1. Chapter 2. Atoms Molecules & Ions KEY CONCEPTS
Chemical laws
Dalton's atomic theory
Atom
Molecules
Ions
Average atomic mass
Periodic table
Formula of a substance
Name of a Substance
Molecular or Covalent compounds
Ionic compounds

2. Alchemist
Early scientist observed changes of matter. They called these changes chemical reactions when there are changes in substances or the physical properties of the matter. They also observed a pattern or a repeatable observation in chemical reactions.

3. Three Chemical Laws: Law of Conservation of Mass:
Law of Constant Proportions:
Law of Multiple Proportions:

4. Law of Conservation of Mass
Total mass after a chemical reaction is same as before the reaction.
H2 + 1/2 O2 ---> H2O
Hydrogen (4g) + oxygen (32g) ----> water 36g after the reaction.

5. Law of Constant (Definite) Proportions
A given chemical compound always contains exactly the same proportion of elements by weight.
36g of water contains
4g of hydrogen and 32g of oxygen
take any other chemical compund.

6. Law of Multiple Proportions
When two elements make a series of chemical compounds, the ratio of the masses of the second elements that combine with 1 gram of the first element can always be reduced to simple whole numbers.
C O
E.g. carbon monoxide g g
carbon dioxide g g

8. Dalton’s atomic theory
All matter is composed of atoms -- the smallest particle of an element that takes part in a chemical reaction.
All atoms of an element are alike.
Compounds are combinations of atoms of one or more elements. The relative number of atoms each element is always the same.
Atoms cannot be created or destroyed by a chemical reaction. They only change how they combine with each other.

9. Models of matter
Models are commonly used to help visualize atoms and molecules.
Atom - The smallest unit of an element that has all of the properties of an element.
Molecule - The smallest unit of a pure substance that has the properties of that substance.
It may contain more that one atom and more than one element.
Ions - Charged particles formed by the transfer of electrons between atoms or molecules

10. What is an Atom? Very small particle.
Smallest particles of elements and molecules
There about 110 types of elements or Atoms.
Different atoms have different physical properties and chemical reactivity

11. Structure of the atom Atoms have a specific arrangement.
Nucleus Small, dense, positively charged
in the center of an atom.
Electrons Surround the nucleus.
Negative charge.

12. What are these? a) atoms b) nucleus c) electrons d) protons
e) neutrons
f) isotope
g) atomic number(Z)

13. Atomic Symbols Each element is assigned a unique symbol.
arsenic As potassium K
barium Ba nickel Ni
carbon C nitrogen N
chlorine Cl oxygen O
hydrogen H radon Rn
helium He titanium Ti
gold Au uranium U
Each symbol consists of 1 or 2 letters. The first is capitalized and the second is lower case.
Symbol may not match the name - often had a different name to start with.
Some elements (about 11) the names were not in English. E.g., Sodium-Na (natrium-latin),
potassium-K(kalium-latin).

14. Atomic masses Atoms are composed of protons, neutrons and electrons.
Almost all of the mass of an atom comes from the protons and neutrons.
All atoms of the same element will have the same number of protons. The number of neutrons may vary - isotopes.
Most elements exist as a mixture of isotopes.

15. Isotopes *Atoms of the same element but having different masses.
*All isotopes of an element have same atomic number
*Each isotope has a different number of neutrons.
Isotopes of hydrogen H H H
Isotopes of carbon C C C 1 2 1 3 1 13 6 14 6 12 6

16. The atomic symbol & isotopes
Isotopic symbol: atomic symbol showing atomic number (Z) and mass number (A)
Determine the number of protons, neutrons and electrons in each of the following. P 31 15 Ba
138 56 2+ U
238 92

17. Isotopes Most elements occur in nature as a mixture of isotopes.
Element Number of stable isotopes H C O Fe Sn
This is one reason why atomic masses are not whole numbers. They are based on averages.

18. Average atomic masses Most elements exits as a mixture of isotopes.
Each isotope may be present in different amounts.
The masses listed in the periodic table reflect the world-wide average for each isotope.
One can calculate the average atomic weight of an element if the abundance of each isotope for that element is known.

