chem

• Dmitri Mendeleev 

  • In 1869, Mendeleev and Lothar Meyer (Germany) published nearly identical classification schemes for elements known to date. 
  • The periodic table is based on the similarity of properties and reactivities exhibited by certain elements.
• Henry Moseley (1914)

  • revision of the periodic table by rearranging the elements by their ATOMIC NUMBER. 
  • number of protons = atomic number
• Elements.

  • originally Aristotle’s theory of Air, Water, Fire, and Earth. 
  • Scientists have identified 90 naturally occurring elements and created 28 others
  • The elements make up our bodies, the world, the sun, and the entire universe
• Valence Electrons

  •  Valence electrons are the electrons in the outer energy level of an atom.
  •  These are the electrons that are transferred or shared when atoms bond together
• Metals

  • Conductors of heat and electricity
  • Shiny, ductile, malleable
• Non-Metals

  • Poor conductors
  • Not ductile nor malleable
  • Solid of these are brittle and can break easily
  • Dull
• Metalloids

  • Metal like (metal + nonmetal)
  • Solids that can be shiny or dull
  • Conducts heat and electricity but not as well as metals
  • Ductile and malleable
• FAMILIES

  • similar but not identical properties. 
  • same # of valence electron
• PERIODS (row)

  • not alike in properties
  • 1st element solid-like
  • last elements gas-like
• What does it mean to be reactive? Elements that are reactive bond easily with other elements to make compounds
• Reactive element


  Has an incomplete valence electron level
• Hydrogen

  • its own family
  • It has one proton and one electron
  • Only needs 2 electrons to fill up its valence shell
• Alkali Metals

  • found in the first column of the periodic table
  • have a single electron. Group 1A
  • most reactive
  • gives valence e
• Alkaline Earth Metals

  • uncombined in nature (madaling mag react)
  • 2 valence electrons. Group 2A
• Transition Metals

  • Elements have 1 or 2 valence electrons, which they lose when they form bonds with other atoms.
  • Some of its elements can lose electrons in their next-to-outermost level
  • mostly metals
  • brightly colored when grinded
• Boron Family

  • 3 Valence Electrons. Group 3A
• Carbon Family

  • 4 valence electrons. Group 4A
  • This family includes a non-metal (carbon), metalloids, and metals.
  • kakabit or kakabitan
  • carbon = “basis of life” (organic chem)
• Nitrogen Family

  • This family includes nonmetals, metalloids, and metals. 
  • 5 valence electrons. Group 5A
  • tagatanggap ng electrons
  • nitrogen = 78% of the atmosphere
• Oxygen Family

  • 6 valence electrons. Group 6A
  • Most elements in this family share electrons when forming compounds
• 
  

  Noble Gases

  • They are inactive because their outermost energy level is full. 
  • Because they do not readily combine with other elements to form compounds these are called inert. 
  • colorless
• Halogen Family

  • 7 valence electrons (Group 7A), 
  • most active non-metals.toms only need to gain 1 electron to fill their outermost energy level
  • react with alkali = salt
  • does not exist on its own
• Rare Earth Elements

  • One element of the lanthanide series and most of the elements in the actinide series are called trans-uranium, which means synthetic or man-made
• Electron Configuration 

  • Arrangements of electrons in an atom. 
• Aufbau’s Principle (bottom up)

  • Electrons fill orbitals starting from the lowest energy orbital, then proceed to fill each lower energy orbital, one electron at a time, before filling a higher energy level.
• Pauli Exclusion Principle

  • An atomic orbital can contain at most two electrons with opposite spins. It may be clockwise or counter-clockwise. Electrons spins are represented by arrows up or down. An orbital with two electrons is written as 
• Hund’s Rule

  • When electrons occupy orbitals of equal energy, one electron enters each orbital until all degenerate orbitals contain one electron with parallel spins     
• Periodic Trends

  • The properties of an element are determined largely by the electron configuration of the outermost electrons, and by how far away those electrons are from the nucleus. 
• Atomic Radius

  • Is difficult to measure because an atom does not have a precise outer boundary. What can be measured is the distance between the nuclei of atoms (internuclear distance)
  • Note that from left to right across a period of elements, it decreases. This is so because as you go through a period, the energy level remains the same as electrons are being added
• Ionic Radius

  • Metals readily lose electrons to form positive ions. The loss of outermost electrons results in increased attraction by the nucleus for the remaining electrons. As a result, the cations are smaller than the atoms from which they are formed.
• Ionization Energy

  • When an atom loses or gains electrons to form ions, the process is called ionization. The energy a gaseous atom absorbs to remove an electron in its ground state is called ionization energy (1).
  • The symbol I, stands for the first ionization energy-the energy required to remove or electron from a neutral gaseous atom. I, is the second ionization energy-the energy requires to remove an electron from a gaseous ion with a charge of +1.
• Electron Affinity

  • amount of energy released when an electron is added to a gaseous atom. 
  • example: energy is released when a chlorine atom gains an electron to become a negative ion (anion).
• Electronegativity

  • ability of an atom in a compound to attract electrons to itself. 
  • The larger the electronegativity value, the greater the tendency of the atom to attract electrons.
• Chemical Bonding

  • Almost everything that you can see or touch is a result of a chemical bond.
  • Atoms bonded to one another produce chemical compounds.
  • Chemical bonding happens when there is a transfer or sharing of electrons.
  • 2 major types of chemical bonds:
• IONIC (METAL + NONMETAL)

  • electrons are transferred between valence shells of atoms 
  • Ionic compounds are made of ions. Ionic compounds are called Salts or Crystals
  • Always formed between METALS (lose)  and NON-METALS (gain)
  • hard solid at 22 oC. 
  • High mp. Temperatures. 
  • Nonconductors of electricity in solid phase.
  •  Good conductors in liquid phase or dissolved in water (aq)
• COVALENT (NONMETAL ^2)

  • Forms Molecules. 
  • SHARE electron
  • Low m.p. temp and b.p. temperature. 
  • soft solids as compared to ionic compounds 
  • Nonconductors of electricity in any phase
• NON POLAR

  • When two atoms of the same element share electrons, they are shared equally
  • equally charged
• POLAR (dipoles)

  • when the atoms of different elements share electrons, they tend to be pulled closer
  • opposite charge
• Metallic Bonding

  • Occurs between like atoms of a metal in the free state 
  • Valence e- are mobile (move freely among all metal atoms) 
  • Positive ions in a sea of electrons ▪Metallic characteristics ▪
  • High m.p. temps, ductile, malleable, shiny. Hard substances. Good conductors of heat and electricity 
• Lewis Dot Structure


  • Symbol represents the KERNEL of the atom (nucleus and inner e-)
  • Dots represent valence e-
• Steps for drawing Lewis Dot Structure

  1. Count total valence e- involved

  2. Connect the central atom (usually the first in the formula) to the others with single bonds

  3. Complete valence shells of outer atoms

  4. Add any extra e- to the central atom