1. Examples:
N2O dinitrogen monoxide
(also called dinitrogen oxide)
CO carbon monoxide
Cl2O dichlorine heptaoxide
P4O tetraphosphorous decaoxide
PF phosphorous pentafluoride
N dinitrogen

2. Chemical Reactions and Chemical Equations – An introduction
Chemical reaction: The transformation of one or more chemicals into different compounds.

3. Examples: 2 Na + 2 H2O 2 NaOH + H2 (sodium) (water) (sodium dihydrogen hydroxide)

4. Examples: 2 Na + 2 H2O 2 NaOH + H2 (sodium) (water) (sodium dihydrogen hydroxide) Cu + 4 HNO3 Cu(NO3)2 + 2 NO2 (copper) (nitric acid) (copper (II) nitrate) (nitrogen dioxide) + 2H2O

5. Examples: 2 Na + 2 H2O 2 NaOH + H2 (sodium) (water) (sodium dihydrogen hydroxide) Cu + 4 HNO3 Cu(NO3)2 + 2 NO2 (copper) (nitric acid) (copper (II) nitrate) (nitrogen dioxide) + 2H2O Zn + 2HCl(aq) ZnCl2 + H2

6. Reactants: The starting substances in a chemical reaction. E. g
Reactants: The starting substances in a chemical reaction. E.g. Na and H2O in the first reaction.

7. Reactants: The starting substances in a chemical reaction. E. g
Reactants: The starting substances in a chemical reaction. E.g. Na and H2O in the first reaction.
Products: Substances as a result of a chemical reaction. E.g. Cu(NO3)2, NO2, and H2O in the second reaction.

8. Reactants: The starting substances in a chemical reaction. E. g
Reactants: The starting substances in a chemical reaction. E.g. Na and H2O in the first reaction.
Products: Substances as a result of a chemical reaction. E.g. Cu(NO3)2, NO2, and H2O in the second reaction.
Balanced equation: The number of atoms of each element (free or part of a compound) is the same on both sides of the equation.

9. The three equations given are all balanced.
An example of a non-balanced equation is:
H O H2O
The commonly employed definition of balanced equation is the smallest set of integers that leads to a balanced equation.

10. So,
H ½ O H2O
and
4 H O H2O
are not balanced equations. The balanced equation is:
2 H O H2O

11. The coefficients tell us how many molecules of each species are reacting.

12. Measurements and Units
Most physical quantities we encounter in chemistry have units.
Old system: The cgs system – centimeter, gram, second were some of the key units in use – hence the abbreviation cgs

13. The SI system:

14. The SI system:
Basic Quantity Name of Unit Symbol

15. The SI system:
Basic Quantity Name of Unit Symbol
length meter m

16. The SI system:
Basic Quantity Name of Unit Symbol
length meter m
mass kilogram kg

17. The SI system:
Basic Quantity Name of Unit Symbol
length meter m
mass kilogram kg
time second s

18. The SI system:
Basic Quantity Name of Unit Symbol
length meter m
mass kilogram kg
time second s
temperature kelvin K

19. The SI system:
Basic Quantity Name of Unit Symbol
length meter m
mass kilogram kg
time second s
temperature kelvin K
amount of
substance mole mol

20. Derived Units: These are obtained from the basic units just given.
Examples: The SI unit of volume is derived as follows:
volume = length3 (think about the volume of a cube).
Since the unit of length is m, then the unit of volume is m3

21. Hence the units of density are kg m-3
Derived Units: These are obtained from the basic units just given.
Examples: The SI unit of volume is derived as follows:
volume = length3 (think about the volume of a cube).
Since the unit of length is m, then the unit of volume is m3
The SI unit of density:
Hence the units of density are kg m-3

22. There are some non-SI units that are in common use in chemistry
There are some non-SI units that are in common use in chemistry. A common unit of volume is the liter.
1 liter = 1 cubic decimeter (1 dm3)
= (10 cm)3
= cm3
and liter = ml,
that is 1 ml = 1 cm3

23. The basic units are not always convenient for reporting measurements
The basic units are not always convenient for reporting measurements. Various prefixes are often employed.
Examples: 1 kg = 1000 g
1 km = 1000 m
1 nm = m
1 cm = 0.01 m

24. The factor-dimensional method of calculation (also called the factor-label method)

25. The factor-dimensional method of calculation (also called the factor-label method)
Suppose we want to convert 45.6 m into centimeters. From the conversion factor
100 cm = 1m

26. The factor-dimensional method of calculation (also called the factor-label method)
Suppose we want to convert 45.6 m into centimeters. From the conversion factor
100 cm = 1m
we can write or
These are unit conversion factors.

27. Number of centimeters =
Multiplication of a quantity by a unit conversion factor will change its units, and adjust the value of the quantity for the new unit system used.
For the present problem we have:
Number of centimeters =
= x 103 cm

28. If we take the wrong conversion factor, the result would be:
Number of centimeters =
= x 10-1 m2 cm-1

29. If we take the wrong conversion factor, the result would be:
Number of centimeters =
= x 10-1 m2 cm-1
In this case we can see that the units on the right-hand side do not match with what we expect on the left-hand side of the expression.

