Periodic Table of Elements

Elements

� Science has come along way since Aristotle’s theory of Air, Water, Fire, and Earth.

� Scientists have identified 90 naturally occurring elements, and created about 28 others.

Elements

� The elements, alone or in combinations, make up our bodies, our world, our sun, and in fact, the entire universe.

Mendeleev – Father of the Periodic Table

� In 1869, Dmitri Mendeléev created the first accepted version of the periodic table.

� He grouped elements according to their atomic mass, and as he did, he found that the families had similar chemical properties.

� Blank spaces were left open to add the new elements he predicted would occur – scandium, germanium.

Glenn T. Seaborg (1912 – 1999)

�Reference: U.S. DOE. (2012). Glenn T. Seaborg

�Contributions to Advancing Science. Research and

Development from the

U.S. DOE.

Downloaded from: “http://www.osti.gov/accomplishments/seaborg.html”

Seaborg, McMillan, Kennedy, & Wahl, isolated element 94, plutonium (Pu) in 1940.

Isolated Uranium-233 in 1941 & established Th's nuclear fuel potential.

1944, Seaborg formulated the 'actinide (Ac) concept' of heavy element electronic structure which predicted the actinides – -specifically, the first eleven trans-uranium elements. This resulted in a rearrangement of the periodic table.

The Ac series formed a transition series analogous to the rare earth series of the lanthanide (La) elements.

Considered most significant change to the periodic table since Mendeleev designed the table in the mid-1800’s. The actinide concept showed how the trans-uranium elements fit into the periodic table.

From 1944 -1958, Seaborg identified eight additional elements –Americium (95), Curium (96), Berkelium (97), Californium (98), Einsteinium (99), Fermium (100), Mendelevium (101), and Nobelium (102). Element 106, Seaborgium, bears his name.

In 1951, Seaborg shared the Nobel Prize in Chemistry with E.M. McMillan for "for their discoveries in the chemistry of the trans-uranium elements"

Names of Elements

� Property – Sodium from suda (Arabic for headache)

� Mineral in which found – Lithium, Boron, Argentum

� Place of discovery –Germanium, Californium, Yttrium, Ytterbium, Terbium, Erbium

� Person (scientist) –Curium, Einsteinium, Seaborgium

Chemical Symbols - Berzelius

� Shorthand for element

� Use first letter of name (capital letter)

� Second letter is important in name - if needed (lower case)

� IUPAC has final say: International Union of Pure and Applied Chemistry (Seaborgium story)

Key to the Periodic Table

� Elements are organized on our modern periodic table according to their atomic number, usually found near the top of the square.� The atomic number refers to

how many protons an atom of that element has.

� For instance, hydrogen has 1 proton, so it’s atomic number is 1.

� The atomic number is unique to that element. No two elements have the same atomic number.

Atomic Number� This refers to how many

protons an atom of that element has.

� No two elements, have the same number of protons.

Bohr Model of Hydrogen AtomQuantum Mechanical

Wave Model

Valence Electrons

� The number of valence electrons an atom has can be found from the last digit of the group number for the s and p blocks. (Arsenic is in group 15 – has 5 valence electrons)

� Valence electrons are the electrons in the outer energy level of an atom.

� These are the electrons that are transferred or shared when atoms bond together.

Properties of Metals

� Metals are good conductorsof heat and electricity.

� Metals are shiny “luster”.� Metals are ductile (can be

stretched into thin wires).� Metals are malleable (can be

pounded into thin sheets).� Metals lose electrons to

become stable. They are the “losers” and become positively charged cations.

� They lose their outer energy level, becoming smaller as a cation than they were as an atom.

Properties of Non-Metals� Non-metals are poor

conductors of heat and electricity.

� Non-metals are not ductile or malleable.

� Solid non-metals are brittle and break easily.

� They are dull.� Many non-metals are

gases.� They gain electrons to

form negative anions. They become larger when they gain elctrons.Sulfur

Properties of Metalloids� Metalloids are located

along the stairstep line.� Metalloids (metal-like)

have properties of both metals and non-metals.

� They are solids that can be shiny or dull.

� They conduct heat and electricity better than non-metals but not as well as metals. Si and Ge are used in computer chips.

