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Prior Knowledge: An electron configuration tells us where the electrons are in an atom. Example: Ca: 1s22s22p63s23p64s2 ↓ 2 1s ↓ Energy level the orbital are in How many e- in the orbital Shape of the orbital Valence Electrons (Outermost electrons/Why we learn e- configurations) Notice that in electron configurations, the last two s & p orbitals are in the furthest energy level from the nucleus. If you add up the electrons in the outer most s & p orbitals you get a number between 1-8 These electrons are “Valence Electrons” (e-’s in the outer-most e.l.) Valence Electrons: - are responsible for chemical bonding - All of the other electrons are called “core electrons”. - Help tell us see some physical and chemical properties - Helps predict the charges on ions Valence Electrons: Counting Valence Electrons Group A # = number of valence electrons (only exception Helium = 2 e-’s) Examples: Ca = 2 e-’s Nitrogen = 5 e-’s Argon = 8 e-’s Si: 1s22s22p63s23p2 outer most energy level: 3 s & p orbital e-’s: 2 and 2 Valence electrons: 2+2= 4 Cl: outer most energy level: s & p orbital e-’s: Valence electrons: Octet Rule (and ions) The s & p orbitals allow for a total of 8 electrons. When the outer-most s & p are full (8e-) the atom is stable. -atoms will gain or lose electrons in order to reach a “full octet” (this makes the atoms an ion) Predictable Ions (Groups 1A-8A): 1. Locate the nearest noble gas 2. Count how many “places” it is away 3. If you count backwards to the left you have a + charge if you count forward, you have a negative charge Practice Problems: How many electrons are gained or lost when forming an ion from the following elements? a) Magnesium: ____ (gained or lost) b) Iodine: ____ (gained or lost) c) Gallium:____ (gained or lost) d) Boron:____ (gained or lost) Practice Problems: How many electrons are gained or lost when forming an ion from the following elements? 1. Magnesium: ____ (gained or lost) 2. Iodine: ____ (gained or lost) 3. Gallium:____ (gained or lost) 4. Boron:____ (gained or lost) Expressing Valence Electrons Electron dot structures represent electrons as dots located around the symbol of the element in the pattern shown below. 4 7 3 8 X 6 Examples: Nitrogen = N 1 5 2 Hydrogen = (important exception.... Helium = He ) H Valence Electrons On The Periodic Table 1 8 2 3 4 5 6 7 The Modern Periodic Table Dmitri Mendeleev constructed the 1st periodic table Mendeleev: -left blank spaces for “missing elements”. Later when these elements were discovered, the gaps could be filled. - arranged the elements in columns and rows according to their properties. Elements with similar properties were in the same horizontal row. -was able to accurately predict the properties of the missing elements based on the properties of the elements in similar rows. - ordered the elements by increasing atomic mass The Modern Periodic Table - In 1913, Henry Moseley determined the atomic number, (# of p+), of the elements. - He then arranged the elements in the periodic table by increasing atomic number. - This switched the position of some elements. This is how the modern periodic table is arranged today. -In 1942 Glenn T. Seaborg rearranged the periodic table to include the 5f row. He won a Nobel Prize after publishing his new periodic table. Features of the Modern Periodic Table Periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number is called periodic law. The Modern Periodic Table Columns of elements are called groups. Rows of elements are called periods. Elements in groups 1,2, and 13-18 possess a wide variety of chemical and physical properties and are called the representative elements. Elements in groups 3-12 are known as the transition metals. The Modern Periodic Table Elements are classified as metals, non-metals, and metalloids. Metals are elements that are generally • shiny when smooth and clean • solid at room temperature • good conductors of heat and electricity • Malleable The Modern Periodic Table Alkali metals are all the elements in group 1 except hydrogen, and are very reactive. Alkaline earth metals are in group 2, and are also highly reactive. The transition elements are divided into transition metals and inner transition metals. The two sets of inner transition metals are called the lanthanide series and actinide series and are located at the bottom of the periodic table. The Modern Periodic Table Non-metals are elements that are generally: • gases or brittle • dull-looking solids • poor conductors of heat and electricity. Group 17 is composed of highly reactive elements called halogens. Group 18 gases are extremely unreactive and commonly called noble gases. The Modern Periodic Table Metalloids have physical and chemical properties of both metals and non-metals: • Semiconductors -(which can be used in computer chips) Examples: Silicon & Germanium Figure 11.35: Classification of elements as metals, nonmetals, and metalloids. Parts of the Periodic Table Trends on the Periodic Table The Periodic Table is organized so that it follows many trends: Trends you have learned: Trends you will learn in the next few slides: •Increasing Atomic Number •Increasing valence electrons across row •Groups have similar properties •Atomic Radius •Atomic Radius vs Ion radius •Ionization Energy •Electronegativity Trends: Atomic Radius Atomic size is a periodic trend influenced by electron configuration. Atomic radius is the distance from the nucleus to the outer most electrons. Atomic Radius There is a general decrease in atomic radius from left to right: • caused by increasing positive charge in the nucleus. • As you add more electrons and protons in the same energy level, it causes a stronger attractions which shrinks the e- cloud. As you move down a group the size generally increases: • Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons. Atomic Radius Atomic radius generally increases as you move down a group. The outermost orbital size increases down a group, making the atom larger. Figure 11.36: Relative atomic sizes for selected atoms. Ionic Radius An ion is an atom or bonded group of atoms with a positive or negative charge. When atoms lose electrons and form positively charged ions, they always become smaller for two reasons: 1.The loss of a valence electron can leave an empty outer orbital resulting in a small radius. 2.Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius. Ionic Radius When atoms gain electrons, they can become larger, because the addition of an electron increases electrostatic repulsion. Ionic Radius Both positive and negative ions increase in size moving down a group. Ionization Energy Ionization energy is the energy required to remove the outer most electron in an atom. •The energy required to remove the first electron is called the first ionization energy. •As you move down a group less energy is required to remove electrons. Therefore, ionization energy decreases. It is easier to remove an electron if it is farther away from the nucleus. These e-’s are not as attracted to the nucleus. In general, the larger the atom, the less attracted it is to its e-’s. Ionization Energy The ionization at which the large increase in energy occurs is related to the number of valence electrons. • First ionization energy increases from left to right across a period. • First ionization energy decreases down a group because atomic size increases and less energy is required to remove an electron farther from the nucleus. Ionization Energy Ionization Energy Ionization Energy Discrepancies When an orbital is half full the ionization energy increases a little bit. When an orbital is full the ionization energy increases When the energy level is full the ionization energy is the highest Ionization Energy The octet rule states that atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons. The octet rule is useful for predicting what types of ions an element is likely to form. Electronegativity The electronegativity of an element indicates its relative ability to attract electrons in a chemical bond. Electronegativity decreases down a group and increases left to right across a period. As we move down a group, electronegativity decreases. -This occurs because the electron is farther away from the nucleus. As we move across a period, electronegativity increases. -This is because the atoms get smaller moving across the period and the electrons are more strongly attracted to the nucleus