is an arrangement of the elements according to their properties. It enables chemists to classify the elements so that it is possible to identify patterns and trends in their properties. Many scientists have made significant contributions to the development of the modern Periodic Table,
Arranged elements in groups of three (triads) For example: Cl (mass 35) Br (mass 80) I (mass 127) Note: Average mass of the extreme elements =mass of the middle element
Father of the Modern Periodic table. Based on increasing atomic mass Chemical properties of elements reoccur in a periodic fashion.
Like Meyer prepared a table based on increasing atomic mass. Consisted of 17 columns with 63 elements Widely accepted Constantly revised as new masses were discovered Left vacant spots for new elements (Ga, Cd,Ge,Tc, Re, Po) Certain discrepancies as new masses were found: Ar and K, Co and Ni, Te and I
English scientist used x-rays to determine atomic number (number of protons) Discovered that the properties of elements were periodic functions of their atomic numbers, Many of the discrepancies disappeared.
Based on increasing atomic number Periodic Law: properties of the elements depend on the structure of the atom and vary with the atomic number in a systematic way Horizontal rows are called periods (#1-7) Indicates the number of principal energy levels. Note: Lanthanide series part of period 6—long period and Actinide series part of period 7— unfinished period Properties of elements change within a period Vertical columns are called groups or families (#1- 18) Indicates the number of valence electrons. Elements within a group have similar properties. Ex.Na and Li
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What is a trend? A trend is the general direction in which something tends to move.
The elements on the Periodic Table of Elements show many trends in their physical and chemical properties. Across the rows (periods or series) Down the columns (groups or families)
½ the distance between the nuclei of 2 like atoms in a diatomic molecule The atoms of the 8 main groups are shown here.
Atomic radius increases as you move down a group Why? More electrons in more Principal Energy Levels Atomic size increases
Atomic radius decreases as you move across a period Why? (-) electrons increase, but so do (+) protons !!! Increased (+) nuclear charge pulls the (-) electrons closer to the nucleus Atomic size decreases
The size trend in periods is less pronounced than in groups because of the electron shielding effect.
Reduction in effective nuclear charge on an electron that is caused by the repulsive forces of other electrons between it and the nucleus In an atom with one electron, that electron experiences the full charge of the positive nucleus. However, in an atom with many electrons, the outer electrons are simultaneously attracted to the positive nucleus and repelled by the negatively charged electrons.
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What are ions? Ions are charged atoms, either + or - Cations are positive ions Cations form when atoms lose electrons Anions are negative ions Anions form when atoms gain electrons
Cations are smaller than their parent atoms. Why? By losing their valence electrons, they lose their entire valence shell Cations are formed by the metals on the left side of the Periodic Table
Ionic size increases as you move down a group for the same reason atomic size increases Number of principal energy levels increases
Anions are larger than their parent ions Why? When extra (-) electrons are added, extra (+) protons are NOT added to the nucleus Effective nuclear attraction is less for the increases number of electrons
Ionic size increases as you move down a group for the same reason atomic size increases Number of principal energy levels increases Cations are formed by nonmetals on the right side of the Periodic Table
JJust like their parent atoms… CCations get smaller as you move from left to right AAnions get smaller as you move from left to right IIncreased (+) nuclear charge pulls the (-) electrons closer to the nucleus
Energy is needed to remove an electron from an atom The energy needed to overcome the attraction of the nuclear charge and remove an electron (from a gaseous atom) is called the Ionization Energy
The energy needed to remove the 1 st electron from an atom is the 1 st Ionization Energy
Atomic Radius Smaller atoms hang on to valence electrons more tightly, and so have higher ionization energy
Charge The higher the positive charge becomes, the harder it is to pull away additional electrons Second ionization energy is always higher than the first
Orbital Type It's easier to remove electrons from p orbitals than from s orbitals, which are “deeper”
Electron Pairing Within a subshell, paired electrons are easier to remove than unpaired ones Reason: repulsion between electrons in the same orbital is higher than repulsion between electrons in different orbitals
Electron Pairing – Example On the basis of gross periodic trends, one might expect O to have a higher ionization energy than N. However, the ionization energy of N is 1402 kJ/mol and the ionization energy of O is only 1314 kJ/mol. Taking away an electron from O is much easier, because the O contains a paired electron in its valence shell which is repelled by its partner.
Subsequent electrons require more energy to remove than the first electron How much more energy is needed depends on what energy levels and orbitals the electrons are in
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Ability of an atom to attract electrons toward itself in a chemical bond
The difference between the electronegativities of two atoms will determine what kind of bond they form Linus Pauling used an element's ionization energy and electron affinity to predict how it will behave in a bond. The more energy it takes to pull off the outer electron of an atom, the less likely it is to allow another atom to take those electrons. The more energy the atom releases when it gains an electron, the more likely it is to take electrons from another atom in bonding. These two energies were used to compute a numerical score.
Electronegativity increases going left to right across the periodic table. Fluorine's high nuclear charge coupled with its small size make it hold onto bonding electrons more tightly than any other element. Lithium has a lower nuclear charge and is actually larger than fluorine. Its valence electron is not tightly held and it tends to surrender it in chemical bonds. LiBeBCNOF
Electronegativity decreases going down a group The bonding electrons are increasingly distant from the attraction of the nucleus H 2.1 Li 1.0 Na 0.9 K 0.8 Rb 0.8 Cs 0.7 Fr 0.7
The electron affinity is a measure of the energy absorbed when an electron is added to a neutral atom to form a negative ion. Most elements have a negative electron affinity. This means they do not require energy to gain an electron; instead, they release energy. Atoms more attracted to extra electrons have a more negative electron affinity. The more negative the value, the more stable the ion is.
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Electron affinity is essentially the opposite of the ionization energy.
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PowerPoint: Adela J. Dziekanowski