� They are ductile and malleable.Silicon

Families Periods� Columns of elements are

called groups or families.

� Elements in each family have similar properties.

� For example, lithium (Li), sodium (Na), potassium (K), and other members of family IA are all soft, white, shiny metals.

� All elements in a family have the same number of valence electrons and outer electron configuration ending.

� Each horizontal row of elements is called a period.

� Elements in the period have their valence electrons in the same energy level.

� The elements in a period are not alike in properties.

� In fact, the properties change greatly across any given row.

� The first element in a period is always an extremely active solid. The last element in a period, is always an inactive gas.

Hydrogen

� Hydrogen sits on top of Group 1, but it is not a member of that family. Hydrogen is in a class of its own.

� It’s a gas at room temperature.

� It has one proton and one electron in its one and only one energy level.

� Hydrogen only needs 2 electrons to fill up its valence shell, but is actually more stable with no electrons at all.

Alkali Metals

� The alkali metal family is found in the first column of the periodic table.

� Atoms of the alkali metals have a single electron in their outermost level, in other words, 1 valence electron and end in s1 in the same energy level as the period.

� They are shiny, have the consistency of clay, and are easily cut with a knife.

Alkali Metals� They are the most reactive

metals.� They react violently with

water. More violent as you go down group.

� Alkali metals are never found as free elements in nature. They are always bonded with another element.

� They lose 1 electron and form a +1 cation.

�Alkali video

What does it mean to be reactive?� We will be describing elements according to their

reactivity. � Elements that are reactive bond easily with other

elements to make compounds.� Some elements are only found in nature bonded

with other elements. � What makes an element reactive?

� An incomplete valence electron level.� All atoms (except hydrogen) want to have 8 electrons in

their very outermost energy level (This is called the octet rule.)

� Atoms bond until this level is complete. Atoms with few valence electrons (4 or less) lose them during bonding. Atoms with 6, 7, or 8 valence electrons gain electrons during bonding. They can also share electrons.

Elements, Compounds, Mixtures

� Sodium is an element.

� Chlorine is an element.

� When sodium and chlorine bond they make the compound sodium chloride, commonly known as table salt.

�Compounds have different properties than the elements that make them up.

�Table salt has different properties than sodium, an explosive metal, and chlorine, a poisonous gas.

5

Alkaline Earth Metals

� Have 2 valence electrons, lose them, become +2 cation.

� Harder, more dense, stronger, and have higher melting points than Group 1.

� Not as reactive as Group 1.

� They are never found uncombined in nature.

Halogen Family

� The elements in this family are fluorine, chlorine, bromine, iodine, and astatine.

� They react with alkali metals to form salts.

� Halogens have 7 valence electrons, which explains why they are the most reactive non-metals.

� Found only as compounds in nature.

�Halogen atoms only need to gain 1 electron to fill their outermost energy level, form -1 anion.�All are diatomic elements –F2, Cl2, Br2, I2

Noble Gases

� Noble Gases are colorless gases that are extremely un-reactive. � One important property of the noble gases is their inactivity.

They are inactive because their outermost energy level is full. � Because they do not readily combine with other elements to form

compounds (none in nature), the noble gases are called inert. Only xenon and krypton can be forced to form compounds in the lab.

� The family of noble gases includes helium, neon, argon, krypton, xenon, and radon.

� All the noble gases are found in small amounts in the earth's atmosphere.

Transition Metals

� Transition Elements include those elements in the middle d block.

� These are the metals you are probably most familiar: copper, tin, zinc, iron, nickel, gold, and silver.

� They are good conductors of heat and electricity.

Transition Metals

� The compounds of transition metals are usually brightly colored and are often used to color paints.

� Transition elements have 1 or 2 valence electrons, which they lose when they form bonds with other atoms. Some transition elements can lose electrons in their next-to-outermost d sublevel.

� Most form +2 cations and others

Transition Elements

� Transition elements have properties similar to one another and to other metals, but their properties do not fit in with those of any other family.

� Many transition metals combine chemically with oxygen to form compounds called oxides.

Inner Transition Metals - “Rare Earth Elements”

� The thirty inner transition elements are composed of the lanthanide and actinide series.

� One element of the lanthanide series and most of the elements in the actinide series are called trans-uranium, which means synthetic or man-made.

