Chapter 23

1. Chapter 23 Chemical Bonding Presented by April Senger

2. Valence Shells & Lewis Dot Structures • The valence electrons are the outermost electrons that are available for bonding • For the purpose of understanding how atoms bond, we can ignore the inner electrons • To draw the model more simply we use a dot structure as shown on page 389 for groups 1, 2, 13, 14, 15, 16, 17 & 18 • Electrons shown side by side are bonding pairs and single dots are unpaired electrons available for bonding Presented by April Senger

3. Concept Check • What is a valance electron? • Any electron in the outermost shell that is available for chemical bonding with another atom Presented by April Senger

4. Ions • Ionization is the process of adding electrons to or removing electrons from an atom or groups of atoms • Atoms like to have 8 electrons. If it doesn’t it will try to add or lose electrons to get there • When an atom gains or loses electrons it becomes an ion…these have electrical charges • If you gain an electrons you are negative called anion • If you lose electrons you are positive called cations • If unstable Li loses one electron, it will only have 2 like He and that is stable Presented by April Senger

5. Ions & More Ions • If unstable Ca loses two electrons, it will only have 18 and can act like stable Ar • If unstable Br takes an electron it becomes negative and can act like stable Kr • Anions are written like F- • Cations are written like Ca 2+ • Metals that form + and nonmetals – bond together • The Alkali and Halogens commonly bond Presented by April Senger

6. Concept Check • What type of ion does t he magnesium atom, Mg, tend to form? • Mg forms 2+ Presented by April Senger

7. Ionic Bonds • Ionic Bonds are a bond formed by the attraction between oppositely charged ions. • When a positive and negative ion are attracted together • Nonmetals and Metals!!! Presented by April Senger

8. Ionics are Ornery • Atoms do not always share electrons equally and transfer them to one atom more than another • Chlorine is the greediest-it steals the Sodium electrons in salt to its side a little more. • The Na becomes a little more positive because it loses electrons and Cl becomes a little more negative with extra electrons Presented by April Senger

9. Facts Continued… • The positives and negatives are attracted to form a bond to form a NETWORK called salt • Salt is a series of NaCl molecules in a ratio of 1:1 • We call the NaCl a formula unit • CaF2 is a ratio of 1:2 to even the charges • Ionic bonds are held in place and can’t conduct electricity unless dissolved in water where they can move freely • The NETWORKS make them strong and have high melting points Presented by April Senger

10. Concept Check • Is the transfer of an electron from a sodium atom to a chlorine atom a physical or chemical change? • Electron transfer results in the formation of a new substance and therefore is a chemical change • What is the chemical formula for the ionic compound magnesium oxide? • MgO Presented by April Senger

11. Metallic Bonds • Metallic bonds are a bond formed by the attraction between positively charged metal ions and the electrons around them. • Bonds between positive metal ions and electrons in the environment Presented by April Senger

12. Metals have unique properties such as bending, stretching, flexibility, etc A metal is attracted to another’s electrons and draws them very close together When this close together, the orbits overlap and electrons can flow from atom to atom Electricity is flowing electrons so metals naturally allow electricity to flow Metals are flexible because the atoms can slide past each other Metal Facts Presented by April Senger

13. Covalent Bonds • Covalent bonds are a kind of bond formed when atoms share one or more pairs of electrons. • A bond where the electrons are shared • Nonmetal and Nonmetal bond • Don’t Forget!!! Hydrogen is a nonmetal on the metal side • An atom can have single, double or triple covalent bonds Presented by April Senger

14. Caring Is Sharing • Examples of covalent bonds are sugar, water, and networks like silicon dioxide • Covalent bonds only occur between nonmetals • Covalent bonds can be solids, liquids or gases • They do not conduct electricity or have a charge • They dissolve easily in water, can move freely if melted, and have low melting points • The molecules themselves are strong but the attractions between the molecules is what gives them the low melting point • All gases are in 2-packs & are covalently bonded Presented by April Senger

15. Nonpolar & Polar Covalent Bonds • Nonpolar covalent bonds are when the electrons are shared evenly like in Cl-Cl gas • Polar covalent bonds are when the electrons are not shared evenly. • Elements to the right and closer to the top “want” the electrons more = electronegativity • They pull the shared electrons closer giving a slightly negative charge. The other element becomes slightly positive. Ex: Water Presented by April Senger

16. Concept Check • How many electrons make up one covalent bond? • Covalent bonds share 2 electrons • Why do polar covalent bonds have a slightly higher boiling point than nonpolar bonds? • They have more attraction which means they require more energy to break them apart Presented by April Senger

17. Polyatomic Ions • Polyatomic ions are an ion made of two or more atoms that are covalently bonded and that act like a single ion. • It is a group of atoms with covalent bonds that act as an ion (have a charge) • An example is Sodium hydrogen carbonate or NaHCO3 • Positive ions are attracted to negative ions • Ammonium likes to be with Nitrate or Sulfate Presented by April Senger

18. Hurray For Facts! • Parentheses are used to remind you that they act as a single ion • Without parentheses the numbers of atoms can become incorrect such as (NH4)2SO4 • -ate is used for ions with the most oxygen atoms • -ite is used for ions that have the least oxygens Presented by April Senger