Chapter 5

1. Chapter 5 Atoms and Bonding

2. Atoms, Bonding, and the Periodic Table • What does the periodic table tell you about the atoms of elements? • What are the trends in the Periodic Table?

3. Valence Electrons and Bonding • Valence electrons- electrons on the outermost energy level. • The number of valence electrons in an atom of an element determines how the atom can bond with other atoms. • Electron Dot Diagram – shows an atoms valence electrons as dots around the elements symbol

4. How the Periodic Table Works • The periodic table reveals how many valence electrons an element has: Group 1 = 1 valence, Group 2 = 2 valence, Group 13 = 3 etc • Groups 3-12 vary

5. The Periodic Table • Across a period, the number of valence electrons increases by 1 until the maximum number of valence electrons is reached. • What is the maximum number of valence electrons? 8 Group 18 elements are the inert gases. All but Helium (2), have 8 valence electrons.

6. Chemical bond • Atoms like to be stable. • An atom is most stable when it has 8 valence electrons. • Some atoms will transfer electrons and others will share electrons between them in order to become more stable. • Chemical bond- force of attraction that holds two atoms together as a result of the rearrangement of electrons between them. • Two types of chemical bonds: ionic and covalent.

7. Ionic Bonds • How do ions form bonds? • What are the properties of ionic compounds? Atoms with 5, 6, or 7 electrons usually become more stable when they gain electrons, while atoms with 1, 2, or 3 electrons become more stable when they lose electrons. Recall that protons are “+” and electrons are “-”

8. Ionic Bonding

9. Ions • Ions: atoms or groups of atoms that have lost or gained electrons. • When atoms lose an electron it loses a negative charge and becomes a positive ion (cation). (lose weight +) • When an atom gains an electron it gains a negative charge and becomes a negative ion (anion). (gain weight -)

10. Ionic Bonds • Ionic bond: the attraction between positive and negative ions.

11. Bond Formation • The positive sodium ion and the negative chloride ion are strongly attracted to each other. • The compound sodium chloride, or table salt, is formed. • A compound is a pure substance containing two or more elements that are chemically bonded.

12. Chemical Formulas • Chemical formula – combination of symbols that shows the ratio of elements in a compound • Ex. MgCl2 = 1 Mg: 2 Cl a ratio of 1:2 • The 2 is a subscript. If no number is present we assume a 1 • When ionic compounds form, the ions come together in a way that balances out the charges on the ions. Chemical formulas reflect this balance. What are the ratios of the following chemical formulas? • NaCl = • P2O5 =

13. Writing Chemical Formulas Identify the charges of the ions from the periodic table. Then write the formula. Sodium bromide – Na = +1, Br = -1 NaBr Try these: • Lithium oxide • Aluminum fluoride • Potassium nitrate • Li2O • AlF3 • KNO3

14. Chemical Names For ionic compounds, the name of the positive ion comes first, followed by the names of the negative ion. -ide endings for single negative ions (chloride) -ate or –ite endings for negative polyatomic ions (carbonate) MgO = Magnesium oxide or NH4NO3 = Ammonium nitrate Try these: 1. K2S 2. NH4Cl 1. Potassium sulfide 2. Ammonium chloride

15. Covalent Bonds • What holds covalently bonded atoms together? The force that holds atoms together in a covalent bond is the attraction of each atom’s nucleus for the shared pair of electrons. Molecule – a neutral group of atoms joined by covalent bonds

16. Covalent bonding

17. Covalent Bond • Covalent bond: the attraction of each atom’s nucleus for the shared pair of electrons.

18. Examples • The oxygen atom in water and the nitrogen atom in ammonia each have eight valence electrons as a result of forming covalent bonds with hydrogen atoms.

19. Multiple bonds • Double and triple bonds can form when atoms share more than one pair of electrons.

20. Properties of Molecular Compounds • Low melting points and boiling points • Poor electrical conductivity

21. Polar Covalent Bonding

22. Polar covalent bonding • Some atoms pull more strongly on the shared electrons than other atoms. • Example: Fluorine forms a nonpolar bond with another fluorine atom because each atom pulls on the electrons equally. However, in hydrogen fluoride, fluorine pulls on electrons more strongly than hydrogen does, so it forms a polar bond (pulls on electrons unequally).

23. Polar Covalent Bonding • A carbon dioxide molecule is a nonpolar molecule because of its straight-line shape. • In contrast, a water molecule is a polar molecule because of its bent shape.

24. Bonding in Metals Metals – consist of just one element Alloys – mixture made of two or more elements, at least one of which is a metal Alloys are generally stronger and less reactive than the pure metals from which they are made. Physical properties – retain many of physical properties of metals Chemical properties – not as reactive as pure metals

25. Physical properties – alloys retain many of physical properties of metals Pure gold is shiny and easily bent, but when pure gold is mixed a harder element (gold alloy), it still retains its shininess, but also is much harder and more difficult to bend. drill rods Chemical properties – not as reactive as pure metals Iron often rusts, but when mixed with other elements to make steel, it becomes stronger and less likely to rust.

26. Metallic Bonding Metal atoms combine in regular patterns in which the valence electrons are free to move from atom to atom. Metallic bond – an attraction between a positive metal ion and the many electrons surrounding it.

27. Metallic Properties • Malleability – pound into sheets • Ductility – drawn into wires • Luster – shiny and reflective • Electrical conductivity – electrons can move freely among atoms • Thermal conductivity - heat