 Russian chemist Dmitri Mendeleev placed the known elements in order of increasing atomic mass.  When he did this he noticed that the elements’ properties repeated in a regular pattern, or a periodic pattern.  Mendeleev placed the known elements in a table, where he arranged elements into columns with similar properties.

 Mendeleev predicted the properties of several elements that were unknown at the time.  Many of his predictions were correct and were well accepted by the scientific community.

 The only changes to Mendeleev’s periodic table were the addition of newly found elements and that the table was organized by atomic number rather than atomic mass.

 The horizontal rows on the periodic table are called periods. This is because the properties begin to repeat in each new row.  The columns on the periodic table are called groups (or family) and the elements in the groups have similar properties.  Groups are designated with a number and the letter A or B.

 The main group elements can be found in Groups 1A through 8A.  The main group elements (Group A elements) are also called representative elements.

 The group B elements are called the transition elements.

 Elements can be divided into three main classes.  Those classes are: › Metals › Metalloids › Nonmetals

 Metals have the following properties. They are generally: › Shiny solids › Good conductors of heat and electricity  Group 1A has the name alkali metals  Group 2A has the name alkaline earth metals.

 Elements to the right of the heavy stair- step line on the periodic table are called nonmetals.  Nonmetals generally have the properties that they are generally gases or brittle solids at room temperature.  Group 7A are called halogens.  Group 8A are called the noble gases.

 Many of the elements that border the stair step are metalloids.  Metalloids share properties in between those of metals and nonmetals.

 Column A 1. Strontium 2. Chromium 3. Iodine 4. Nitrogen 5. Argon 6. Rubidium 7. Silicon  Column B a. Halogen b. Noble gas c. Alkaline earth metal d. Metalloid e. Alkali metal f. Representative element g. Transition Element

 Scientists now understand that the repetition of properties of elements occurs because the electron configurations of atoms repeat.  The arrangement of elements in the periodic table reflects the electron structures of atoms.

 For representative elements, the Group Number in front of the A tells the number of valence electrons in the atoms in the column.  Also the period number (or row number) of a representative element tells the energy level of the valence electrons.

 The periodic table is divided into blocks that correspond to the energy sublevel being filled as you move across a period.  Groups 1A and 2A are the s block.  Groups 3A- 8a are the p block.  The B elements represent the d block; however remember to go down one for the first quantum number.  The rows that are removed from the table and placed at the bottom represent the f block. Go down 2 numbers for the row number.

 Without using the periodic table, determine the group, period, and block in which an element with each of the following electron configurations is found. 1. [He] 2s 2 2p 5 2. [Ar] 4s 2 3. [Kr]5s 2 4d 10 5p 3 4. [Ar]4s 2 3d 3

Periodic Trends

 The electron structure of an atom determines many of its chemical and physical properties.  There are several trends that can be observed using the periodic table.

 The atomic radius is a measure of the size of an atom. The larger the radius, the larger the atom.  As you move across a period the atom decreases in size.  This is due to the increasing positive charge of the nucleus while electrons are being added, but the orbitals are close in energy.  The increased nuclear charge pulls the outermost electrons closer to the nucleus, making the atom smaller.

 As you move down a group, atomic radii increases.  This is due to the addition of a larger energy level each time.  Electrons in higher levels are located farther from the nucleus than those in the lower energy levels.

 From each of the following pairs, predict which atom is larger. 1. Mg, Sr 2. Sr, Sn 3. Ge, Sn 4. Ge, Br 5. Cr, W

 The driving force that makes reactions happen is ion formation.  Ions form when atoms gain or lose electrons.  When electrons are gained or lost the resultant ion has a positive or negative charge.

 When atoms lose electrons and form positively charged ions, they become smaller, and are called cations.  They become smaller because the loss of the electrons means that the number of protons is greater than the number of electrons.  Therefore, the electrons will be pulled more tightly to the nucleus, and the outer electrons will feel the pull of the nucleus more strongly than before.

 When atoms gain electrons they form negatively charged ions, called anions.  The added electron makes the ionic radius increase, because of repulsion and the weaker pull of the nucleus on each electron.

 As you move from left to right across the periodic table, your positively forming ions, Groups 1A-4A get smaller, but Groups 5A- 8A get larger.  As you move from top to bottom the ions get larger in each column because of the addition of energy levels.

 To form a positive ion, an electron must be removed from a neutral atom.  Removing the electron requires energy. That energy must overcome the attraction between the positive charge in the nucleus and the negative charge of the electron.  This energy, known as ionization energy is defined as the energy required to remove an electron from an atom in the gaseous state.

 The first ionization energy is the amount of energy required to remove the first electron from the outer shell of the atom.  Remember that as you move across a period, you increase atomic number and therefore add more positive charge to the atom.  That addition makes it more difficult to remove the electron and therefore more energy is required.

 For the first ionization energy, as you move across a period, from left to right, the ionization energy increases.  As you move down a group, the ionization energy generally decreases. Why?

 It is also possible to remove electrons after removing the first, or furthermost electron from the electron cloud.  However, once you remove the first electron, much more energy is required to remove the second, and third, and fourth, etc.

 Remember, when Newland tried to design his periodic table he came up with a law of octaves, which wasn’t accepted.  We see using ionization energies that his predictions were every correct.  When atoms/ions have eight electrons in their outer shell, they are much more stable.

 The octet rule has come to be one of the most important principles in chemistry.  It says that atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons.  Use this to predict what kinds of ions will form: elements on the left will form positive ions and elements on the right negative ions.

 The electronegativity of an element tells about its ability to attract electrons in a chemical bond.  The more electronegative an element, the more electron loving it is.  Noble gases form few compounds and so virtually lack electronegativity.

 As you move from top to bottom down a group, electronegativity generally decreases.  As you move from left to right, electronegativity generally increases.  So what is the most electronegative element?