19. Average atomic masses Example.
Silicon exists as a mixture of three isotopes. Determine it’s average atomic mass based on the following data.
Isotope Mass (u) Abundance
28Si %
29Si %
30Si %

20. What is a mass spectrometer?
Mass spectrometer measures masses of different isotopes of an elements and their fractional or percentages abundance. (figure 7.7, page 283)
As a reference point, we use the
atomic mass unit (u) - 1/12th mass of a 12C atom.

21. Diagram of mass (positive ion) spectrometer

22. How do you calculate average Atomic Mass?
Ma x a Mb x b
= AAM
100
Ma = mass of isotope a
Mb = mass of isotope b
a = percent abundance of a
b = percent abundance of b
AAM = Average atomic mass (Reported
on the Periodic Table)

23. How do you calculate average Atomic Mass?
Ma x a Mb x b = AAM
Ma = mass of isotope a
Mb = mass of isotope b
a = fractional abundance of a
b = fractional abundance of b
AAM = Average atomic mass (Reported
on the Periodic Table)

24. Calculate the weighted average atomic mass of gallium.
Gallium in nature consists of two isotopes, gallium-69, with a mass of u and a fractional abundance of 0.601; and gallium-71, with a mass of u and a fractional abundance of
Calculate the weighted average atomic mass of gallium.
1) Ma x a Mb x b = AAM
Ma x a(%) Mb x b(%)
2) = AAM
100

25. Ma (69Ga ) = u,
a = percent abundance of 69Ga = x 100
Mb (71Ga ) = u,
b = percent abundance of 71Ga = x 100
We can obtain an equation with one unknown, AAM.
AAM = x(0.601 x 100) x(0.399x100)
100
AAM (Ga) =
AAM (Ga) = =
AAM (Ga) = 69.7 u (amu)

26. Periodic Table
Periodic table is an arrangement of all known element according to their atomic number and chemical properties.

27. Who is Dmitri Mendeleev?
Mendeleev, Dmitri ( ): Russian chemist
Mendeleev is best known for
his work on the periodic table;
arranging the 63 known
elements into a Periodic
Table based on Atomic Mass

28. Dimitri Mendeleev created this, the original,
periodic table.

29. Modern periodic table * * + + H He Li Be B C N O F Ne Na Mg Al Si P S
I A II A III A IV A V A VI A VIIA 0 H He 1 2 3 4 5 6 7 Li Be B C N O F Ne
III B IVB V B VIB VIIB VIII B IB IIB Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe * Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn + Fr Ra Lr * Gd Cm Tb Bk Sm Pu Eu Am Nd U Pm Np Ce Th Pr Pa Yb No La Ac Er Fm Tm Md Dy Cf Ho Es +

30. Information that may be in the tableAg
107.87
Silver 47
Atomic number
Name of the element
Elemental Symbol
Atomic mass

31. Vertical columns- groups,families
Horizontal columns- periods
Elements in a group have similar
chemical properties
group IA - alkali metal: Li, Na, K Rb, Cs, Fr
group IIA- alkaline earth metals:
Be, Mg, Ca, Sr, Ba, Ra
group VIIA - Halogens: Cl, Br, I, At
group 0 - Noble gases: He, Ne, Ar, Kr, Xe, Rn

32. A group or family Groups are assigned Roman numerals with an A or B H
I A II A III A IV A V A VI A VIIA 0 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar
III B IVB V B VIB VIIB VIII IB IIB K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Lr La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

33. A row or period 1 2 3 4 5 6 7 Periods are assigned numbers H He Li BeC N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Lr La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

34. Common group names He H Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca
I A II A III A IV A V A VI A VIIA 0 He H Li Be B C N O F Ne Na Mg Al Si P S Cl Ar
III B IVB V B VIB VIIB VIII B IB IIB K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Lr La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

35. Elemental states at room temperature
Solid
Liquid
Gas H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe * Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn + Fr Ra Lr * La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb + Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

36. The known elements 112 elements are currently known 89 are metals
31 are radioactive
22 are synthetic (all radioactive)
11 occur as gases
2 occur as liquids
Let’s take a look at them on the table.

