Unit 3 Notes: Periodic Table Notes John Newlands proposed an organization system based on increasing atomic mass in 1864. He noticed that both the chemical and physical properties repeated every 8 elements and called this the ____Law of Octaves ___________. In 1869 both Lothar Meyer and Dmitri Mendeleev showed a connection between atomic mass and an element’s properties. Mendeleev published first, and is given credit for this. He also noticed a periodic pattern when elements were ordered by increasing ___Atomic Mass _______________________________. By arranging elements in order of increasing atomic mass into columns, Mendeleev created the first Periodic Table. This table also predicted the existence and properties of undiscovered elements. After many new elements were discovered, it appeared that a number of elements were out of order based on their _____Properties_________. In 1913 Henry Mosley discovered that each element contains a unique number of ___Protons________________. By rearranging the elements based on _________Atomic Number___, the problems with the Periodic Table were corrected. This new arrangement creates a periodic repetition of both physical and chemical properties known as the ____Periodic Law___. Periods are the ____Rows_____ Groups/Families are the Columns Valence electrons across a period are There are equal numbers of valence in the same energy level electrons in a group. 1 When elements are arranged in order of increasing _Atomic Number_, there is a periodic repetition of their physical and chemical properties Family (Group): ___Columns (vertical)______; tells the number of electrons in the _Outer___ Energy level, called __Valence Electrons________ (only for representative elements) Period (Series): __Rows (horizontal)____; tells the number of ____Energy Levels__________ an atom has; the number of electrons __Increases__ across a period Representative Elements: Groups __1A through 8A _ (called the s and p blocks) (Columns 1, 2, 13, 14, 15, 16, 17, and 18) Valence Electrons: e- in the ___outer most energy level____; farthest away from the __nucleus (protons)___; the e- with the ___most reactive____ Energy; the e- involved with ___Bonding____ (transferring or sharing) Metals: most of the periodic table, located to the __Left___ of the “stair-step” Properties- good conductors of _heat_ and _Electricity_; they also are __ Malleable___; __ Ductile____; _ High Density, BP and MP_____ Nonmetals: to the Right of the “stair-step”, located in the upper corner of P.T._ Although five times more elements are metals than nonmetals, two of the nonmetals—hydrogen and helium—make up over 99 per cent of the observable Universe Properties- mostly _ Brittle __, but a few _low luster______ and _poor conductors__; they have _ low density, low Melting Point and Boiling Point__ Metalloids: also called _semi-metals__, located _along_ the “stair-step” Properties - __ similar __ to both metals and nonmetals Some metalloids are shiny (silicon), some are not (gallium) Metalloids tend to be brittle, as are nonmetals. Metalloids tend to have high MP and BP like metals. Metalloids tend to have high density, like metals. Metalloids are semiconductors of electricity – somewhere between metals and nonmetals. This makes them good for manufacturing computer chips. 2 Valence electrons Valence electrons the electrons that are in the highest (outermost) energy level that level is also called the valence shell of the atom they are held most loosely The number of valence electrons in an atom determines: The properties of the atom The way that atom will bond chemically As a rule, the fewer electrons in the valence shell, the more reactive the element is. When an atom has eight electrons in the valence shell, it is stable. Our discussion of valence electron configurations leads us to one of the cardinal tenets of chemical bonding, the octet rule. The octet rule states that atoms become especially stable when their valence shells gain a full complement of valence electrons. For example, Helium (He) and Neon (Ne) have eight outer valence electrons in their outer shells which means it is completely filled, so they have a tendency to neither gain or lose electrons. Therefore, Helium and Neon, two of the so-called Noble gases or Inert gases Group # Group Name # of valence electrons 1 Alkali Metals 1 2 Alkaline Earth Metals 2 3-12 Transition Metals 1 or 2 13 Boron Group 3 14 Carbon Group 4 15 Nitrogen Group 5 16 Oxygen Group 6 17 Halogens 7 18 Noble Gases 8 The number of valence electrons increases as you go across the periodic table from left to right. 