The Periodic Table & Periodic Law Unit 6 The Periodic Table & Periodic Law

6.1 Development of the Modern Periodic Table The periodic table developed over time as scientists discovered more useful ways to compare and organize the elements

Development of the Periodic Table Antoine Lavoisier – late 1700s Compiled first list of elements known at the time (33 of them) Organized elements into four categories Gases Metals Nonmetals Earths

By 1860 60 elements had been discovered but scientists had no way of organizing them

J.W. Dobereiner classified elements that had similar properties into triads, organizing them by atomic mass

Triads were useful because they grouped elements with similar properties and revealed an orderly pattern in some of their physical and chemical properties which are related to related to atomic mass.

John Newlands – 1864 Noticed when elements were arranged by increasing atomic mass their properties repeated every 8th element Pattern is periodic (repeats in a specific manner) Named this the law of octaves Law did not work for all known elements

Dmitri Mendeleev – 1869 Arranged elements in order of increasing atomic mass into columns with similar properties Predicted the existence and properties of undiscovered elements that were later found

Mendeleev’s Periodic Table

When elements are arranged according to mass, some elements end up in groups with properties different from their own Henry Moseley – 1913 Arranged table by atomic number

Periodic law – when elements are arranged according to increasing atomic number, there is a periodic repetition of chemical and physical properties

The Modern Periodic Table Columns = groups Numbered 1 – 18 Correspond to the number of outermost electrons Have similar properties Some have special names Rows = periods Numbered 1 – 7 Correspond to outermost energy level

Representative elements – in s & p block Transition elements – in d block

What element is in group 4 & period 5? Is this a representative element? What element is in group 18 & period 6?

6.2 Classification of the Elements

Physical States and Classes of Elements Most elements are solid at room temperature Br & Hg are liquid N, O, F, Cl, and Noble gases are gas

Metals Shiny Solid at room temperature Good conductors of heat & electricity Malleable Ductile Found to the left of the staircase 1 exception =

Alkali metals – group 1 elements Except for Very reactive # of valence electrons =

Alkaline earth metals – group 2 elements Also highly reactive # of valence electrons =

Transition metals – group 3 – 12 (d block) Number of valence electrons varies

Inner transition metals – f block Number of valence electrons varies Lanthanide & actinide series

Nonmetals Gases or dull looking solids 1 exception = Poor conductors of heat and electricity Found to the right of the staircase

Halogens – group 17 Highly reactive # of valence electrons =

Nobel gases – group 18 Extremely unreactive # of valence electrons =

Metalloids Border the staircase Have properties of both metals and nonmetals Semi-conductors = conduct electricity only under certain conditions

Valence electrons – electrons in the highest energy level of an atom determine an atoms properties All atoms in same group have same number of valence electrons and therefore…

Valence electrons and period Period an element is in = energy level of its valence electrons Li = Ga=

Valence electrons of representative elements

1 = s1 (1 valence electron) 2 = s2 (2 valence electrons) 13 = s2p1 (3valence electrons) 14 = s2p2 (4 valence electrons) 15 = s2p3 (5 valence electrons) 16 = s2p4 (6 valence electrons) 17 = s2p5 (7 valence electrons) 18 = s2p6 (8 valence electrons)

Transition Metals Outermost s and nearby d sublevels contain electrons All transition metals usually have 1, 2, or 3 valence electrons

Inner Transition Metals Outermost s and nearby f sublevels contain electrons

Valence electrons & ion formation Representative elements lose or gain electrons in order to obtain the same electron configuration as a noble gas How many valence electrons does a nobel gas have?

Octet rule = an atom tends to gain, lose, or share electrons in order to acquire a full set of eight valence electrons

Neutral atom Na 1s22s22p63s1 B 1s22s22p1 P 1s22s22p63s23p5 F 1s22s22p5

Ion Na1+ 1s22s22p6 (Ne) B3+ 1s2 (He) P3- 1s22s22p63s23p6 (Ar) F1-

1 = s1 (1 ve, easily lost, ion formed = ) 13 = s2p1 14 = s2p2 15 = s2p3 16 = s2p4 17 = s2p5 18 = s2p6

6.3 Periodic Trends

Nuclear Charge Proton number Increases as atomic number increases

Shielding Electrons Electrons between the nucleus and valence electrons Shield the valence electrons from the force of the nucleus Total electrons – valence electrons Na S

Shielding Effect The further an energy level is from the nucleus, the more it is shielded from the attraction to the nucleus The outer energy level is shielded by the inner energy levels

Effective Nuclear Charge The nuclear charge that the valence electrons actually feel ENC = total electrons – shielding electrons

Potassium Arsenic Total Valence Shielding ENC Total Valence Shielding

ENC and Group Group 1 Group 2 Group 13 Group 14 Group 15 Group 16 Na Group 2 Mg Group 13 Al Group 14 Si Group 15 P Group 16 S Group 17 Cl Group 18 Ar

Atomic Radius Within period – decreases from left to right More valence electrons = more attracted to nucleus Higher ENC shrinks outer energy level Na vs Cl

Within group – increases from top to bottom More energy levels = bigger atom ENC for all atoms in a group is the same Smallest atom = Largest atom =

Ionic Radius ALWAYS SMALLER THAN NEUTRAL ATOM Cation – ALWAYS SMALLER THAN NEUTRAL ATOM When Cations form the outer most energy level is lost

ALWAYS LARGER THAN THE NEUTRAL ATOM Anions ALWAYS LARGER THAN THE NEUTRAL ATOM When atoms gain electrons they repel each other more

Within period – increases from left to right Anions are larger, where are anions located? Within group – increases from top to bottom More energy levels = bigger atom

Ionization energy Energy required to remove an electron from a gaseous atom Determined by how strongly the valence electrons are attracted to the nucleus Hard to remove electrons from atoms with outer energy level close to nucleus Hard to remove electrons when ENC is large

Within period – increases from left to right More valence electrons = more attraction to nucleus (higher ENC) Within group – decreases from top to bottom More energy levels = valence electrons further from nucleus = less attraction

Electronegativity The ability of an atom to attract electrons in a chemical bond

Within period – increases from left to right Higher ENC attracts electrons more easily Within group – decreases from top to bottom Closer outer energy levels attract electrons more

Highest electronegativity Lowest electronegativity