37. What are these? Transition Metals Actinides Lanthanides
Semimetals or Metalloids
Ionic Charges
Poly atomic ions and their charges

38. Naming Chemical Compounds
Binary compunds: Combinations of two differerent elements
Common names: Water-H2O
Vinger, Ammomia, Chloroform
Systematic Names: Names given starting from
element names and/or following certain rules
Naming depend on the type of compound
Ionic Compounds:
Name:Sodium chloride
Formula:NaCl
Molecular or Covalent Compounds:
Carbon tetrachloride- CCl4

39. Types of Compounds Ionic compounds: Metal + non-metal
Type I : ionic compound
(fixed charge) NaCl
Type II ionic compound
FeCl2 and FeCl3, SnCl2 and SnCl4
Covalent Compounds:
non-metal + non-metal
sulfur dioxide: SO2

40. Formula Molecular Weight ? and Formula Weight?
Formula are used to represent elements and compound.
For molecular compounds, formula tell how many of each kind of atom are in a molecule.
For ionic compounds, formula tell the simples ratio of cations and anions.
Molecular Weight ? and Formula Weight?

41. What is Empirical Formula
Simple whole number ratio of each atom expressed in the subscript of the formula.
Molecular Formula = C6H12O6 of glucose
Empirical Formula = CH2O
Emiprical formula is calculated from % composition

42. Ions
Ions are charged particles formed by the transfer of electrons between elements or combinations of elements.
Cation - a positively charged ion.
Ca Ca e-
Anion - a negatively charged ion.
F2 + 2e F-

43. Ionic compounds Some simple ions Na+ Mg2+ Al3+ Cl- O2- N3- Cations
Anions
Formula for some ionic compounds
NaCl MgCl2 AlCl3
Na2O MgO Al2O3
Na3N Mg3N2 AlN

44. Naming Ionic compounds
When an element forms only one compound with a given anion.
name the cation
name the anion using the ending (-ide)
NaCl sodium chloride
MgBr2 magnesium bromide
Al2O3 aluminum oxide
K3N potassium nitride

45. Metals with multiple charges
Transition metals.
Here it is easier to list the ones that have a single common oxidation state.
All Group 3B
Ni, Zn, Cd Ag
Lanthanides and actinides

46. Ionic compounds: a) Type I ionic compound NaCl, CaCl2, AlCl3
sodium chloride, calcium chloride,
Aluminum chloride
b) Type II ionic compound
FeCl2 and FeCl3
SnCl2 and SnCl4
iron(II)chloride, iron(III)chloride, tin(II)chloride, tin(IV)chloride

47. Examples FeCl2 iron(II) chloride FeCl3 iron(III) chloride
SnS tin(II) sulfide
SnS2 tin(IV) sulfide
AgCl silver chloride
CdS cadmium sulfide
Note
Some transition metals only have a single state
so the roman numeral may be omitted.
Some main group metals, with high atomic number
have more than one state, roman numbers are used.

48. Naming Hydrates
Hydrates are substances that include water into their formula.
The water is not actually part of the chemical substance and this is reflected in the way the formula is written.
Example: CuSO4 . 5 H2O

49. Example Name the following compounds Na2S \______ sodium sulfide HCl
\______ hydrogen chloride
CaH2
\______ calcium hidride
SrO
\______ strontium oxide

50. Polyatomic ions Name the following AgNO3 \______ silver nitrate
Mg(C2H3O2)2
\_______ magnesium acetate
Fe2(SO4)3
\____ iron(III) sulfate

51. Polyatomic ions
A special class of ions where a group of atoms tend to stay together.
NH4+ ammonium
NO3- nitrate
SO42- sulfate
OH- hydroxide
O22- peroxide
Your book contains (Table 2.4, page 56) a more complete list.

