1. E03 Chemical Bonding and Naming
Jenny Wu

2. Ted-ed: How atoms bond

3. Atoms bind together to form bonds
Atoms want a full outer electron shell to achieve stability, like the noble gases.
For compounds not involving transition metals or rare earth metals, that means 8 outer electrons (to fill the S2 and P6 orbitals); this is known as the octet rule.
The outermost “ring” of electrons are called valence electrons

4. Atoms bind together to form bonds
3 types of bonds – to get a full valence shell, atoms form:
1) Metallic bonds (formed between metals) – atoms pool electrons together to form a “sea” of electrons, co-shared by every atom
2) Ionic bonds (formed between a metal and nonmetal) – atoms give/take electrons from each other in order to attain a full valence shell
3) Covalent bonds (formed between nonmetals) – atoms “share electrons in order to attain a full valence shell

5. 1) Metallic bonding
Metals bond through a “sea of electrons” where electrons flow freely between the all the individual metal atoms.
The nucleus of each atoms stays put, and the electrons are collectively shared.
The freedom electrons have to freely move about is why metals conduct electricity.

6. 1) Metallic bonding Alloys: A mixture of 2 or more metals
Alloys are made by melting two or more pure metals together.
This is possible because metals form a “sea of electrons” where the individual atoms in a block of metal all share the same “pool” of electrons between them.

7. 1) Metallic bonding Examples of Alloys
Bronze: Tin and Copper
Brass: Copper and Zinc
Steel: Iron and Carbon
Alloys have long been made because they have qualities in malleability, strength or conductivity that is superior to their component pure metals.

8. Compounds and Molecules
Compound: a chemical where 2 or more elements are chemically bonded together.
Ex: H2O, NaCl
Molecule: a single unit of a chemical sharing covalent bonds; Cl2, O2, H2O

9. Compounds and Molecules
2)Ionic Compounds: are held together by ionic bonds, where electrons are gained/lost
Forms crystal lattices
Held together by electrical attraction (+ / – )
3) Covalent Compounds/ molecules: are held together by covalent bonds, where electrons are shared
Forms molecules
H2O Water
Held together by physical overlap of electron orbitals
Covalent bonds are stronger than ionic bonds

10. 2) Ionic bonds
Each atom in an ionic compound gains or loses electrons in order to gain a full valence shell
Ex: Table salt, NaCl

11. 2) Ionic bonds – more than 2 atoms can join together to form an ionic compound

12. How do you know how many electrons atoms want to gain or lose?

13. Charge numbers can be found by a pattern on the periodic table

14. Charge numbers and what they mean
+1 means that those atoms like to lose 1 electron to attain a filled valence shell (lose 1 electron, has 1 more proton than electron, therefore charge = +1)
-2 means those atoms like to gain 2 electrons
Etc.

15. Ions Cations: ions with a positive charge. They have lost electrons.
Ex: Na+ H+ Ca2+ Al3+
Anions: ions with a negative charge. They have gained electrons.
Ex: F- Cl- O2- N3-

16. 2) Ionic Compounds
Elements like to combine into electrically neutral compounds.
Examples:
Na+ + Cl- → NaCl
Mg2+ + Cl- → MgCl2
Li+ + O2- → Li2O
Al3+ + O2- → Al2O3
Number in superscript represent charge (ex. Ca2+); number in subscript represent number of atoms CaCl2

17. Format of a chemical equation:
Reactants → (yields) Products

18. Forming binary ionic compounds Practice Problems
Na + F →
H + Cl →
Mg + F →
Mg + O →
Al + S →
Ca + Cl →
Li + N →

19. Forming binary ionic compounds Practice Problems
Na + F → NaF
H + Cl → HCl
Mg + F → MgF2
Mg + O → MgO
Al + S → Al2S3
Ca + Cl → CaCl2
Li + N → Li3N

20. Polyatomic ions
Covalently bonded atoms that move together as one unit, and acts as one ion. It has one collective charge.
Examples:
ClO3- Chlorate
CO32- Carbonate
HSO4- Hydrogen Sulfate

21. Polyatomic ions
They combine the same way as single ions:
Examples:
H+ + ClO3- → HClO3
Hydrogen + Chlorate → Hydrogen Chlorate
Ca2+ + CO32- → CaCO3
Calcium Carbonate
… except you don't need to change the name of the anion to end in ---ide.

