Chapter 6 and 7 notes

6.1 Development of the Modern Periodic Table objectives Trace the development and identify key features of the periodic table

It is called the periodic table because properties of the elements in the table repeat in a periodic way

History of the periodic table’s development 1790’s Antoine Lavoisier compiled a list of elements known at the time, 23 elements, including silver, gold, carbon and oxygen.  

First table published in 1869 by Dmitri Mendeleev (1834-1907) – Arranged in rows (periods) of increasing atomic mass – Arranged in columns (groups or families) by similarities in physical and chemical properties – Holes were left in the table to allow for the discovery of new elements

Modified in 1913 by Henry Moseley (1887-1915) into the modern Periodic Table – Arranged in rows (periods) of increasing Atomic Number – that is, increasing number of protons – Arranged in columns (groups or families) by repetition of physical and chemical properties The group number represents the energy level being occupied

The periodic table has been continually refined over time as new elements have been discovered

Glenn Seaborg: Identified the Lanthanide and Actinide Series while working on the Manhattan Project during World War II. He is credited with the discovery of 8 new elements

We are now up to 117 elements.   2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p We are now up to 117 elements.

Moseley explains his table in term of the Periodic Law – The chemical and physical properties of the elements are periodic functions of their atomic numbers; properties of the elements occurred at repeated intervals called periods

Periodic law: the properties of elements repeat every so often Period: horizontal row; (7)- same energy level Group (family): vertical column; (18) same number of valence electrons

Describing the Periodic Table periodic law: the properties of elements repeat every so often period: horizontal row; there are 7 group (family): vertical column; there are 18 18 17 16 15 14 13 12 9 10 11 8 7 6 5 4 1 3 2 1 2 3 4 6 7 5

Regions of the Table Metals: left side of Table; form cations Properties: good conductors of heat and electricity, shiny when clean, solid at room temperature, malleable and ductile   Nonmetals: right side of Table; form anions Properties: good insulators. gases or brittle solids Metalloids (semimetals): “stairs” between metals and nonmetals Properties: in-between those of metals and nonmetals “semiconductors” Si and Ge  computer chips

Regions of the Table (cont.) metals nonmetals

Alkali metals – Group 1: Li, Na, K, Rb, Cs, Fr most reactive group (combine with other elements to become more stable) have 1 electron in an s orbital solid, silvery metals +1 charge   Alkaline-earth metals – Group 2: Be, Mg, Ca, Sr, Ba, Ra second most reactive group of metals have 2 electrons in an s orbital +2 charge

Transition elements – elements in Groups 3 – 12 many of the most commonly recognized metals are in these groups Groups 13-16 contain metals, nonmetals, and all of the metalloids characteristics depend more on their period than on their group Halogens – Group 17: F, Cl, Br, I, At most reactive nonmetals because they only need 1 electron nonmetals that react with metals to form salts -1 charge

Noble gases – Group 18: He, Ne, Ar, Kr, Xe, Rn all of the elements in this group are gases all of the elements in this group are nonmetals least reactive elements   Lanthanide &Actinide series – belong to Periods 6 and 7 called inner transition metals many are radioactive (have an unstable arrangement of protons and neutrons in the nucleus)

main block (representative) elements: groups 1, 2, 3A-8A.

Modern Periodic table – The table consists of boxes, each containing an element name, symbol, atomic number and atomic mass. -The boxes are arranged in order of increasing atomic number into a series of columns called groups or families and rows called periods. -Beginning with hydrogen in period 1 through 8, followed by the letter A or B. The groups designated with an A are often referred to as the main group or representative elements because they posses a wide range of chemical and physical properties. The groups designated with a B are referred to as the transition elements

6.2 Classification of the elements objectives 1. explain why elements in the same group have similar properties 2. identify the four blocks of the periodic table based on electron configuration

Shortcut to finding valence electrons (does not work for transition metals): Group 1 – 1 valence electron Group 2 – 2 valence electrons Group 3A – 3 valence electrons Group 4A – 4 valence electrons Group 5A – 5 valence electrons Group 6A – 6 valence electrons Group 7A – 7 valence electrons Group 8A – 8 valence electrons  

Organizing the elements by electron configuration   Valence Electrons ** atoms in the same group have similar chemical properties because they have the same number of valence electrons. Electrons available to be lost, gained, or shared in the formation of chemical compounds. Found in the s and p orbitals of the highest energy level. Often located in incompletely filled energy levels. To find the number of valence electrons, underline the largest number as often as it occurs and add the superscripts.

