (1A): Li, Na, K, Rb, Cs, and Fr. The name for the family is derived from the Arabic word meaning “ashes” because we find Na and K in the ashes of burnt plants. • *Properties* These metals are: malleable, ductile and good conductors with low densities and low melting points. They are “soft” and are able to be cut with a knife and when exposed to the air, tarnish easily. • THESE METALS ARE EXTREMELY REACTIVE!!! • The react readily with water and air and must be stored in evacuated containers… or placed in oil.

ALKALINE EARTH METALS • This comprises the elements in Group 2 (2A). Be, Mg, Ca, Sr, Ba, and Ra. • *Properties* These metals have higher densities and higher melting points as well as higher ionization energies. Ergo they are not as reactive as the alkali metals. They all have 2 valence electrons which are fairly easy to lose, forming a 2+ ion. In order for the Alkaline earth metals to combine with halogens you will need 2 halogens for every 1 alkaline earth metal (MX2).

Hydrogen: One of a Kind • It has 1 valence electron with a high ionization energy but will form a + ion. It is a nonmetal and a gas at room temperature which is odorless, colorless and also diatomic(H2).“Most of the Earth’s hydrogen is combined with oxygen as water”. Remember… it also combines readily with C to make hydrocarbons (fuels). It is the 9th most abundant element on earth but the absolute most abundant element in the universe! • We get H from hydrocarbons, commercially. And we use H in ammonia, which can be used as a fertilizer, and in other organic compounds like alcohols, gases (methane), etc.

The d-Block: groups 3-12 • Transition metals play an important role in living organisms, and are also extremely valuable as strong, structurally useful materials. From iron captured hemoglobin and cobalt in vitamin B12 to gold and silver jewelry, there are many different and important uses for the transition metals. • *Properties* They vary from family to family but most have high densities and melting points, are malleable, ductile and good conductors.

The p- Block: groups 13-18 • P-block & s-block elements constitute the main-group elements. • Properties- very greatly… includes all of the nonmetals and 6 metalloids. • Metalloids- aka semiconductors, fall on both sides of the line that separate nonmetals and metals. • Other metals • Other non-metals • Halogens • Nobel gases

THE HALOGENS • The name comes from the Greek meaning “salt former” because most will form salts. They each have 7 valence electrons and will accept ONE electron to form a – ion. ALL of the halogens are diatomic and react with most metals and some nonmetals. They are HIGHLY REACTIVE and are found combined with other elements in nature. • WHY? Because all they need in 1 more electron and will get it from any element that is willing to give or share it!

THE NOBLE GASES • The most abundant is Argon and it was the first be discovered (1894). However, it does not form any compounds (argon means “lazy one” in Greek). Now… a whole new, previously unknown and not predicted group of element has been discovered, too. Soon after the other noble gases were found and added to the PTE. • It was originally though that the noble gases never formed compoundsuntil 1962 when XeF2, XeF4, and XeF6 were formed. Then came KrF2, but to date there are no know compounds of He, Ne, or Ar. • He is the most commercially important noble gas. It is used as a coolant for performing experiments at low temps. • Ar is used in light bulbs to conduct heat away from the filament. • Ne makes the red light in “neon” signs.

The f- Block: Lanthanides & Actinides • The elements of the 4f series are generally called the lanthanides, after the element lanthanum (La), which is the first member of the series. Similarly, the elements of the 5f series are called actinides, after the first member of that series, actinium (Ac).

THE LANTHANIDES • *Properties* they are very similar to each other and easily lose 3 electrons forming 3+ ions. All are silvery and quite soft but heavier and less reactive than the alkali metals. They tarnish easily but react slowly in water. They are found together in nature and are difficult to separate. • *USES*some alloys are used to make special steel. Some compounds are found in the color in your television.

THE ACTINIDES • These elements are all radioactive- EVERY LAST ONE OF THE ISOTOPES OF EVERY ELEMENT! Only Th and U exist in nature… all others are primarily man made.

5.3 Electron Configuration and Periodic Properties • As you move across the periodic table you begin to see trends/predictable patterns in the chemical and physical properties of the elements. This phenomenon is known as PERIODIC TRENDS.

Atomic Radius • Atoms have a radius called the ATOMIC RADIUS- the distance from the center of the atom to the farthest most/ highest energy level. This distance is not exact.We estimate. • Contrary to popular opinion… atoms with more electrons do not always have a larger atomic radius. The trend is: • As you move from TOP to BOTTOM down the PTE the radius grows larger • As you move from RIGHT to LEFT across the PTE the radius grows larger

ATOMIC RADII Largest

The first trend has a simple explanation: as you move down through a group/family you are actually adding additional principle quantum numbers(ex group 1: 1s, 2s, 3s, 4s, 5s, 6s, and 7s) with corresponding increases in the number of electrons from 1, 3, 11, 19, 37, 55, and 87. Ergo the radius must increase to accommodate these increases in electrons. • The second trend is more complex. As you move across a period you are increasing both the # of electrons as well as the number of protons. This allows the protons to pull more strongly on the electrons in the same energy level thus shrinkingthe radius.

Ionic Size • We know atoms can gain or lose electron to form Ions. Depending whether they gain (becoming negatively charged) or lose (becoming positively charged) will depend on whether ionic size will shrink or grow. • When you gain electrons- the ionic size grows • As you lose electrons- the ionic size shrinks(more protons pulling on fewer electrons= shrinking). • The trend is complex- Figure 5-19 p.149

Ionization Energy • The energy needed to strip an atom of an electron is called IONIZATION ENERGY. Typically measured in J/atom you can think of it as a measure of how strongly the atom holds on to its valence electrons. • Its trend is exactly opposite the atomic radius trend: • As you move from BOTTOM to TOP the ionization energy increases • As you move from LEFT to RIGHT the ionization energy increases

Ionization Energy Hardest Easiest

Successive Ionization Energies What if you want to pull more than one electron from an atom? What happens to the ionization energy for the 2nd, 3rd, or even 4th electron? • Successive ionization energies refer to the energy required to remove the 2nd, 3rd and more electrons. The energy required increases, partially because of electron-electron repulsion, however it is not a linear increase. • *note* there is always a dramatic increase in ionization energy when you start extracting electrons from the inner noble gas core.

Electron Affinity • If you can pull an electron from an atom… you can ADD one to an atom. When an atom gains an electron energy is REQUIRED! They are reported in J/mol(NOT ATOMS). Not all energy used in the addition of an electron is gained by the atom that the electron is being added to. It is possible to have + and – electron affinities. If the value is + then the atom gaining the electron is also gaining energy. If the value is – then the atom gaining the electron is releasing energy.

General rule for the electron affinity trend- • Nonmetals have more NEGATIVE electron affinities than the metals. • The exception is the noble gases. • ** OCTET RULE** atoms tend to gain, lose or even share electrons in order to have a full set of valence electrons (8)- to have a full s and p orbital. More Negative

Electronegativity • This is the ability of an atom to attract an electron in a chemical bond. This is a unit less value. • The least electronegative elements are in the LOWER LEFT CORNER of the PTE. The trend is the same as Ionization energy. • When bonds occur… the atom with the higher electronegativity would attract the electron more strongly. Depending on the discrepancy of the attraction there will form different types of bonds.

More electronegative

Chapter 5: The Periodic Law