3 Element Lithium Germanium Sulfur Symbol Li Ge S Group # 1A(1) 4A(14) 6A(16) # of valence e- 1 4 6 Period # 2 4 3 # of E levels 2 4 3 Type of element M ML NM Periodic Trends: 1. Atomic Size - __Decreases__ from left to right across a period (smaller) - __Increases___ from top to bottom down a group (larger) Why? - as you go across a period, (same __energy level__), e- are _added_but _pulled closer to the nucleus___ - as you go down a group, you add ___energy levels___ 2. Ionization Energy: the amount of E needed to _remove _ an electron - __Increases__ from left to right across a period - __Decreases____ from top to bottom down a group Why? 4 - as you go across a period, e- feel stronger attraction from nucleus (protons)___, _Energy___ to remove e-, ____Ionization___ E necessary as you go down a group, __Energy_, _Decreases_ to remove outermost e- because they are further away from the Nucleus (protons) 3. Electronegativity: the tendency for an atom to __attract___ electrons; exclude Noble Gases! - __Increases__ from left to right across a period (except Noble Gases) - __Decreases____ from top to bottom down a group Why? - as you go across a period, e- feel ___more__ attraction from nucleus _Protons_____ to pull in more e- - as you go down a group, more _shielding__ from inner e-, __hinders the nucleus ability__ to attract more e- 4. Ionic Size: Cations:__positive_ ions; metal atoms that ___lose__ electrons 5 - __smaller__ than corresponding neutral atom Why? - __fewer__ e-, so it’s _easier_ for protons to pull in remaining e- Anions:__Negative___ ions; nonmetal atoms that _gain_ electrons - ___larger____ than corresponding neutral atom Why? - _more_ e-, so it’s __harder_ for protons to pull in outermost e- Shielding: The ability of the _inner (lower levels)_ electrons to _shield (reduce)_ the pull of the _protons_ on the _outer (higher levels)__ electrons. “Shielding effect”_increase_ as you add Energy levels (move down a group) Quantum Model Notes Heisenberg's Uncertainty Principle‐ Can determine either the _velocity or the position of an electron, cannot determine both. Schrödinger's Equation ‐ Developed an equation that treated the hydrogen atom's electron as a wave. o Only limits the electron's energy values, does not attempt to describe the electron's path. Describe probability of finding an electron in a given area of orbit. The Quantum Model‐ atomic orbitals are used to describe the possible position of an electron. Orbitals The location of an electron in an atom is described with 4 terms. 6 o Energy Level‐ Described by intergers. The higher the level, the more energy an electron has to have in order to exist in that region. o Sublevels‐ energy levels are divided into sublevels. The # of sublevels contained within an energy level is equal to the integer of the energy level. o Orbitals‐ Each sublevel is subdivided into orbitals. Each orbital can hold 2 electrons. o Spin‐ Electrons can be spinning clockwise (+) or counterclockwise (‐) within the orbital. Periodic Table Activity: Complete the table on page 21 with the information found on pages 18‐20. When complete color each group in a different color in the periodic table. The Periodic Table Notes: Historical development of the periodic table: Highlights Mendeleev (1869): Put the elements into columns according to their properties. Generally ranked elements by increasing atomic mass. Moseley (1911): Periodic table arranged by atomic number Top table: Metals, nonmetals, and metalloids Metals: Explain the electron sea theory, and as you explain each of the properties below, discuss how they are explained by the electron sea theory. Also make sure to explain that these are general properties and may not be true for all metals. o Malleable: Can be pounded into sheets. o Ductile: Can be drawn into wires o Good conductors of heat and electricity o High density (usually) o High MP and BP (usually) o Shiny o Hard Nonmetals: Explain how the bonds between the atoms are highly localized, causing each of the properties below. Again, emphasize that these are general properties and may not be true for all nonmetals. o Brittle o Poor conductors of heat and electricity o Low density o Low MP and BP (many are gases)! Metalloids: The bonding in metalloids is between that of metals and nonmetals, so metalloids have properties of both. o Some metalloids are shiny (silicon), some are not (gallium) o Metalloids tend to be brittle, as are nonmetals. o Metalloids tend to have high MP and BP like metals. o Metalloids tend to have high density, like metals. 7 o Metalloids are semiconductors of electricity – somewhere between metals and nonmetals. This makes them good for manufacturing computer chips. Structure of the periodic table Families/groups (the terms are synonymous and will be used interchangeably) o These are elements in the same columns of the periodic table. o Elements within families/groups tend to have similar physical and chemical properties. o They have similar chemical and physical properties because they have similar electron configurations. Example: Li = [He] 2s1, Na = [Ne] 3s1 – each has one electron in the outermost energy level. o Explain that s‐ and p‐electrons in the outermost energy level are responsible for the reactions that take place.