22. Polyatomic ions
They combine the same way as single ions:
Examples:
Mg2+ + ClO3- → Mg(ClO3)2
Magnesium + Chlorate → Magnesium Chlorate
Al3+ + CO32- → Al2(CO3)3
Aluminum Carbonate
… except you don't need to change the name of the anion to end in ---ide.

24. Naming ionic compounds
Name = cation name + anion name in ---ide
Example:
NaCl Sodium Chloride
MgO Magnesium Oxide
Li2O Lithium Oxide
CaF2 Calcium Floride
Just change the ending of the anion name to end in ---ide; the subscripts in the chemical formula do not change the name

25. Naming ionic compounds
Transition (d-block) metals get a roman numeral after their name to specify the charge.
Example:
Iron (II) Sulfide would be FeS
Iron (III) Sulfide would be Fe2S3

26. Naming ionic compounds
Some transition metals have a second, old, latin name. Expect to see this.
Example: Fe2O3
Regular Name: Iron (III) Oxide
Old Latin Name: Ferric Oxide
Example: FeO
Regular Name: Iron (II) Oxide
Old Latin Name: Ferrous Oxide

27. Naming ionic compounds
Some transition metals have a second, old, latin name. Expect to see this.
Example: CuS
Regular Name: Copper (II) Sulfide
Old Latin Name: Cupric Oxide
Example: Cu2S
Regular Name: Copper (I) Sulfide
Old Latin Name: Cuprous Sulfide

28. Naming ionic compounds Practice Problems
KCl
MgF2
K2O
Al2O3
BeCl2
FeO
CuCl
CuCl2

29. Practice Time Names  Chemical Formulas
Sodium Fluoride
Sodium Oxide
Sodium Nitride
Magnesium Oxide
CaI2
Li2S
Al2S3
K3N

30. List of common Polyatomic Ions to memorize
PO43- Phosphate
SO42- Sulfate
CO32- Carbonate
NO3- Nitrate
NH4+ Ammonium
OH- Hydroxide
HCO3- Hydrogen Carbonate (Bicarbonate)

31. Extended list of polyatomics/diatomics
Polyatomic Ions Chromate CrO4-2 Dichromate Cr2O7-2 Ferricyanide [Fe(CN)6]3− Ferrocyanide [Fe(CN)6]4− Permanganate MnO4− Manganeous Mn+x Chlorate ClO3- Sulfite SO3-2
Common Chem/
Diatomics
Ammonia NH3
Nitrogen N2
Oxygen O2
Fluorine F2
Chlorine Cl2
Bromine Br2 Iodine I2
Hydrogen H2

32. Practice Time Names  Chemical Formulas
Calcium Carbonate
Potassium Sulfate Jason
Ammonium Chloride Gyoung Na
Sodium Nitrate Kelly
CaSO4
LiOH Takato
Mg3(PO4)2 Margaux
(NH4)2S Claire

33. Ionic compounds are “salts”
Properties of salts:
High solubility
High melting point and boiling point (mp/bp),
does not conduct electricity unless molten or dissolved in solution,
forms crystal lattice

34. Solubility of ionic compounds/salts
Ions have High solubility

35. 2) Ionic compound nomenclature
Hydrate (adj): an ionic compound with water molecules incorporated into their crystal lattice
Anhydrous (adj): an ionic compound that has been heated to evaporate off all water

36. Vid -Chemical Bonds Song

38. Ionic Bonds
Atoms give and take electrons, acquire a charge, and is held together by electrical attraction.

39. The covalent bond
A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms.
– the sharing of electrons allows each atom to attain the equivalent of a full outer shell (8 e-); stability
A molecule is an electrically neutral group of two or more atoms held together by covalent bonds.

43. CLOUD

44. Examples of Molecules Water – H2O Oxygen – O2 Ozone – O3
Carbon Dioxide – CO2
Glucose (a part of sugar) – C6H12O6
*Note that nonmetal atoms bond together using covalent bonds

53. Polar vs. Non-polar covalent bonds
In the above picture, the dots represent the nucleus and the blue represent the electron clouds
In non-polar covalent bonds, electrons are shared equally; in polar ones, they are shared but unequally; in ionic bonds, electrons are given and taken.

