1. Laws of Chemical
Combination
Chem1 SY 10-11
-kaaferrer

2. What makes compounds different?

3. Law of Constant Composition
 1799 Joseph Proust
 a chemical compound contains the same elements
in exactly the same proportions (ratios) by mass
regardless of the size of the sample or source of the
compound

4.  For example, water always consists of oxygen and
hydrogen atoms, and it is always 89 percent oxygen
by mass and 11 percent hydrogen by mass

5. Does mass change during a chemical
reaction?

6. Law of Conservation of Mass
 Lavoisier heated a measured amount of mercury to
form the red oxide of mercury. He measured the
amount of oxygen removed from the jar and the
amount of red oxide formed. When the reaction was
reversed, he found the original amounts of mercury
and oxygen.

7. Law of Conservation of Mass
 1744 Antoine Lavoisier
 matter can not be created or destroyed in ordinary
chemical or physical changes.
 the mass of the reactants (starting materials) equals
the mass of the products
2Mg (s) + O2 (g) → 2MgO (s)
48.6 g 32.0 g 80.6 g

8. Example
 10 grams of CaCO3 on heating gave 4.4g of CO2
and 5.6 of CaO. Show that these observations are in
agreement with the law of conservation

9. Law of Multiple Proportions
 1803 John Dalton
 States that when two elements combine to form
more than one compound, the masses of one
element which combine with a fixed mass of the
other element are in ratios of small whole
numbers

10. Example:
 Carbon monoxide (CO): 12 parts by mass of carbon
combines with 16 parts by mass of oxygen.
 Carbon dioxide (CO2): 12 parts by mass of carbon
combines with 32 parts by mass of oxygen.
 Ratio of the masses of oxygen that combines with a
fixed mass of carbon (12 parts) 16: 32 or 1: 2

11.  Water has an oxygen-to-hydrogen mass ratio of
7.9:1.
 Hydrogen peroxide, another compound consisting of
oxygen and hydrogen, has an oxygen-to-hydrogen
mass ratio of 15.8:1.
 Ratio of the masses of oxygen that combines with a
fixed mass of hydrogen is 7.9: 15.8 or 1: 2

12. Chemical Formula Relationships

13. How many number of ATOMS?
 Ca(NO)3
 NaNO3
 Ba(IO3 )2

14. Molecular and Formula Weights
 The sum of all the atomic weights of each atom in its
chemical formula.
Formula mass = ∑ atomic masses in the formula
unit

15. SEATWORK!

16. Percentage Composition from Formulas
 Percentage by mass contributed by each element in
the substance.
%element= (# of atoms of that element)(atomic mass of that element)
------------------------------------------------------------------------
formula mass of the compound

17. Stoichiometry
 The quantitative relationships between the
substances involved in a chemical
reaction, established by the equation for
the reaction

18. Terminologies
 ATOM
◦ Smallest particle of an element
 MOLECULE
◦ Smallest unit particle of a pure substance
 ION
◦ An atom or group of bonded atoms with electrical charge
because of an excess or deficiency of electrons
 ELEMENT
◦ Pure substance; CANNOT be broken down into 2 or more
pure substances by chemical means
 COMPOUND
◦ Pure substance; CAN be broken down into 2 or more pure
substances by chemical means

19. Atomic Mass
 Atomic Mass = used to numerically indicate the mass of
an atom in its ground state, it is expressed in the non SI
unit of u
u = refers to unified atomic mass unit (formerly known as
atomic mass unit or amu)
1 amu = 1/12 the mass of carbon-12 atom,
 therefore the mass of C-12 atom is made EQUAL to 12
amu
 Carbon-12 atom is an isotope of carbon
1 amu = 1.66 x 10-24 g

20. Atomic Mass
subatomic charge Mass (amu)
particle
Neutron None 1.0087 ≈ 1
Proton Positive 1.0073 ≈ 1
Electron Negative 5.486 x 10-4 ≈ 0
 Mass of e = 1/1800 of mass of p and n so it is
negligible making the equation
Atomic mass of 12C = mass of p + mass of n

21. Atomic Mass vs.
Average Atomic Mass
Atomic Mass Average Atomic Mass
For carbon it is 12 u not 12.01 u *Used to relate the fact that the
numerical value assigned to each
element in the periodic table
reflects the average abundances
of the atoms that compose a
naturally occurring element
*Related to isotopes
*For carbon it is 12.01 u
*Chemists often will use the term
“atomic mass” when they are
actually referring to average
atomic mass of an atom.

22. Average Atomic Mass (calculation)
 Solve for the Average Atomic Mass of the element
Boron
Isotope Mass (u) Percent
Abundance
11B 11.009305 80.1
10B 10.012937 19.9
 Average atomic mass
= ∑ (mass x percent abundance)
 Where ∑ means “sum”

23. Relative Atomic Mass &
Atomic Weight
 Have different meaning from atomic mass but
synonymous with each other, although of different
historical origins
 Relative Atomic Mass
 “Ratio” of the average mass of the atom to the unified
atomic mass; dimensionless
 Atomic Weight
 The name Dalton used in the early 19th century to
numerically describe the weight of atoms relative to each
other

24. Molecular Mass
 The average mass of the molecules of a binary
compound (non metal-non metal)
 Unit: u
 Ex.: The molecular mass of carbon dioxide gas, CO2 is
28.01 u.

25. Formula Mass
 The average mass of the molecules of an ionic
compound (metal-non metal)
 Unit: u
 Ex.: The formula mass of barium chloride, BaCl2 is 208.2
u.

26. The Mole Concept
 Used to describe the number of particles (atoms,
molecules, etc.) that make up sample of matter
“One mole is the amount of any substance that
contains the same number of units as the number
of atoms in exactly 12 grams of carbon-12.”
1 mol of any substance = 6.02 x 1023 units of that
substance,
where 6.02 x 1023 Avogadro’s Number, N

27. Molar Mass
 Mass of 1 mole of substance
 SI unit: g/mol
 Provides a bridge between mass and amount

28. Conversion
Use
GRAMS FORMULA
molar MOLES Use N
UNITS
mass

29. Comparison of Atomic and Molar Mass
Atomic mass Molar mass
Unit: Unit:
amu; u g/mol
Atomic Numerically Macroscopic
level equivalent level
Total mass Mass of 1
of p+ and no mole

30. Empirical Formula
and
Molecular
Formula

31. EMPIRICAL FORMULA
 Shows the relative number of atom of each
element in the compound

32. Problem 1
A compound is found out to contain
20.0% carbon, 2.2% hydrogen, and 77.8%
chlorine. Determine EF of the compound.

33. MOLECULAR FORMULA
 Shows the actual number of atom of each
element in the compound

34. Relationship between EF & MF
MF = (EF)n
where n is an integer
therefore,
n = MMF
MEF
where M is molar mass

35. Problem 2
A compound with the empirical formula
C2H5 has a molar mass of 58.12 g/mol.
Find the molecular formula of the
compound.

36. Problem 3
The molar mass of a compound that is
54.6% carbon, 9.0% hydrogen and 36.4%
oxygen is 88 g/mol. Find MF of the
compound.

37. The Mole
 One mole is the amount of any substance that
contains the same number of units as the number of
atoms in exactly 12 grams of carbon-12.
 6.02 x 1023  Avogadro’s Number

38. MOLAR Mass
 The mass in grams of one (1) mole of a substance.
 g/mol, grams per mole
 The molar mass of any substance in grams per mole
is always numerically equal to the atomic, formula, or
molecular mass of the substance in amu.

39. Avogadro
’s Formula
atoms gram
mole weight
number