Example: Cl – 1s2, 2s2, 2p6, 3s2, 3p5 – 7 valence electrons Example: Mg - 1s2, 2s2, 2p6, 3s2 – 2 valence electrons   Example: Kr – 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6 – 8 valence electrons Example: U – 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6, 6s2, 4f14, 5d10, 6p6, 7s2, 5f4 – 2 valence electrons

Note: The number of valence electrons for the transition elements can have different values based on the conditions of chemical reactions. This is also true for a small number of the metals in Groups 3A-6A

S block – Groups 1A and 2A and hydrogen and helium S block – Groups 1A and 2A and hydrogen and helium. In this block the valence electrons only occupy the s orbitals. Group 1A elements have a partially filled s orbital containing one valence electron. Group 2A have a completely filled s orbital P block – after the s level is filled the valence electrons next occupy the p sublevel. The p block is comprised of groups 3A-8A. The three p orbitals can hold a maximum of 6 electrons. Group 8A or the noble gases are unique because of their stability. Both the s and p orbitals corresponding to the principal energy level are filled. D block – contains the metals and is the largest block. With several exceptions, d block elements are characterized by a filled outermost s orbital and filled or partially filled d orbital. The 5 d orbitals can hold a total of ten electrons. F – block – contains the inner transition metals. The f block is characterized by a filled or partially filled outermost s orbital and filled or partially filled 4f and 5 f orbitals. The 7 f orbitals can hold up to 14 electrons.

Orbitals Being Filled Groups 1 8 2 1s 1 3 4 5 6 7 1s 2s 2 2p 3 3s 3p 3 4 5 6 7 1s 2s 2 2p 3 3s 3p Periods 4s 3d 4p 4 4d 5p 5 5s La 5d 6p 6 6s Ac 6d 7 7s 4f Lanthanide series 5f Actinide series Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 345

** Thus the s, p, d, and f blocks determine the shape of the periodic table. As you proceed down through the periods, the principal energy level increases as does the number of energy sublevels containing electrons Period 1 contains only s block elements, periods 2 and 3 contains both s and p block elements. Periods 4 and 5 contains s, p, and d block elements and periods 6 and 7 contains s,p,d, and f block elements

6.3 periodic trends 1. compare period and group trends of several properties 2. relate period and group trends in atomic radii to electron configuration

Periodic Law – The chemical and physical properties of the elements are periodic functions of their atomic numbers; properties of the elements occurred at repeated intervals called periods. This defines the property of periodicity  

while there are many periodic trends, we will focus on properties that show patterns when examined across the periods or vertically down the groups while there are many periodic trends, we will focus on atomic radii (the plural of radius) ionization energy electronegativity

Distance between nuclei Atomic Radii – one half the distance between the nuclei of identical atoms that are bonded together. decreases across periods increases down groups Distance between nuclei atomic radii increases

Atomic Radii = 1 Angstrom IA IIA IIIA IVA VA VIA VIIA Li Be B C N O F 1.52 1.11 1.86 1.60 2.31 1.97 2.44 2.15 2.62 2.17 0.88 0.77 0.70 0.66 0.64 1.43 1.17 1.10 1.04 0.99 1.22 1.22 1.21 1.17 1.14 1.62 1.40 1.41 1.37 1.33 1.71 1.75 1.46 Na Mg Al Si P S Cl K Ca Ga Ge As Se Br Rb Sr In Sn Sb Te I Cs Ba Tl Pb Bi = 1 Angstrom

Atomic Radius The trend for atomic radius in a vertical column is to go from smaller at the top to larger at the bottom of the family. Why? With each step down the family, we add an entirely new energy level to the electron cloud, making the atoms larger with each step.