� Not as rare as once thought to be.

Chalcogen or Oxygen Family

� Atoms of this family have 6 valence electrons, gain 2 to be

-2 anion.

� Most elements in this family also share electrons when forming compounds.

� Chalcogen means “chalk-former”

� Oxygen is the most abundant element in the earth’s crust. It is extremely active and combines with almost all elements.

�Tom Lehrer “Elements” song

Boron Family

� The Boron Family is named after the first element in the family.

� Atoms in this family have 3 valence electrons, lose them to form +3 cations.

� This family includes a metalloid (boron), and the rest are metals.

� This family includes the most abundant metal in the earth’s crust (aluminum).

Carbon Family� Atoms of this family have 4

valence electrons, form +2, +4 cations. Carbon can form carbide, a -4 anion.

� This family includes a non-metal (carbon), metalloids, and metals.

� The element carbon is called the “basis of life.” There is an entire branch of chemistry devoted to carbon compounds called organic chemistry.

Nitrogen Family

� The nitrogen family is named after the element that makes up 78% of our atmosphere.

� This family includes non-metals, metalloids, and metals.

� Atoms in the nitrogen family have 5 valence electrons. They tend to share electrons when they bond, but can gain 3 electrons to form -3 anions.

� Other elements in this family are phosphorus, arsenic, antimony, and bismuth.

ELECTRON CONFIGURATION can easily be determined from the layout of the periodic table

We can find:

� Sublevel blocks

� Configuration ending

� Main (highest) energy level for elements

� Number of valence electrons

SUBLEVEL BLOCKS

� s sublevel

groups 1 and 2: first two groups and helium

� p sublevel

groups 13-18: last six groups

� d sublevel

groups 3-12: middle elements (transition)

� f sublevel

Bottom two rows: lanthanide and actinide series (inner transition)

VALENCE ELECTRONS

� The electrons in the highest energy level (s and p sublevels) are called the valence electrons

� Generally, these are the electrons used for bonding

� The number of valence electrons can be determined from the last digit of the group number for s and p blocks (middle all have 2)

� For example: arsenic is in group 15 so it has 5 valence electrons

Stability depends on valence electrons

� 1. Octet Rule: having 8 valence electrons makes an atom unreactive. How does helium fit this rule? (full outer level)

� 2. Filled or half-filled sublevel is more stable (less reactive)

� Cr: [Ar]4s23d4 → [Ar]4s13d5

� Cu: [Ar]4s23d9 → [Ar]4s13d10

Unit 4 – Periodic Table

WXHS, Honors Chemistry

PERIODIC PATTERNS

ATOMIC RADIUS

� Position in the table and properties of the elements are due to its electron configuration.

� Radius is the distance from the center of the nucleus to the “edge” of the electron cloud. The “edge” is considered the 90% probability calculation of finding the electron.

� Atomic radii are usually measured in picometers (pm) or angstroms (Å). An angstrom is1x10-10 m, and a pm = 1x10-12 m.

ATOMIC RADIUSCovalent Bond Determination

� Since a cloud’s edge is difficult to define,

scientists calculate the atomic radii in

one of two manners. For covalently

bound atoms, the radius is simply ½ the

distance from the center of one atom to

the other (in the diatomic molecule).

� BROMINE = Br2

�114 pm�114 pm

�228 pm

ATOMIC RADIUSTrends in the Periodic Table

ATOMIC RADII TRENDS

� As you move down a

family, or a group, a new

energy level is added.

� In other words, a new

floor is added to the

building with each new

period.

� Shielding also increases

(more electrons in

between )

� DOWN A FAMILY OR GROUP � WHY?

Increases

ATOMIC RADII TRENDS

� As you go across a period the

number of protons increases.

Thus, the nuclear charge

increases which pulls the

electrons in closer to the

nucleus.

� This decreases the radius.

� ACROSS A PERIOD � WHY?

�DECREASES

Radius size is affected by:

� Distance from nucleus – more energy levels = bigger atom

� Size of nuclear charge – more protons in same period = smaller atom

� Shielding effect of inner electrons –lowers the attraction of the nucleus for the outer electrons

As more electrons are added to atoms, the inner layers of electrons shield the outer electrons from the positive charge of the nucleus.

