1. Chemical Bonds

2. Atom – the smallest unit of matter “indivisible”
Helium
atom

3. electron shells Atomic number = number of Electrons
Electrons vary in the amount of energy they possess, and they occur at certain energy levels or electron shells.
Electron shells determine how an atom behaves when it encounters other atoms

4. Electrons are placed in shells according to rules:
The 1st shell can hold up to two electrons, and each shell thereafter can hold up to 8 electrons.

5. Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons
C would like to
N would like to
O would like to
Gain 4 electrons
Gain 3 electrons
Gain 2 electrons

6. Why are electrons important?
Elements have different electron configurations
different electron configurations mean different levels of bonding

8. Electron Dot Structures
Symbols of atoms with dots to represent the valence-shell electrons
H He:
          
Li Be  B   C   N   O  : F  :Ne :
       
          
Na Mg  Al  Si  P S :Cl  :Ar :
       

9. Chemical bonds: an attempt to fill electron shells
Ionic bonds
Covalent bonds
Metallic bonds

10. Learning Check A. X would be the electron dot formula for
A. X would be the electron dot formula for
1) Na 2) K 3) Al
 
B  X  would be the electron dot formula
1) B 2) N 3) P

11. IONIC BOND bond formed between two ions by the transfer of electrons

12. Formation of Ions from Metals
Ionic compounds result when metals react with nonmetals
Metals lose electrons to match the number of valence electrons of their nearest noble gas
Positive ions form when the number of electrons are less than the number of protons
Group 1 metals  ion 1+
Group 2 metals  ion 2+
Group 13 metals  ion 3+

13. Formation of Sodium Ion
Sodium atom Sodium ion
Na  – e  Na +
( = Ne)
11 p p+
11 e e-

14. Formation of Magnesium Ion
Magnesium atom Magnesium ion 
Mg  – 2e  Mg2+
(=Ne)
12 p p+
12 e e-

15. Some Typical Ions with Positive Charges (Cations)
Group 1 Group 2 Group 13
H+ Mg2+ Al3+
Li+ Ca2+
Na+ Sr2+
K+ Ba2+

16. Learning Check A. Number of valence electrons in aluminum
1) 1 e ) 2 e- 3) 3 e-
B. Change in electrons for octet
1) lose 3e ) gain 3 e ) gain 5 e-
C. Ionic charge of aluminum
1) ) ) 3+

17. Solution A. Number of valence electrons in aluminum 3) 3 e-
B. Change in electrons for octet
1) lose 3e-
C. Ionic charge of aluminum
3) 3+

18. Learning Check Give the ionic charge for each of the following:
A. 12 p+ and 10 e-
1) 0 2) 2+ 3) 2-
B. 50p+ and 46 e-
1) 2+ 2) 4+ 3) 4-
C. 15 p+ and 18e-
2) ) 3- 3) 5-

19. Ions from Nonmetal Ions
In ionic compounds, nonmetals in 15, 16, and 17 gain electrons from metals
Nonmetal add electrons to achieve the octet arrangement
Nonmetal ionic charge:
3-, 2-, or 1-

20. Fluoride Ion     1 - : F  + e : F :     2-7 2-8 (= Ne)
unpaired electron octet
   
: F  e : F :
   
(= Ne)
9 p p+
9 e e-
ionic charge

21. Ionic Bond
Between atoms of metals and nonmetals with very different electronegativity
Most ionic compounds are called salts
Bond formed by transfer of electrons
Produce charged ions all states. Conductors and have high melting point.
Examples; NaCl, CaCl2, K2O

22. Ionic Bond
Active metals readily loses its valence electrons, usually to a nonmetal atom
Oppositely changed ions are formed by this process of transferring electrons
Cation: positively charged ion
Anion: negatively charged ion

23. Ionic Bond Types of ions:
Monoatomic: ion formed from one atom
Polyatomic: ion made up of 2 or more atoms bonded together that acts like a single unit
Ionic Compound: composed of positive and negative ions so that the positive and negative charges equal in number

25. Ionic Bonds: One Big Greedy Thief Dog!

26. 1). Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions.