52. Polyatomic ions
When a compound contains a polyatomic ion, you simply use the given name.
NH4Cl ammonium chloride
NaOH sodium hydroxide
KMnO potassium permanganate
(NH4)2SO ammonium sulfate

53. Naming Ionic Compounds Summary
Simple rules that will keep you out of trouble most of the time.
Groups IA, 2A, 3A (except Tl) only have a single oxidation state that is the same as the group number - don’t use numbers.
Most other metals and semimetals have multiple oxidation states - use numbers.
If you are sure that a transition group element only has a single state, don’t use a number.

54. Molecules
When atoms of nonmetals combine to form compounds, molecules result.
A molecule
is the smallest unit of an element or a compound.
has the chemical properties of the element or compound.
does not have a net electrical charge.

55. Types of molecules Diatomic Molecules that contain two atoms.
Homoatomic
Two or more atoms of one element (elements).
Hetroatomic
Contain at least two atoms of two or more elements (compounds).

56. Naming Molecular or Covalent compounds
Naming is similar to ionic compound but prefixes are added in front of the name of elements to indicate number of atoms in the molecule.
Dinitrogen pentoxide- N2O5

57. Naming Molecular Compounds
A simple set of rules can be used.
Naming is similar to ionic compound
name elements in the order they appear in the formula.
use the ending (-ide) for the second element listed in the formula
But;
use prefixes to indicate how many atoms there are of each type.
mono = 1 tri = 3
di = 2 tetra = 4.

58. Naming covalent compounds
N2O5
CO2 CO
SiO2
ICl3
P2O5
CCl4
dinitrogen pentoxide
carbon dioxide
carbon monoxide
silicon dioxide
iodine trichloride
diphophorous pentoxide
carbon tetrachloride
The rule may be modified to improve how a name sounds.
Example - use monoxide not monooxide.

59. Naming Acids
formula starts with H
HCl
HNO3
H2SO4
HClO3
H3BO3

60. HClO hypochlorous ClO¯ ”hypochlorite
HClO2 chlorous ClO2¯ chlorite
HClO3 chloric ClO3¯ chlorate
HClO4 perchloric ClO4¯ perchlorate
HNO3 nitric NO3 ¯ nitrate
HNO2 nitrous NO2 ¯ nitrite

61. Naming bases formula ends with OH NaOH sodium hydroxide
Ba(OH) barium hydroxide
KOH potassium hydroxide
NH4OH
Ca (OH)2

62. Nomenclature overview
Now that a large number of nomenclature rules have been introduced, we need to review them.
Simple binary ionic compounds
Ionic compounds of metals with multiple charges
Compounds containing polyatomic ions
Simple molecular compounds
It’s useful to be able to identify which system to use by looking at the chemical.

63. A bit more on nomenclature
When the first element is a metal then usually:
If only one other element is present and
the second element is a non-metal -
name the metal first - as element.
Name non-metal second with -ide ending
If more than one other element is present -
The rest is most likely a polyatomic ion
so use the name from the table in book.

64. A bit more on nomenclature
Is a metal present
as the first element? No
Use prefixes
(mono, di, tri ...)
Yes
Can the metal have
more than one
oxidation state? No
Roman numerals
are not needed.
Yes
Use Roman numerals

65. More examples Give the chemical name of CaCO3
First element is a metal - Ca
Ca is NOT a transition metal - calcium
- Roman numerals are not necessary
CO3 is a polyatomic ion - carbonate
Proper name calcium carbonate

66. Even more examples BCl3 \____ boron trichloride Na3PO4
\___ sodium phosphate
AlBr3
\_____ aluminum bromide
UF6
\______ uranium(VI) fluoride
MgNH4PO4
\_ magnesium ammonium phosphate

67. Getting the formula from the name
iron(III) hydroxide
\_______ Fe(OH)3
sulfur dioxide
\_______ SO2
silicon hexachloride
\_______ SiCl6
silver cyanide
\_______ AgCN

68. Models
Scientists often rely on models to help them understand their observations.
Chemists’ models of atoms and molecules are much larger than the real thing.
Models can be mathematical or even imaginary.

69. Molecular representations
H2O - water
CH3CH2OH - ethyl alcohol