Periodic Law

Periodic Table Prior to 1860 no agreement/method to accurately determine masses of atoms. First International Congress of Chemists – 1860 Stanislao Cannizzaro presented method for accurately measuring atomic masses Looked for relationships between atomic masses and other properties of elements

First tables arranged elements by atomic weight Could not agree on atomic weights therefore tables were different

John Newlands Noticed elements properties repeated every 8th element when arranged by atomic mass Named this phenomenon “the Law of Octaves” Did not work for all elements

Julius Lothar Meyer Developed first modern table Consisted of 28 elements divided into 6 families Families (groups) had similar chemical and physical properties Discovered all elements in same family had same number of valence e- -- outermost electrons in highest energy level Why?

Dmitri Mendeleev Noticed that properties repeat themselves at certain intervals Arranged all known elements into one table based on properties– 1869 1871 - Proposed the “Periodic Law” Based on the properties spaces were left for unknown elements (Sc, Ga, Ge)

Upon discovery of other elements inconsistencies were found with Mendeleev’s table Atomic masses improved and they no longer arranged the elements by increasing atomic mass

Why can most elements be arranged by atomic mass? What was the reason for chemical periodicity?

Henry Mosely Discovered elements contain unique number of protons (atomic number) - 1911 Arranged elements by atomic number - 1913 Fully explained the Periodic Law

Periodic Law The physical and chemical properties of the elements are periodic functions of their atomic numbers. Aka – when elements are arranged by increasing atomic number, elements with similar properties appear at regular intervals.

Parts of the Periodic Table Noble Gases – added to the table in 1894 after the discovery by Lord Rayleigh and William Ramsey First discovered Argon while studying nitrogen Later discovered Helium Highly inert (unreactive) due to a full octet

Parts…. Alkali metals – group 1 Alkaline earth metals – group 2 Halogens – group 17 Transition metals – d block elements Inner Transition metals Lanthanides (elements 58-71) added in early 1900’s Have very similar properties Actinides (elements 90-103)

Electron Configuration & the Periodic Table

s-Block Elements Groups 1 & 2 All elements in group 1 & 2 will have an electron configuration of ns1 or ns2 where n = highest energy level occupied

Alkali Metals Group 1 elements In the elemental state Soft Silvery metal High melting points Extremely reactive therefore are not found in elemental state in nature React violently with water to produce hydrogen gas

Alkaline – Earth Metals Group 2 elements Outer most s orbital is full Do not exhibit stability (outer p orbital is empty) Properties Harder, denser than group 1 Higher melting points than group 1 Not as reactive but too reactive to be found in nature in elemental form

Burning Mg

Hydrogen & Helium H has same valence electrons as group 1 but does not share any other properties He share same electron configuration (valence e-) as group 2 but does not share same properties Placed with group 18 because it is very stable

d-block elements Transition elements Beginning filling the 3d orbitals Good conductors of electricity High luster Less reactive than s-block elements Can be found in elemental form

Exceptions in the d-block The following elements have odd configurations Cr: [Ar]4s13d5 Cu: [Ar]4s13d10 Ag: [Kr]5s14d10 More stable with half filled s & d orbitals or full d orbital Exceptions follow throughout the d element similar to Chromium and Copper

p-block elements All elements in p block have a full s orbital Properties Contain all non metals except H & He Contain all metalloids (exhibit properties of both metals and non metals) Have semi conducting properties Contains 6 metals Elements in s & p block make up the representative elements

Halogens Group 7A/17 Most reactive non metals (Fluorine is most reactive) Will bond with a metal to form a salt F & Cl are gases at room temp Br is a liquid at room temp I & At are solids at room temp

Periodic Trends

Octet Rule Atoms will gain, lose, or share electrons in order to have eight (8) valence electrons. 3 or less valence electrons – atom likely to lose electrons 6 or more valence electrons – atoms likely to gain electrons 4 or 5 valence electrons – atoms likely to share electrons

Periodic Trends Properties of the elements change in a predictable manner across a period and down a group

Atomic Radius The half distance between nuclei of identical atoms that are chemically bonded together

Atomic Radius Tends to decrease as you go across a period Increase nuclear charge pulls electrons closer to the nucleus (decreasing radius) Zeff Tends to increase as you go down a group New electrons are placed in higher energy levels Shielding: core electrons shield outer electrons from pull from nucleus

Ionic Radius Ions – atom or bonded group of atoms that has a positive or negative charge due to a loss/gain of electrons Positive charge  lost electrons Smaller ionic radius compared to anions Negative charge  gained electrons Larger ionic radius compared to cations

Ionic Radius Tends to decrease across a period Tends to increase down a group

Ionic Radius vs. Atomic Radius Metals - the atomic radius of a metal is generally larger than the ionic radius of the same element. Why? Generally, metals loose electrons to achieve the octet. This creates a larger positive charge in the nucleus than the negative charge in the electron cloud, causing the electron cloud to be drawn a little closer to the nucleus as an ion.

Ionic Radius vs. Atomic Radius cont. Non-metals - the atomic radius of a non-metal is generally smaller than the ionic radius of the same element. Why? Generally, non-metals loose electrons to achieve the octet. This creates a larger negative charge in the electron cloud than positive charge in the nucleus, causing the electron cloud to 'puff out' a little bit as an ion. 

Ionic Radius vs. Atomic Radius

Ionization Energy Energy required to remove an electron from a gaseous atom (J) If an atom has a high ionization energy not likely to form a positive ion Tends to increase across a period Tends to decrease down a group

1st Ionization Energy

Electronegativity Relative ability of an atom to attract electrons in a chemical bond Numerical value of 3.98 Paulings or less Fluorine is the most electronegative atoms positioned closer to F have higher electronegativies Tends to increase across period Tends to decrease down a group

Electronegativity

Reactivity vigorously an atom is to react with other substances. Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually determined by two things:

1) How easily electrons can be removed (ionization energy) from an atom

2) or how badly an atom wants to take other atom's electrons (electronegativity)

The transfer/interaction of electrons is the basis of chemical reactions.

Reactivity of Metals Period - reactivity decreases as you go from left to right across a period. Group - reactivity increases as you go down a group Why? The farther to the left and down the periodic chart you go, the easier it is for electrons to be given or taken away, resulting in higher reactivity.

Reactivity of Non-Metals Period - reactivity increases as you go from the left to the right across a period. Group - reactivity decreases as you go down the group. Why? The farther right and up you go on the periodic table, the higher the electronegativity, resulting in a more vigorous exchange of electron.

Electron Affinity Change in energy that occurs when a neutral atom acquires an electron Hopefully atom is become more stable by acquiring an elcectron Measured with a negative value The more negative the value the easier it is to acquire an electron Tends to become more negative across a period Tends to become more positive down a group

Electron Affinity

Summary of Periodic Trends