Atomic Radius The trend across a horizontal period is less obvious. What happens to atomic structure as we step from left to right? Each step adds a proton and an electron (and 1 or 2 neutrons). Electrons are added to existing energy levels

Atomic Radius The effect is that the more positive nucleus has a greater pull on the electron cloud. The nucleus is more positive and the electron cloud is more negative. The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period.

Effective Nuclear Charge What keeps electrons from simply flying off into space? Effective nuclear charge is the pull that an electron “feels” from the nucleus. The closer an electron is to the nucleus, the more pull it feels. As effective nuclear charge increases, the electron cloud is pulled in tighter.

Ionization Energy – the energy required to remove one electron from a neutral atom of an element. increases across periods decreases down groups as nuclear charge increases, ionization energy increases * octet rule, states that atoms tend to gain , lose, or share electrons in order to acquire a full set of eight valence electrons  

Ionization energy increases

Electronegativity – a measure of the ability of an atom in a compound to attract electrons increases across periods decreases down groups

electronegativity: the tendency for a bonded atom to attract e– to itself Linus Pauling quantified the electronegativity scale. As we go , electronegativity… decreases. As we go , electronegativity… increases. electronegativity increases

Ionic Radius Metallic ions lose electrons and are smaller compared to the atoms from which they come from. For ex., Na+ is smaller than Na since it loses an electron! Nonmetallic ions gain electrons and are LARGER compared to the atoms they come from. Cl- is larger than Cl

Ionic Radius Cations are always smaller than the original atom. The entire outer electron is is removed during ionization. Conversely, anions are always larger than the original atom. Electrons are added to the outer electron level.

Ionic radii

Metallic Character This is simple a relative measure of how easily atoms lose or give up electrons.

Summary of Periodic Trends Shielding is constant Atomic radius decreases Ionization energy increases Electronegativity increases Nuclear charge increases 1A Ionization energy decreases Electronegativity decreases Nuclear charge increases Atomic radius increases Shielding increases Ionic size increases 2A 3A 4A 5A 6A 7A Ionic size (cations) Ionic size (anions) decreases decreases

SIMPLY PUT: M E R I Metal reactivity electronegativity radii (Atomic) ionization energy

CHAPTER 7: THE ELEMENTS 7.1 Properties of s-block elements Objectives explain how elements in a given group are both similar and different Discuss the properties of hydrogen Describe and compare the properties of alkali metals and alkaline earth metals  

Elements in any given group on the periodic table have the same number of valence electrons. The number and location of valence electrons determine the chemistry of an element. Thus, elements within a group have similar physical and chemical properties

Representative elements display the range of possible valence electrons from one in group 1A to eight in group 8A. The valence electrons of representative elements are in s or p orbitals.

Although elements within a group have the same number of valence electrons, they have different numbers of nonvalence electrons. As new levels of electrons are added, the atomic radius increases and the shielding effect increases. As a result, the ionization energy decreases. A lower ionization energy makes it easier for an element to lose electrons.

Remember, that metals tend to lose electrons Remember, that metals tend to lose electrons. Thus, the lower the ionization energy the more reactive the metal. For a group of metals, reactivity increases as the atomic number increases. The opposite is true for nonmetals, because non metals tend to gain electrons. For a group of nonmetals, reactivity decreases as the atomic number increases

Diagonal relationships are the close relationships between elements in neighboring groups

Hydrogen – Hydrogen is placed in group 1A because it has one valence electron. This does not mean it has the same properties as the metals in group 1A. It has metallic and nonmetallic properties and is not considered part of any group. Remember the Hindenburg. It burned with the hydrogen that kept it afloat came in contact with oxygen.

The universe contains more than 90% hydrogen by mass. There are 3 naturally occurring hydrogen isotopes: protium, deuterium and tritium. 99.985% of hydrogen is protium which has no neutrons

Hydrogen’s single valence electron explains its unusual set of metallic and nonmetallic properties. When a hydrogen atom acts like a nonmetal, it gains an electron and achieves the stable electron configuration of helium. When it loses its electron it forms a hydrogen ion

Hydrogen is produced in a lab when a metal reacts with an acid or when electricity is use to separate water into hydrogen and oxygen. Large quantities are produced when water reacts with methane, which is the main ingredient in natural gas.