30. Second example: convert 45. 6 m to inches
Second example: convert 45.6 m to inches. Conversion factors are 1 m = 100 cm and 1 inch = 2.54 cm (exactly)

31. Second example: convert 45. 6 m to inches
Second example: convert 45.6 m to inches. Conversion factors are 1 m = 100 cm and 1 inch = 2.54 cm (exactly)
Number of inches =
= x 103 inches

32. Second example: convert 45. 6 m to inches
Second example: convert 45.6 m to inches. Conversion factors are 1 m = 100 cm and 1 inch = 2.54 cm (exactly)
Number of inches =
= x 103 inches
In this factor label calculation, two conversion factors are strung together. You can multiply any number of conversion factors in a string.

33. Read the sections in the book on: Chapter 1 Chapter 2, Atomic view of matter (sections 2.1-2.4)

34. Atomic number, Mass number, and Isotopes
Atomic Number: The number of protons in the nucleus of each atom of an element.
The standard symbol for the atomic number is Z.

35. Atomic number, Mass number, and Isotopes
Atomic Number: The number of protons in the nucleus of each atom of an element.
The standard symbol for the atomic number is Z.
For a neutral atom, the atomic number also indicates the number of electrons.

36. Example: The atomic number of C is 6
Example: The atomic number of C is 6. Thus, carbon has 6 protons and 6 electrons.

37. Mass number: The total number of protons and neutrons in the nucleus of each atom of an element.
The symbol employed is A.

38. Mass number: The total number of protons and neutrons in the nucleus of each atom of an element.
The symbol employed is A.
The number of neutrons present in the nucleus = A - Z

39. Mass number: The total number of protons and neutrons in the nucleus of each atom of an element.
The symbol employed is A.
The number of neutrons present in the nucleus = A - Z
For example, the mass number of fluorine is 19 and the atomic number is 9, so the number of neutrons present in the fluorine nucleus is
= 10.

40. Isotopes: Atoms having the same atomic number, but different mass numbers.

41. Isotopes: Atoms having the same atomic number, but different mass numbers.
The common symbol for denoting the atomic number and mass number of element X is

42. Examples of isotopes: hydrogen has three common isotopes
hydrogen deuterium tritium

43. Examples of isotopes: hydrogen has three common isotopes
hydrogen deuterium tritium
Hydrogen is the only element for which special symbols are used for different isotopes.
D for deuterium
T for tritium

44. Key Point: The chemical properties of an element are primarily determined by the number of protons and electrons, not by the number of neutrons present. For this reason, isotopes of the same element are chemically similar.

45. Atomic mass scale
By international agreement, the reference adopted for an atomic mass scale is the assignment of exactly a mass of 12 for the most abundant isotope of carbon,

46. Atomic mass scale The atomic mass unit (amu) is defined by
By international agreement, the reference adopted for an atomic mass scale is the assignment of exactly a mass of 12 for the most abundant isotope of carbon,
The atomic mass unit (amu) is defined by
1 amu =

47. Most naturally occurring elements contain more than one isotope – this means that when the atomic mass of a naturally occurring element is determined, it is an average quantity that is determined.

48. Most naturally occurring elements contain more than one isotope – this means that when the atomic mass of a naturally occurring element is determined, it is an average quantity that is determined.
Example: the atomic masses of and
are … amu (exact number) and amu and their natural abundances (relative amounts of the isotopes present) are 98.89% and 1.11% respectively.

49. The weighted average
atomic mass of carbon
= amu

50. Molar mass
Chemists frequently use the term molecular weight or for an atomic system, atomic weight.

51. Molar mass
Chemists frequently use the term molecular weight or for an atomic system, atomic weight.
The molar mass is the mass of 1 mol of a substance expressed in units of g/mol.

52. The mass of a single atom of carbon -12 is determined from:
= 1.99 x g

53. Molecules
The simplest molecules consist of only two atoms – these are called diatomic molecules.

54. Molecules
The simplest molecules consist of only two atoms – these are called diatomic molecules. Examples: O2, H2, and Cl2 These are dioxygen, dihydrogen, and dichlorine. Unfortunately, it is also very common to refer to these as oxygen, hydrogen, and chlorine.

55. Molecules containing more than two atoms are called polyatomic molecules.

56. Molecular mass
The molecular mass of a molecule is the sum of the atomic masses of the atoms present in the molecule. Example: The molecular mass of H2O is given by 2 x( ) = amu

57. The molar mass (frequently termed the molecular weight) of a compound is the molecular mass expressed in grams/mole (units abbreviated as g/mol).

58. The molar mass (frequently termed the molecular weight) of a compound is the molecular mass expressed in grams/mole (units abbreviated as g/mol). Example: The molecular mass of H2O is amu and hence the molar mass of water is g/mol.

59. Laws of Chemical Combination
Law of Definite Proportions: In a given compound, the elements are always combined in the same proportion by mass.

60. Laws of Chemical Combination
Law of Definite Proportions: In a given compound, the elements are always combined in the same proportion by mass. Example: water contains hydrogen and oxygen in the proportion of g to g.