54. Dipole moments
Some elements on the periodic table have more affinity than other elements for electron; this creates an unequal sharing of electrons and partial charges at either end of a molecule called “dipole moments.” These molecules are “polar.”

55. Electronegativity --the secret behind polarity
Linus Pauling won the Noble prize for chemistry by discovering “electronegativity” of atoms.
Electronegativity – is a chemical property that describes the tendency of an atom or a functional group to attract electrons (or electron density) towards itself.
An atom with a higher electronegativity have a stronger attraction for electrons.

56. Electronegativity on the Periodic Table
These are electronegativity numbers for the elements on the periodic table. 4 is the most electronegative, 0 is the least.

57. Electronegativity on the Periodic Table
Differences of >0.4 between atoms mean a polar covalent bond, differences of >1.7 mean an ionic bond.

58. Electronegativity pattern on the periodic table

59. Ted-ed Vid How polarity makes water behave strangely - Christina Kleinberg

60. Polar vs. Non polar Molecules
Some molecules have dipole moments, some do not.
Examples of non polar molecules: Methane (CH4), oil, wax
Examples of polar molecules: Water (H2O), Alcohol (Ethanol – C2H6O), Nail polish remover (Acetone – C3H6O)
Polar and non polar molecules do not mix.

61. Polarity - Like dissolves like
Polar substances will dissolve other polar substances.
Ex: Water dissolves salt, alcohol, etc.
Non-polar substances will only dissolve nonpolar substances.
Ex: Crayon dissolves in oil.

62. Water, H2O, is a polar molecule
Water has a partial negative charge (δ+) at the oxygen end and a partial positive charge at the hydrogen end (δ-)

63. Hydrogen bonding
Water molecules forms “hydrogen” bonds with each other. The positive ends of each water molecule (H), gets attracted to the negative ends (O) of other water molecules.

64. Hydrogen bonding
Warning – A hydrogen bond is NOT just a bond with hydrogen involved.

65. Hydrogen bonding
Anytime a hydrogen is bonded to any of the 3 highly electronegative atoms – Nitrogen(N), Oxygen(O), or Fluorine(F) – it creates a highly polar bond.
Any molecule that contains one of these elements bonded to H can form a “hydrogen bond” with other molecules of their own kind.

66. Hydrogen Bonding
Hydrogen bonding gives water has 4 properties that make it very special for life
Cohesion
Adhesion
Capillary action
Light solid phase density

67. Cohesion
Cohesion (n. lat. cohaerere "stick or stay together") or Cohesive attraction or Cohesive force is the action or property of molecules sticking together, being mutually attractive.

68. Cohesion

69. Adhesion
Adhesion is the tendency of dissimilar particles or surfaces to cling to one another (cohesion refers to the tendency of similar or identical particles/surfaces to cling to one another).
--usually due to polarity

70. Capillary action
Capillary action is the ability of a liquid to flow in narrow spaces without the assistance of, and in opposition to, external forces like gravity.
– trees do this in their trunks in order to get water up to the leaves
If water was not polar, plants would not be able to get the water they need against gravity.

71. Polar bonding causes liquid H2O to be more dense than Solid H2O
This is why ice floats in water.
The frozen ice layer over lake act as a blanket to insulate the rest of the water underneath, so that the whole lake doesn't freeze solid.
This makes it possible for fish to live in the winter.

72. Polar bonding causes liquid H2O to be more dense than Solid H2O
This is why ice floats in water.
The frozen ice layer over lake act as a blanket to insulate the rest of the water underneath, so that the whole lake doesn't freeze solid.
This makes it possible for fish to live in the winter.

73. Ted-ed Vid Why does ice float in water?

74. The reason behind why atoms share electrons

75. Nobel Gas electron configurations
All elements on the periodic table want to achieve stability.
They want to have the electron configuration of their nearest noble gas on the periodic table.

76. Valence Electrons
The number of valence electrons determine the reactivity of an element.
Atoms with complete valence shells (noble gases) are the most chemically non-reactive,
Those with only one electron in their valence shells (alkali metals) or just missing one electron from having a complete shell (halogens) are the most reactive.

77. The Octet Rule
The octet rule states that atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas.

78. Properties of Metals and Alloys