2. To bond or not to bond….

3.  Chemical Bond: is a strong attractive force between atoms or ions in a compound.  Drawn as a stick or dots of electrons

4. Electrons bond  Valence electrons: electrons in the outer shell or orbital  Octet rule: most atoms prefer to have 8 electrons in their outer shells :

5. Chemical formula  Chemical formula: for a compound reveals the elements and their ratios in that compound  Atoms are chemically stable when their outer shells are full  NaClCO  CaCl 2 ZnO CO 2 Fe 2 O 3 FeO

6. Ions  Ionic bond: a force of attraction between two oppositely charged ions  Ions: are charged particles that form when atoms gain or lose electrons These atoms are said to have charges  Calculate charges by subtracting the number electrons of from the number of protons

7. Anion  Atoms that gain electrons form negative ions  Anion: Negatively charged ion

8. Cation  Atoms that lose electrons form positive ions  Cation: positively charged ion  Cats are positively cwazy

9. Ions  Charged ions that bond form a neutral compound as in salt Na + Cl -  Energy is required and can either be given off or absorbed when electrons move In ionic compounds the ions will often disassociate in water Ions are important in your nerve and muscle cells – often called electrolytes

10. Covalent bonds  Covalent bonds: Atoms share electrons Molecules are formed from at least two covalently bonded atoms Water has a covalent bond :.

11. Covalent Bond  Between nonmetallic elements of similar electronegativity.  Stable non-ionizing particles, they are not conductors at any state  Examples; O 2, CO 2, C 2 H 6, H 2 O, SiC

13. Bonds in all the polyatomic ions and diatomics are all covalent bonds

14. Electonegativity  Electronegativity: the pull one atom has for another atom's electrons during the bonding process. Electronegativity determines what type of bond will be formed

15.  when electrons are shared equally NONPOLAR COVALENT BONDS

16. Diatomic again  Diatomic molecules are two atoms bonded together O, H, N, and all halogens often called diatomic elements

17. 2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons. Oxygen Atom Oxygen Molecule (O 2 ) Oxygen Molecule (O 2 )

18. Polar express  Polar Molecule has a positive end and a negative end  This is caused by uneven sharing of electrons Nonpolar molecules do not have positive or negative ends- they share evenly

19. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

20. Polar vs. Nonpolar

21. Oxidation numbers  Oxidation number: shows the combining ability of an element in a compound  Na +1Mg +2  Cl –1B +3  Some elements have more then one oxidation state  Fe, Cu, Cr, Pb, Sn  We must indicate the oxidation state when writing these compounds or elements  Fe (III), Cu (II), Sn (II),

22. .

23. Bond type is determined by electonegativity

24. Naming Covalent Compounds Covalent compounds are named by adding prefixes to the element names. The compounds named in this way are binary covalent compounds. ‘Binary’ means that only two atom are present. ‘Covalent’ (in this context) means both elements are nonmetals. A prefix is added to the name of the first element in the formula if more than one atom of it is present. (The less electronegative element is typically written first.) A prefix is always added to the name of the second element in the formula. The second element will use the form of its name ending in ‘ide’.

25. Naming Covalent Compounds Prefixes SubscriptPrefix 1mono- 2di- 3tri- 4tetra- 5penta- SubscriptPrefix 6hexa- 7hepta- 8octa- 9nona- 10deca- Note: When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.

26. Naming Binary Covalent Compounds: Examples N 2 S 4 dinitrogen tetrasulfide NI 3 nitrogen triiodide XeF 6 xenon hexafluoride CCl 4 carbon tetrachloride P 2 O 5 diphosphorus pentoxide SO 3 sulfur trioxide 1mono 2di 3tri 4tetra 5penta 6hexa 7heptaa 8octa 9nona 10deca * Second element in ‘ide’ from * Drop –a & -o before ‘oxide’

27. Writing formulas  1.Symbol with a positive oxidation number is first (metals always have positive oxidation numbers)  2.Symbol with negative oxidation number  3.Write subscripts so oxidation number of the compound totals zero  Use parantheses around polyatomic ions if using multiples of these

28. Polyatomic  A group of covalently bonded atoms which have charge.  Usually negatively charged  Treat polyatomic ions as if they are one thing, don’t mess with their subscripts.

29. Naming  Write the name of the positive ion  Write the root of the negative ion-( ox, flor, chlor, iod,etc.) and add ide.  For polyatomics just write the name of the ion  For multivalent metals tell valence state using roman numerals in paranthesis. Ex. Iron (II) chloride

30. Practice FormulasNaming  Calcium hydroxide  Ca(OH) 2  Aluminum bromide  AlBr 3  Iron (III) sulfate  Fe 2 (SO 4 ) 3  Silver nitrite  AgNO 3  CuO  Copper (II) oxide  LiNO 3  Lithium nitrate  (NH 3 ) 2 CrO 4  Ammonium chromate  Sr(C 2 H 3 O 2 ) 2  Strontium acetate