Periodic table

1. Periodic table Lec. 3

2. Atoms vs. Ions • Atomsare neutral; they contain the same number of protons as Electrons. • Ionis an electrically charged particle produced by either removing electrons from a neutral atom to give a positive ion or adding electrons to a neutral atom to give a negative ion. In ions: • The number of protons does not change . • Positively charged ions by removing one or more electrons, ex:Na+ ion , charge of +1( known as cation). • Negatively charged ions By adding one more electron ,ex: Cl- ion ,charge of -1(known as anion).

4. Compounds Elements combine to form chemical compounds that are often divided into two categories: 1-Ionic compounds Metals often react with nonmetals to form ionic compounds. These compounds are composed of positive and negative ions formed by Adding or subtracting electrons from neutral atoms and molecules. 2Na+Cl2 → 2NaCl 2-Molecular compounds Nonmetals combine with each other to form molecular compounds (covalent compounds) which exist as neutral molecules 2H2+O 2 → 2 H2O

5. Types of chemical bonds A Molecule is a chemical species that is formed when two or more atoms join together chemically, with covalent or ionic bonds . *Ionic bond: A chemical link between two atoms caused by the electrostatic force between oppositely-charged ions in an ionic compound. An ionic bond is formed by the transfer of electrons form one atoms to the other. One atom then takes on a positive charge (the cation) and the other a negative charge (the anion). *Covalent bond A chemical link between two atoms in which electrons are shared between them.

6. Introduction to the periodic table • The periodic table lists all the known elements in numerical order, starting with the lightest (hydrogen) and • proceeding to the heaviest (uranium, among naturally occurring elements). • The list is broken into seven rows. Each row is placed below the previous row in a way that places elements with similar chemical properties in the same column of the table. • Moving across a row of the periodic table, the elements generally increase in mass and change dramatically in their chemical properties. • Moving down a column, mass also increases, but the elements • have similar chemical properties.

8. The elements can be divided into categories: metals, nonmetals, and metalloids(semimetals). • Examples of each appear in the Figure except for hydrogen, all the elements in the left and central regions of the periodic table are metals. • Metals display several characteristic properties. For example, they are good conductors of heat and electricity and usually appear shiny. • The first column of the periodic table, Group 1, contains elements that are soft, shiny solids. These alkali metals include lithium, sodium, potassium, rubidium, and cesium. At the other end of the table, fluorine, chlorine, bromine, iodine, and astatine appear in the next-to-last column. These are the halogens. • The elements in the second column of the table (Group 2) are the alkaline earth metals. These resemble the alkali metals in their appearance, but they have different chemical properties.

9. The last column of the periodic table contains the noble gases, or Group 18 elements, all of which occur in nature as gases. With a few exceptions, these elements do not undergo chemical reactions. • The elements in Groups 3 through 12 are known as transition metals. • A list of the chemical elements arranged into Groups (the columns going down) and Periods (the rows going across). • The group number gives the number of valence (outer shell) electrons around the atom. • The Period number is the same as that of the outer energy shell. • Elements increase in atomic number across each period by one. • EXAMPLE: Calcium is in the 2nd group therefore it has 2 electrons in the outer shell and a valiancy of +2. It is in the fourth period therefore it’s outer Electrons are in the fourth energy shell. Ca: 2,8,8,2. • Element Configuration Shorthand: • Hydrogen H )1e- Lithium Li )2e- )1e- • Fluorine F )2e- )7e- Sodium Na )2e- )8e- )1e-

10. Electron Configurations for Selected Elements in period two Group IA VIA VIIA VIIIA Lithium Oxygen Fluorine Neon Chemical “Families” Going from left to right on the periodic table, you will find metals, then metalloids, and finally nonmetals. • Group IA are called alkali metals because the react with water to from an • alkaline solution. 2. Group IIA are called the alkali earth metals because they are reactive, but not as reactive as Group IA. They are also soft metals like Earth. 3. Group VIIA are the halogens These need only one electron to fill their outer shell They are very reactive. 4. Group VIIIA are the noble gases as they have completely filled outer Shells.

11. . Trends in atomic radius in the Periodic Table • Period Trends • Atomic radius decrease from left to right across a period • The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus. • The strong positive charge causes the electrons to be pulled inward; creating a decrease in radius • Group Trends • Atomic radius of the main group element increase down a group. • As electrons occupy sublevels in successively higher main energy levels located farther from the nucleus, the sizes of the atoms increase. Give reason The atoms get bigger as you go down groups.(that an extra layer of electrons are added.)

12. Ionization Energy • Energy required to overcome the attraction of the nuclear charge to remove • an electron(highest energy) from an atom. • Decrease down the groupdue to Size of atoms increases, outer electron farther from nucleus, So the electron will be easier to remove due to nucleus has weaker attraction. • Increase across the period due to Nuclear charge increasing, greater attraction of nucleus for electrons. • The ionization energy may be an indicator of the reactivity of an element. Elements with a low ionization energy tend to be reducing agents and form cations.

13. Electron Affinity • The electron affinityof an element is the energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion . • A fluorine atom in the gas phase, for example, gives off energy when it gains an electron to form a fluoride ion. • F(g) + e- F-(g) ∆Ho = -328.0 kJ/mol • Electron affinities generally become smaller as we go down a column of the periodic table for two reasons: • First, the electron being added to the atom is placed in larger orbitals, where it spends less time near the nucleus of the atom. • Second, the number of electrons on an atom increases as we go down a column, so the force of repulsion between the electron being added and the electrons already present on a neutral atom becomes larger.

15. Electro negativity • Electro negativity is a measure of the tendency of an atom to attract a bonding pair of electrons. • Fluorine : the most electronegative element. • Argon. It doesn't have electro negativity, because it doesn't form bonds. • Electro negativity across a period increases because the number of charges on the nucleus increases. That attracts the bonding pair of electrons more strongly. • Note: [Atomic radius decreases → Ionization energy increases → Electronegativity increases]

16. Electro negativity down a group decreases because the bonding pair • of electrons is increasingly distant from the attraction of the nucleus. • The attraction that a bonding pair of electrons feels for a particular nucleus depends on: 1. The number of protons in the nucleus; 2. Distance from the nucleus 3. The amount of screening by inner electrons.