The Periodic Table

History of the Periodic Table In 1869, Dmitri Mendeleev arranged the elements according to atomic mass in an attempt to classify them. He left spaces for elements that he predicted existed, yet hadn’t been discovered yet. This set him apart from others, and he was later proven correct.

Henry Moseley In 1913, Henry Moseley developed the concept of atomic numbers. He determined the frequencies of X-rays emitted as different elements were bombarded with high energy electrons. Each element produces X-rays of unique frequency, and the frequency increased as atomic mass increased.

Atomic Number Moseley arranged these frequencies in order by assigning a unique whole number, which he called the atomic number. He correctly identified the atomic number as the number of protons in the nucleus of the atom. This clarified some problems that Mendeleev had; Ar and K, Te and I.

Mendeleev and the Periodic Law When elements are arranged in order of increasing atomic number, their physical and chemical properties show a pattern or trend.

Effective Nuclear Charge Electrons are (-) charged and nuclei have a (+) charge. Many properties of the atoms depend on their electron configurations and how strongly the valence electrons are attracted to the nucleus.

Effective Nuclear Charge, continued A valence electron in an atom is attracted to the nucleus, and is repelled by the other electrons in the atom. The inner electrons (core) are effective in shielding or blocking the nuclear attraction. (Electron shielding effect.)

Effective Nuclear Charge = +1.

Effective Nuclear Charge, continued Zeff = Z - S Z is the number of protons in the nucleus S is the screening constant The value of S is usually close to the number of core electrons in the atom. The effective nuclear charge increases as we move across a period.

Periodic Table Arrangement Periods: Horizontal Rows Indicates the energy level where the valence electrons are located. Example: An element in Period 4 has its valence electrons in the 4th energy level.

Periodic Table Arrangement Groups/Families: Vertical columns Elements in the same column have similar valence electron configurations and therefore similar properties. Example: All elements in Group I have 1 valence electron.

Hydrogen is NOT a part of Group 1. It is a non-metal and a gas.

Groups 1 and 2 Group 1 – Alkali Metals Group 2 – Alkali Earth Metals 1 valence electron VERY reactive Group 2 – Alkali Earth Metals 2 valence electrons Reactive, but not as much as Group 1

Both Group 1 and Group 2 Reactivity increases as you go down a group Form cations Found as compounds only , in nature Low ionization energy and low electronegativity Large radii Radii of positive ions are smaller than the radii of atoms.

Radius of Sodium atom versus its ion Na Na 2e- 8e- 8e- 1e- Sodium ion with 10 electrons Sodium atom with 11 electrons

Groups 3 – 11 Transition Elements Metallic characteristics Have multiple (+) ions Colored compounds

Group 17 – Halogens 7 valence electrons These are the most active of the non-metals! Gains electrons easily and forms (-) ions F2 is the most reactive and I2 is the least reactive F2 is the most electronegative and has the most ionization energy

Group 18 – Noble Gases 8 valence electrons Inactive gases not found in compounds in nature

Inner Transition Elements (Bottom Row) Lanthanides – Lanthanoid series – ones that follow the element Lanthanum (#57) Rare earth elements that are less than 0.01% in nature Actinides – Actinoid series – ones that follow Actinum (#89) All are radioactive and unstable Those after U are not found in nature Elements above 92 are synthetic

Broad Categories of the Periodic Table

Metals The elements on the left side of the table. Malleable Ductile Lustrous Conducts heat and electricity Low ionization energies and low electronegativities Form cations because they lose electrons Mostly solids Reactivity increases as you go down a group.

Non-Metals The elements on the right side of the table Brittle as solids, with no luster Poor conductors of heat and electricity High ionization energies and high electronegativities Reactivities decrease as you go down a group Form anions because they gain electrons They can be gases, liquids or solids

Semi-Metals (Metalloids) Some properties are metallic, some are non-metallic and some are in-between They are found adjacent to the darkened line

Other information, continued Mono-atomic elements – He and all the Noble gases in Group 18 7 Diatomic Elements - BrINClHOF Allotropes – Two different structures of the same element Example – Carbon Graphite diamond buckyballs

Buckyball

Electronegativity This is the ability to attract electrons Based on Fluorine, which is 4.0 F has the greatest electronegativity (Foxy Fluorine steals other elements electrons!)

Ionization Energy The first ionization energy (I1) is the ENERGY needed to remove the first electron in a neutral atom. The second ionization energy (I2) is the energy needed to remove the second electron and so on. Measured in kJ/mol The greater the ionization energy the more difficult it is to remove an electron.

Trends in the Periodic Table

Trends of Atomic Radii (across a period) The trend increases from LEFT to RIGHT The radius decreases because the pulling power of the nucleus increases (due to an increase in the effective nuclear charge.) Atomic # 13 Atomic # 14 Atomic # 11 Atomic # 15 Na Al Si P Radius = 132 Radius = 128 Radius = 190 Radius = 143

Trend in Atomic Radii

Trends in Atomic Radii (down a group) As you go down (↓) a group, the atomic radius increases. Li 155 Na 190 K 235 Rb 240 Because . . . There is an increase in principal quantum numbers (n).

Radii of Neutral Atom versus Its Ion Na 2e- 8e- 8e- 1e- 1 e- removed Sodium ion with 10 electrons (Na +1) Radius = 120 pm Sodium atom with 11 electrons (Na 0) Radius = 190 pm

Radii of Neutral Atom versus Its Ion 1 e- added Fluorine ion with 10 electrons (F -1) Radius = 95 pm Fluorine atom with 9 electrons (F 0) Radius = 57 pm

The difference between the non-bonding atomic radius and the bonding atomic radius. The bonding atomic radius is shorter than the non-bonding radius.

Atomic Radii and Bond Lengths Knowing atomic radii allows us to estimate the bond lengths between different atoms in molecules. Ex: The Cl ̶ Cl bond length in Cl2 is 1.99 Å, so a radius of 0.99Å is assigned to Cl. In CCl4, the measured length of the C ̶ Cl is 1.77Å, close to the sum of the atomic radii of C and Cl (0.77Å + 0.99Å).

Isoelectronic Series This is a group of ions all containing the same number of electrons. The series: O-2, F-, Na+, Mg2+, and Al3+ all have 10 electrons. Nuclear charges increase as you move through a series, and the atomic radius of the ion decreases.

Trends of Electronegativity and Ionization Energy (across a period) Electronegativity INCREASES Ionization Energy INCREASES The increased effective nuclear charge makes it harder to remove electrons. WHY?

Trends of Electronegativity and Ionization Energy (down a group) WHY? Ionization Energy DECREASES Electronegativity DECREASES Electrons are held less tightly because energy levels increase and there is more electron shielding.

Electron Affinity Most atoms gain electrons to form anions. The energy change that occurs when an electron is added to an atom is called electron affinity because it measures the attraction (or affinity) of the atom for the added electron. For most atoms, energy is released when electrons are added. Cl(g) + e- → Cl-(g) ΔE = -349 kJ/mol

Trends of Metallic Characteristics Down a group – INCREASE in metallic character Across a period – DECREASE In metallic character (There is a decreased attraction for electrons.) (There is an increased attraction for electrons.)

Metallic Character Metals have low ionization energies and therefore tend to form positive ions easily. As a result, metals are oxidized (lose electrons – LEO) when they undergo chemical reactions. Most metal oxides are basic. Those that dissolve in water react to form metal hydroxides. CaO (s) + HOH →Ca(OH)2 (aq)