Unit 4: The Periodic Table How is the periodic table a useful tool?

Introduction to the Periodic Table

Johann Döbereiner : Arranged elements into groups of three (triads)

John Newland : the “Law of Octaves” When elements were arranged in order of increasing atomic weight, any one of these elements showed properties similar to those of the elements 8 places ahead and 8 places behind

Dmitri Mendeleev : published a table of the elements (63 elements at this time) Arranged the elements in order of increasing atomic mass

The Modern Periodic Table 1 In the modern periodic table, elements are arranged in order of increasing atomic number.

Groups vs. Periods 1 Groups: Vertical columns All elements in the same group have the same number of valence e- Lewis Dot Diagrams: represent the configuration of ONLY valence electrons Remember- “2 on top, around the clock” Periods: Horizontal rows All elements in the same period have the same number of occupied principle energy levels (PELs)

Types of Elements 1 Metals Nonmetals Metalloids

Properties of Metals 1 Good conductors of heat and electricity Have luster (shiny) Solid at RT, except for Hg (liquid at RT) Ductile: can be drawn into wires Malleable: can be hammered into thin sheets without breaking High melting points High density

Nonmetals 1 Most are gases at room temp A few are solids (ex. sulfur and phosphorus) Bromine is a liquid at RT Look dull NM solids are poor conductors NM solids brittle and hard Have low densities Have lower melting points

Metalloids (Semi-metals) 1 Have some properties of metals and some properties of nonmetals B, Si, As, Te, Ge, Sb

Atoms vs. Ions 1 ATOMS are NEUTRAL; protons = electrons IONS are charged particles protons DO NOT equal electrons because electrons are lost or gained There are two types of ions….. Cations: ions with a positive (+) charge Have lost electrons, so they become positive Anions: ions with a negative (-) charge Have gained electrons, so they become negative

Ions, cont. 1 Metals tend to LOSE electrons and become positively charged Non-metals tend to GAIN electrons and become negatively charged Noble Gases do not gain/lose electrons because they are stable

Groups of the PT 1 Group 1: Alkali Metals Have one valence electron Highly reactive with water Lose one electron to become a +1 charged ion

Groups of the PT 1 Group 2: Alkaline Earth Metals Have two valence electrons Reactive with water Lose two electrons to become a +2 charged ion **Elements in Groups 1 and 2 are so reactive that they are never found alone in nature (only found as compounds)

Per.5 – Monday, 11/23 1 Periodic Table test tomorrow Online Review HW 11:19 am tomorrow Full packet key will be posted and PP will be posted on shscience.net I will be after school starting at 3 pm

Transition Metals 1 Groups 3-12: Transition Metals Can have multiple oxidation states Colored ions must be transition metals

Inner Transition Metals 1 Lanthanide series: an extension of period 6 Actinide series: an extension of period 7

The Halogens 1 Group 17- The Halogens Have 7 valence electrons Gain one electron to become -1 charged ion

The Noble (Inert) Gases 1 Group 18- Noble (Inert) Gases Stable/ Non-reactive Have a complete valence shell of 8 e- (except He which has 2 valence e-) The Octet Rule: All other elements want a full valence shell of 8 e- like the noble gases

Phases of Elements at STP 1 Two Liquids: Hg and Br Gas: H, N, O, F, Cl, and Noble Gases The rest are solids

What trends exist within the PT? 1 There are trends as you go DOWN a group and ACROSS a period

Atomic Radius 1 Atomic radius: distance from the nucleus to the outermost occupied energy level n=2 n=1 n=3

Ionization Energy 1 Ionization Energy: the amount of energy it takes to remove an electron from an atom Units: kJ/mol

Electronegativity 1 Electronegativity: the ability of an atom to attract electrons toward itself Based on the Pauling Scale ( )

Group Trends 1 As you go DOWN a group: 1) Atomic radius increases Why? Increase in the number of occupied PELs 2) Ionization energy decreases Why? The shielding effect Kernel (inner) electrons reduce the pull the nucleus has for the outermost electrons 3) Electronegativity decreases Why? The shielding effect Shielding reduces the ability for the nucleus to attract electrons toward itself

Atomic Radius Trends 1 ElementAtomic Radius (pm) Li130 Na160 K200 Rb215

Ionization Energy Trends 1 ElementIonization Energy (kJ/mol) Li520 Na496 K419 Rb403

Electronegativity Trends 1 ElementElectroneg. Li1.0 Na0.9 K0.8 Rb0.8

Period Trends 1 As you go ACROSS a period: 1)Atomic radius decreases Why? Increasing nuclear charge pulls electrons closer 2)Ionization energy increases Why? Increasing nuclear charge makes it harder to pull electrons off 3) Electronegativity increases Why? Increasing nuclear charge makes it easier to attract e-

Atomic Radius Trends 1 ElementAtomic Radius (pm) Li130 Be99 B84 C75

Ionization Energy Trends 1 ElementIonization Energy (kJ/mol) Li520 Be900 C1086 N1402

Electronegativity Trends 1 ElementElectronegativity Li1.0 Be1.6 C2.6 N3.0

Trends in Metallic Character 1 Metallic character increases as you go down a group Metallic character decreases as you go across a period

1 Increasing Electronegativity

Summary: Group Trends 1 As you go down a group: Atomic radius increases Ionization energy decreases Electronegativity decreases Why? The shielding effect

Summary: Period Trends 1 As you go across a period: Atomic radius decreases Ionization energy increases Electronegativity increases Why? Nuclear charge effect