28. Properties of Ionic Compounds
Strong Bond
Form Crystal Lattice
Melting Point, Boiling Point, Hardness  due to strong bonds tend to be high
Good conductors of electricity when dissolved in water
Form electrolytes
Formation of ionic bonds is always exothermic

29. Lattice energy
Energy required to separate 1 mol of ions in ionic cmpd. (Strength holding ions in place)
The more negative the lattice energy the stronger the force of attraction
Related to the size of the ions bonded (LiCl has higher lattice energy than LiBr since atoms are closer together)

30. Naming and Formulas
Formula Unit  simplest ratio of ions represented in an ionic compound.
Charges are also known as oxidation state

31. Naming Name cation followed by anion
Monatomic cations use element name
Monatomic anions use their root name followed by –ide
Oxygen is oxide etc.
Transition metals often have more than 1 oxidation number and therefore use Roman Numerals
When using Polyatomic ions simply name ion

32. COVALENT BOND bond formed by the sharing of electrons

33. Covalent Bond
Between nonmetallic elements of similar electronegativity.
Formed by sharing electron pairs
Stable non-ionizing MOLECULES, they are not conductors at any state
Examples; O2, CO2, C2H6, H2O, SiC

34. Covalent Bond May share 1 or more pairs of electrons
Single bond- one pair of shared electrons between 2 atoms
Double bond- two pairs of shared electrons between 2 atoms
Triple bond- three pairs of shared electrons between 2 atoms

35. Covalent Bond
Atoms that are bonded covalently form stable particles called molecules.
Ex: H20, CO2, and C6H12O6
7 Diatomic molecules to know – H2, N2, O2, F2, Cl2, Br2, and I2
Molecular compound: 2 or more atoms are covalently bonded, lowering their potential energy, than its individual atoms

36. Formation of Covalent Bond
These bonds form due to ATTRACTIVE and REPULSIVE forces

37. Covalent Bonds

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

39. when electrons are shared equally
NONPOLAR COVALENT BONDS
when electrons are shared equally
H2 or Cl2

40. 2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons.
Oxygen Atom
Oxygen Atom
Oxygen Molecule (O2)

41. when electrons are shared but shared unequally
POLAR COVALENT BONDS
when electrons are shared but shared unequally
H2O

42. Polar Covalent Bonds: Unevenly matched, but willing to share.

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

44. 9.2 Naming Molecules First element is named using entire element name
Second element uses ending “ide”
Add appropriate prefix to indicate the number of atoms.
Mono = 1
Di = 2
Tri = 3
Tetra = 4
Penta = 5
Hexa = 6
Hepta = 7
Octa = 8

45. Naming Molecules Extras… Do not use mono with the first element
Ex. CO2 = carbon dioxide
Do not leave “ao” or “oo” combination in name
Ex. N2O4 = dinitrogen tetroxide
NO = nitrogen monoxide

46. Naming Molecules Let’s Practice… CCl4 = carbon tetrachloride As2O3 =
diarsenic trioxide
SO2 =
sulfur dioxide
NF3 =
nitrogen trifluoride

47. Naming Molecules
Many molecules were named prior to the development of the modern naming system.
Formula
Common Name
Molecular Compound Name
H2O
water
dihydrogen monoxide
NH3
ammonia
nitrogen trihydride
N2H4
hydrazine
dinitrogen tetrahydride
N2O
nitrous oxide (laughing gas)
dinitrogen monoxide NO
nitric oxide
nitrogen monoxide
table from page 249

48. METALLIC BOND bond found in metals; holds metal atoms together very strongly

49. Metallic Bond Formed between atoms of metallic elements
Electron cloud around atoms
Good conductors at all states, lustrous, very high melting points
Examples; Na, Fe, Al, Au, Co

50. Metallic Bonds: Mellow dogs with plenty of bones to go around.

51. Metallic Bonds, A Sea of Electrons

52. Metals Form Alloys
Metals do not chemically combine with metals. They form alloys which is a solution of a metal in a metal. Examples are steel, brass, bronze and pewter.

53. Formula Weights Formula weight is the sum of the atomic masses.
Example- CO2
Mass, C + O + O
43.999

54. Practice
Compute the mass of the following compounds round to nearest tenth & state type of bond:
NaCl;
= 58; Ionic Bond
C2H6;
= 30; Covalent Bond
Na(CO3)2;
23 + 2( x16) = 123; Ionic & Covalent