The effective nuclear charge on those outer electrons is less, and so the outer electrons are less tightly held.

SHIELDING EFFECT

Example of Shielding Effect

IONS

�Gain electrons becoming negative. Nuclear charge is spread among more electrons, so they are not held as tightly. This makes them larger than the original atom.

i.e. Chlorine – Cl

[Ne]3s23p5

Gains one e- becoming Cl-1 and now has the [Ar] Noble gas configuration.

The anion also obeys the Octet Rule.

� Metals � NonMetals

Lose electrons becoming positive. Lose entire outer level so they become smaller than the original atom.

i.e. Calcium – Ca

[Ar]4s2

Loses 4s2 e-s becoming Ca+2, and the cation now has the electronic configuration of the Noble gas [Ar].

The cation obeys the Octet Rule

IONIC RADII TRENDSRadii values given in (Angstroms) & pm

� When we proceed down a group, the energy level energy level energy level energy level increases with each increases with each increases with each increases with each new periodnew periodnew periodnew period....

� The outer energy level is further away from the nucleus, and the ion’s radius increases.

DOWN A FAMILY OR GROUP Scientific Reasoning

�INCREASES

IONIC RADII TRENDSRadii values given in (Angstroms) & pm

� There is a big jump in size when cross from metal to nonmetal, but otherwise follows same trend as the atomic radii.

� Metals are smaller than the atom because they lose outer level.

� When we look at the nonmetals, the radius increases in size because the atom, or anion, gained electrons; thus, the outer electron, or electrons, is (are) not held as tightly by the nucleus.

ACROSS A PERIOD SCIENTIFIC REASONING

�DECREASES then INCREASES

METALLIC ATOM AND ION COMPARISON*Radii given in units of Angstroms

NONMETALLIC ATOM AND ION COMPARISON * Radii given in units of nm

�Why do the Noble Gases not have an ionic Radius?

ATOM AND ION COMPARISON*Radii given in units of nanometers (nm)

�Why does Hydrogen not have an ionic Radius?

The energy required to remove an electron from an atom is ionization energy. Successive ionization energies are the

respective energies required to remove subsequent electrons from an ion.Ionization energies are generally measured in kilojoules per mole (kJ/mol), and sometimes in electron volts (eV).

IONIZATION ENERGY

PERIODIC TABLE TRENDS with

respect to IONIZATION ENERGY

�The larger the atom is, the easier its electrons are to remove.�Ionization energy and atomic radius are inversely proportional.

�Increases across period – nuclear charge increases so harder to remove�Decreases down group – shielding and energy levels increase with each period�Metals have low IE, nonmetals have high IE�Measured in gas phase.

�IONIZATION ENERGY (cont)

�Ionization Energy (eV) v Atomic Number

�Dips show where atom is more stable with half or filled sublevel. Between s and p blocks (full s) and between p3 and p4 (half-filled p sublevel)

PERIODIC TABLE TRENDS

in IONIZATION ENERGIES

�INCREASES

�INCREASES

Electron Affinity is a measure of the energy change when an electron is added to a neutral atom, in the gaseous state, to form a negative ion.

This measures the ability to attract electrons as a neutral atom.

Decreases down group-More energy levels and shielding decrease attraction felt from nucleus

Increases across period – increased nuclear charge

Energy is released as atom becomes more stable.

ELECTRON AFFINITY

ELECTRON AFFINITY

�Why do the Alkaline Earth Metals and Noble Gases not have measurable Electron Affinities?

Electronegativity is a measure of the tendency of an atom, in a molecule, to attract electrons in a bond when they are being shared).

Trend is same as electron affinity and for the same reasons.

Highest electronegativity – Fluorine (most reactive nonmetal)

ELECTRONEGATIVITY

ELECTRONEGATIVITY of the

Elements in the Periodic Table

Reaction Tendencies are Determined by:

� Number of energy levels

� Shielding effect of inner electrons

� Size of nuclear charge

� Stability of filled or half-filled sublevels

ISOELECTRONIC SPECIES

� When a noble gas atom and ions all have the same configuration

� Examples: N3-, O2-, F-, Ne, Na+, Mg2+, Al3+

METALLIC TREND�INCREASES

Summary of the major trends in the

Periodic Table Trends

�Electronegativity

�Electronegativity