Chapter 13 Chemical Periodicity

Introduction In the 19th century, chemists began to categorize the elements according to similarities in their physical and chemical properties The end result of this was the modern periodic table The periodic table is very useful for understanding and predicting the properties of elements

13.1 The Development of the Periodic Table Newland, an English chemist, published list of elements arranged in order according to their increasing atomic mass. He stated that the elements properties repeated when they were arranged according to increasing atomic mass in groups of eight He called this the arrangement the law of octaves Similar to musical scale that repeats every eighth note Law only works up to Ca

13.1 The Development of the Periodic Table Mendeleev, a Russian chemist, refined and added to the arrangement of elements in a table according to their atomic masses With this arrangement he noticed a regular (periodic) recurrence of their physical and chemical properties

13.1 The Development of the Periodic Table Mosely rearranged the periodic table according to the atomic number of the elements, which is how the modern periodic table is arranged today The periodic table is a valuable organizational tool for chemists

13.2 The Modern Periodic Table The most commonly used modern periodic table, sometimes called the long form (your table) The long form table lists many properties of the elements so that the chemist can check them at a glance

13.2 The Modern Periodic Table The periodic law states that when elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties The horizontal rows of the periodic table are called periods – there are 7 periods in the periodic table

13.2 The Modern Periodic Table The vertical columns are called groups or families identified by number and a letter Groups 1A through 7A and group 0 make up the representative elements (wide variety of properties) Group B elements are the transition metals Two rows of elements below the periodic table are the lanthanides and actinides

13.3 Electron Configurations and Periodicity The electron configuration of an element plays the greatest part in determining it’s physical and chemical properties Most elements within the same group have the identical electron configurations in their outer most energy level (valence level) Elements are classified into 4 different categories according to their valence (outer) electron configuration Noble Gases, Representative Elements, Transition Metals, Inner Transition Metals

13.3 Noble Gases (Group 0) 1. Noble gases (group 0) are elements in which the outermost s and p sublevels are filled Also called inert gases because they do not react with other elements – they are stable on their own Helium has 2 valence electrons (full 1s sublevel) The rest of the noble gases have 8 valence electrons (full s and p sublevels): Ne, Ar, Kr, Xe, Rn

13.3 Representative Elements 2. Representative elements (Group A) Elements whose outermost s or p sublevels are only partially filled Group 1A are known as the alkali metals 1 electron in outermost energy level Very reactive → only in compounds in nature

13.3 Representative Elements 2. Representative elements (Group A) continued Group 2A are known as the alkaline earth metals 2 electrons in outermost energy level Also reactive (but not as reactive as 1A) → only in compounds in nature Group 7A are known as the halogens 7 electrons in outermost energy level Nonmetals that are highly reactive Also called salt formers

13.3 Representative Elements For any representative element, the group number is equal to the number of electrons in the outermost energy level See periodic table 354 –355

13.3 Transition Elements 3. Transition Metals (Group B) Elements whose outermost s sublevel and nearby d sublevel contain electrons The d sublevels overlap with s sublevels – this is why they are transition elements Characterized by having electrons added to the d orbitals Not as reactive as Group A elements

13.3 Inner Transition Metals Elements whose outermost s sublevel and nearby f sublevel generally contain electrons Characterized by the filling of the f orbitals

13.3 Electron Configurations and Periodicity The periodic table can be divided into sections, which correspond to the sublevels that are filled with electrons (on your table) (blocks) Group1A and 2A are in the s block (also Helium) valence level = period # Group 3A, 4A, 5A, 6A, 7A, and 0 belong to p block

13.3 Electron Configurations and Periodicity Transition belong to d block Exception – d sublevel is one less than period # Inner transition belong to f block Exception – f sublevel is 2 less than period # The valence electron configurations can be determined by using the block diagram in figure 13.4 – on your periodic table

13.4 Periodic Trends in Atomic Size Remember that, according to the quantum mechanical model, an atom does not have a specifically defined boundary that sets the limit of its size.

13.4 Periodic Trends in Atomic Size However, there are ways to estimate the relative sizes of atoms. X-ray diffraction – estimates the size of atoms in crystalline solids The distance between the nuclei of diatomic molecules (examples: O2 or Br2) can be used to estimate the atomic radius of an atom. atomic radius – half the distance between the nuclei of two like atoms

13.4 Atomic Size – Group Trends Atomic size generally increases as you move down a group of the periodic table The size increases because electrons are added to higher principle energy levels The added charge of nucleus pulls electrons inward, but the net effect is an increase in size because electrons are further from nucleus

13.4 Atomic Size – Periodic Trends Atomic size generally decreases as you move from left to right across a period The size decreases because electrons are added to the same principle energy level, but the added charge of nucleus pulls electrons inward; the net effect is a decrease in size This trend is less pronounced in periods where there are more electrons in the occupied principle energy levels between the nucleus and the outermost electrons; this is referred to as the shielding effect

13.5 Periodic Trends in Ionization Energy When an atom gains or loses an electron it forms an ion. The energy that is required to overcome the attraction of the nuclear charge and remove an electron from a gaseous atom is called the ionization energy The first ionization energy is the amount needed to remove the first outermost electron The second ionization energy is the amount needed to remove the next outermost electron The third ionization energy is the amount to remove the third and so on (Table 13.1 page 362)

13.5 Periodic Trends in Ionization Energy Ionization energies can be used to predict how many electrons an atom will gain or lose in a chemical reaction 1A vs. 2A – Table 13.1 page 362 Two factors affect ionization energy: nuclear charge and distance from the nucleus

13.5 Ionization Energy – Group Trends In general, the first ionization energy decreases as you move down a group on the periodic table. The size of the atoms increases as you move down; thus the outermost electron is farther from the nucleus and more easily removed This results in a lower ionization energy

13.5 Ionization Energy – Periodic Trends For the representative elements, the first ionization energy generally increases as you move from left to right across a period. The nuclear charge is increasing and the atomic size is decreasing, therefore there is more of an attraction between the nucleus and the outermost electron This results in a higher ionization energy

13.6 Trends in Ionic Size When atoms lose electrons they become positive ions (cations) Cations are always smaller than the atoms from which they are formed There is a stronger attraction between the nucleus (same number of protons) and the remaining electrons (fewer)

13.6 Trends in Ionic Size When atoms gain electrons they become negative ions (anions) Anions are always larger than the atoms from which they are formed There is less of an attraction between the nucleus (same number of protons) and the resulting electrons (more)

13.6 Trends in Ionic Size Periodic Trend – There is a decrease in the size of cations as you move across a period from left to right – when you get to group 4A the anions (which are much larger) start to decrease in size Group Trend – Ionic size (both cations and anions) increases as you go down each group.

13.7 Trends in Electronegativity The electronegativity of an element is the tendency for the atoms of the element to attract electrons when they are chemically combined with another element The Pauling scale uses arbitrary units to express the electronegativity of the all elements (except noble gases) The Pauling scale is based on a number of factors including ionization energies and electron affinities

13.7 Trends in Electronegativity Periodic Trend – As you go across a period from left to right, the electronegativity of the representative elements increases Metallic elements far left have low electronegativities Nonmetallic elements far right have high electronegativities Group Trend - Electronegativity generally decreases as you go down a group *Transition metals do not show as regular trend of electronegativity